Unit 1

1. Atomic Structure and Periodicity

The development of modern atomic theory revealed much about the inner structure of
atoms. It was learned that an atom contains a very small nucleus composed of positively
charged protons and uncharged neutrons, surrounded by a much larger volume of space
containing negatively charged electrons. The nucleus contains the majority of an atom’s
mass because protons and neutrons are much heavier than electrons, whereas electrons
occupy almost all of an atom’s volume. The diameter of an atom is on the order of 10−10
m, whereas the diameter of the nucleus is roughly 10−15 m, about 100,000 times smaller.

In order to understand the structure of an atom, different types of atomic models were
developed using different instruments. Some of the models are:
1.1. Democritus theory (300 BC)
The Greek philosopher Democritus expressed his own postulate and he states that:
Matter consists of very small, indivisible particles which are called as atoms. Atoms are
derived from Greek words ATOMOS which mean indivisible into smaller particle.
Atoms are indivisible particles. Explains certain natural occurrences such as the existence
of elements
Atoms-and the protons, neutrons, and electrons that compose them—are extremely small.
Atomic theory describes about the historical models of the atom, especially the external
structure of atoms and how atoms combine to form molecules. This theory is a scientific
theory that deals the nature of matter; states that matter is composed of discrete units
which called atoms. Atom is the smallest particle into which an element can be divided
and still be the same substance. Atoms are so small until recently, no one had ever seen it.
But ideas, or theories, about atoms have been around for over 2,000 years. The theory of
the atom has had a long history. The ancient Greeks postulated that matter exists in the
 form of atoms. But they did not base their theory on experiment and they cannot develop
additional ideas about atoms. Different scholars (researchers) give their idea on the
atomic theory.
Limitation of Democritus atomic theory
 His theory does not support by experiment but only theoretical view
 Does not have any information about subatomic particles (electron, proton and
neutron). Due to the lack of experiment, Democritus’ idea was not accepted by many
of researchers or scholars because Experimental evidence from early scientific
investigations provided support for the notion of “atomism” and gradually gave rise to
the modern definitions of elements and compounds.
1.2. Dalton’s Atomic Theory
John Dalton (1766-1844), an English schoolteacher, developed the first useful atomic
theory of matter around 1808. His findings were based on experiments and also from
laws of chemical combination. He used fundamental laws of chemical combination just
described as the basis of an atomic theory. His theory involved many assumptions:
 All substances are composed of tiny, indivisible particles which called “atoms”.
 Atoms of the same element are identical in size, mass and properties but, atoms of
different elements have different properties (he doesn’t recognize isotopes).
 Compounds are formed by the union of two or more different elements
 Dalton model is known as the solid sphere model (like billiard balls)
Importance and Improvement on previous model
 Explains how atoms combine to form molecules.
 Explains chemical change better than the particle theory.
 Define conservation of mass and definite proportion.
Dalton’s theory put down a corner stone for modern atomic theories since he uses the
law of chemical combination.
1.2.1. Laws of Chemical Combination
Various chemical reactions take place according to certain laws, known as the Laws of
chemical combination. There are three common laws of chemical combinations. These
are:
A. Law of conservation of mass/the law of mass action: It also known as the law of
indestructibility of matter since this law states that “matter is neither created nor
destroyed in the course of chemical reaction rather it may change from one form to
other”. The total mass of materials after a chemical reaction is same as the total mass
before reaction. In short this law states that “during any physical or chemical changes, the
total mass of product equal to the total mass of reactants”. For example, in an experiment
63.5g of copper combines with 16g of oxygen to give 79.5g of cupric oxide (a black
oxide of copper).

2Na + Cl2 →2NaCl,

2x23+2x35.5=117=2(23+35.5)=117
B. Law of constant or definite composition: According to this law, in a given
compound, the constituent elements are always combined in the same proportions by
mass, regardless of the origin or mode of preparation of the compound. What this law
means is that when elements react chemically, they combine in specific proportions, not
in random proportions. A sample of pure water which obtains river, sea, well, spring
whatever the source, always contains 88.9% by mass of oxygen and 11.1% by mass of
hydrogen.

C. Law of multiple proportions: According to this law, when two elements A and B
combine to form more than one chemical compounds then different weights of A, which
combine with a fixed weight of B, are in proportion of simple whole numbers. Or when
two/more elements combine to form two/more compounds, the mass of one element
combines with a fixed mass of another element. Example: Carbon monoxide (CO): 12
parts by mass of carbon combines with 16 parts by mass of oxygen. Carbon dioxide
(CO2): 12 parts by mass of carbon combines with 32 parts by mass of oxygen. Ratio of
the masses of oxygen that combines with a fixed mass of carbon (12 parts) is 16: 32 or 1:
2

Limitations of Dalton’s atomic theory

 Does not include the existence of the nucleus
 Does not explain the existence of ions or isotopes
  Does not have any information on subatomic particles (electron, proton and
neutron).

Example 1. A nitrogen-oxygen compound is found to have an oxygen-to-nitrogen mass

ratio of 1.142 g oxygen for every 1.000 g of nitrogen. Which of the following oxygen-to-
nitrogen mass ratio(s) is/are possible for different nitrogen-oxygen compound(s)? a 2.285
c 0.571 b 1.000 d 2.500 2. Determine, which of the following oxygen-to-nitrogen mass
ratio(s) is/are also possible for nitrogen-oxygen compound. (Refer Example 2.1 No. 1 for
required information) a 0.612 c 1.713 b 1.250 d 2.856

Solution: 1. The given compound has 1.142 g of oxygen and 1.000 g of nitrogen.
Response (a) has 2.285 g of oxygen for the same 1.000 g of nitrogen. The ratio of the
masses of oxygen, 2.285:1.142, is almost exactly 2:1. Response (a) seems to be correct
possibility, so is response (c). Here the ratio is 0.571:1.142 = 0.500 = 1:2. Responses (b)
and (d) are not possibilities. They yield ratios of 1.000:1.142 = 0.875 and 2.500: 1.142 =
2.189, respectively. Neither of these can be expressed as a ratio of small whole numbers.
2. By the same method, a 0.612:1.142 = 0.536:1 is not possible b 1.250:1.142 = 1.095:1
is not possible c 1.713:1.142 = 1.5:1 or 3:2, is possible d 2.856:1.142 = 2.5:1 or 5:2 is
possible

Exercise 1

1. List the postulates of Dalton that continue to have significance (are retained in modern
atomic theory). 2. Match the atomic theory statements in part A with the matching items
in part B.
1.3. The concept of the atom
2.3.1. Discoveries of subatomic particles

Name location Charge(C) Unit Mass(amu) Mass(g) Discovered by

charge

Electron Outside −1.602× 10−19 -1 0.00055 0.00091 × 10−24 Joseph John

nucleus Thomson

Proton Nucleus 1.602 ×10−19 +1 1.00727 1.67262 × 10−24

Neutron Nucleus 0 0 1.00866 1.67493 × 10−24 James

Chadwick

A. Cathode Rays

In 1879, the English scientist William Crookes (1832-1919) experimented with gas
discharge tubes. When a very high electrical potential (~ 10,000 volts) is applied across a
gas taken in a discharge tube of a very low pressure (~ 0.001 torr) some radiations are
emitted from cathode. These radiations are called cathode rays. Fig 2.1, shows emission
of cathode rays in a discharge tube. At this stage the glass walls of the discharge tube
opposite to the cathode starts glowing with a faint greenish light. It is now known that
this greenish glow on the walls is due to the bombardment of the glass wall with the
cathode rays.

Cathode rays normally travel in straight lines, but are deflected when a magnet is brought
nearby (Figure 2.1b).

Figure 1 (a) Electric discharge in an evacuated tube (b) The cathode ray is bent. in the presence of a
magnet.
An English physicist Joseph John Thomson (1856-1940) in 1897 studied the behaviour of
cathode rays in electric and magnetic fields, Thomson established clearly that the rays
consist of negatively-charged particles. Moreover, his experiments showed that the
particles were identical, regardless of the materials from which the belectrodes were
made or the type of gas in the tube. Thomson concluded that these negatively charged
particles were constituents of every kind of atom. We now call these particles electrons, a
term that had been coined by the Irish Physicist George Stoney in 1891 to describe the
smallest unit of electric charge. Cathode rays are beams of electrons. In 1909, Robert A.
Millikan, an American physicist, determined the charge on the electron by observing the
behaviour of electrically-charged oil drops in an electric field. Based on careful
experiments, Millikan established the charge on an electron as e = –1.602 × 10–19 C.
From this value and the value for me /e, we can calculate the mass of an electron.

B. Radioactivity and Discovery of Nucleus

I. Radioactivity
Radioactivity is the spontaneous emission of radiation from the unstable nuclei of certain
isotopes. Isotopes that are radioactive are called radioactive isotopes or radioisotopes. For
example,

are radioactive isotopes. Radioactive decay is defined as a nuclear breakdown in which

particles or (electromagnetic) radiation is emitted. Shortly after the discovery of
radioactivity, three types of rays were identified in the emanations from radioactive
substances. One type called alpha (α) particles which consist of particles that have a mass
that is about four times that of a hydrogen atom. They also have a charge twice the
magnitude of an electron but positive rather than negative. An alpha particle is now
known to be a doubly-ionized helium atom, that is, He2+.
A second type of radiation was shown to consist of negatively-charged particles, identical
to cathode rays. These particles are called beta (β) particles, which are electrons, coming
from inside the nucleus.
The third type of radiation, called gamma (γ) rays, is a form of electromagnetic radiation
much like the X-rays but of even higher energy. Like X-rays, but unlike alpha and beta
particles, gamma rays are a form of energy and not a form of matter.
II. Discovery of Nucleus
In 1910, the New Zealand chemist and physicist, Ernest Rutherford, who had studied
with J.J. Thomson decided to use α-particles to probe the structure of atoms.
Together with his associate Hans Geiger, Rutherford carried out a series of experiments
using very thin foils of gold and other metals as targets for α-particles from a radioactive
source.

Figure 2. a. Rutherford’s experimental design for measuring the scattering of α-particles by a piece of gold
foil. b. Magnified view of α-particles passing through and being deflected by the nucleus.

C. Discovery of Neutron
Except for the lightest hydrogen isotope, protium (1H), atoms have more mass than is
indicated by the numbers of their protons (Before the 1930’s, protons were considered as
the sole contributors to the mass of an atom). For example, a helium nucleus, with two
protons has a mass four times that of hydrogen. If all the mass came from the protons, a
helium atom would have only twice the mass of a hydrogen atom. The reason for this
“excess” mass puzzled scientist for several years.
In the 1920s and early 1930s, alpha particles were used as projectiles to bombard a
variety of materials. Bombardment of beryllium atoms produced a strange, highly
penetrating form of radiation. In 1932, James Chadwick (1891-1972) showed that this
radiation was best explained as a beam of neutral particles. These particles, called
neutrons, were found to have about the same mass as protons but no electric charge.
This discovery finally provided an explanation for the mysterious excess mass. A helium
atom has two protons and two neutrons. Because protons and neutrons have roughly the
same mass (and electrons have almost no mass) the helium atom should have about four
times the mass of the hydrogen atom. The mass of a neutron, mn = 1.67493 × 10–27 kg, is
about 1840 times the mass of electron.
1.4. Atomic Mass and Isotopes

The number of protons in the nucleus of an atom is its atomic number (Z). This is the
defining trait of an element: Its value determines the identity of the atom. For example,
any atom that contains six protons is the element carbon and has the atomic number 6,
regardless of how many neutrons or electrons it may have. A neutral atom must contain
the same number of positive and negative charges, so the number of protons equals the
number of electrons. Therefore, the atomic number also indicates the number of electrons
in an atom. The total number of protons and neutrons in an atom is called its mass
number (A). The number of neutrons is therefore the difference between the mass number
and the atomic number: A – Z = number of neutrons.

Atoms are electrically neutral if they contain the same number of positively charged
protons and negatively charged electrons. When the numbers of these subatomic particles
are not equal, the atom is electrically charged and is called an ion. The charge of an atom
is defined as follows:

Atomic charge = number of protons − number of electrons

As will be discussed in more detail later in this chapter, atoms (and molecules) typically
acquire charge by gaining or losing electrons. An atom that gains one or more electrons
will exhibit a negative charge and is called an anion. Positively charged atoms called
cations are formed when an atom loses one or more electrons. For example, a neutral
sodium atom (Z = 11) has 11 electrons. If this atom loses one electron, it will become a
cation with a 1+ charge (11 − 10 = 1+).

A neutral oxygen atom (Z = 8) has eight electrons, and if it gains two electrons it will
become an anion with a 2- charge (8 − 10 = 2-).
Atomic mass unit and average atomic mass because each proton and each neutron
contribute approximately one amu to the mass of an atom, and each electron contributes
far less, the atomic mass of a single atom is approximately equal to its mass number (a
whole number). However, the average masses of atoms of most elements are not whole
numbers because most elements exist naturally as mixtures of two or more isotopes.

The mass of an element shown in a periodic table or listed in a table of atomic masses is a
weighted, average mass of all the isotopes present in a naturally occurring sample of that
element. This is equal to the sum of each individual isotope’s mass multiplied by its
fractional abundance.

For example, the element boron is composed of two isotopes: About 19.9% of all boron
atoms are 10B with a mass of 10.0129 amu, and the remaining 80.1% are 11B with a
mass of 11.0093 amu. The average atomic mass for boron is calculated to be:
Calculation of Average Atomic Mass: A meteorite found in central Indiana contains
traces of the noble gas neon picked up from the solar wind during the meteorite’s trip
through the solar system. Analysis of a sample of the gas showed that it consisted of
91.84% 20Ne (mass 19.9924 amu), 0.47% 21Ne (mass 20.9940 amu), and 7.69% 22Ne
(mass 21.9914 amu). What is the average mass of the neon in the solar wind?

Exercise 2. There are two isotopes of lithium found on earth. Isotope 6Li (6.01512 mu)
accounts for 7.42% of the total, and isotope 7Li (7.01600 mu) accounts for the remaining
92.58%. What is the average atomic mass of lithium?

Solution:

2. What is the mass number of an isotope of tin that has 66 neutrons and 50 protons?
3. Calculate the number of protons and neutrons for 24 12Mg and 88 38Sr.
4. Why do isotopes of an element have similar chemical properties?
5. Element X is toxic to humans in high concentration but essential to life at low
concentrations. Identify element X whose atoms contain 24 protons and write the symbol
for the isotope with 28 neutrons
6. Copper (Cu: atomic mass 63.546 mu) contains the isotopes 63Cu (mass = 62.9298 mu)
and 65Cu (mass = 64.9278 mu). What percent of a Cu atom is 65Cu?
7. The element chlorine contains two isotopes: 35Cl, which has a mass of 34.97 mu, and
37
Cl, which has a mass of 36.97 mu. Calculate the percentage of each chlorine isotope.
The average atomic mass of chlorine is 35.5 mu.
8. Carbon exists as the isotopes carbon-12, with a fractional abundance of 0.9890 and a
mass of exactly 12 mu, and carbon-13, with a fractional abundance of 0.0110 and a mass
of 13.00335 mu. Calculate the average atomic mass of carbon
1.5. The Bohr Theory of the hydrogen atom

In 1913, Niels Bohr, a Danish physicist, who had worked with Rutherford, combined
ideas from classical physics and the new quantum theory to explain the structure of the
hydrogen atom. He suggested a model for the hydrogen atom that predicted the existence
of line spectra. In doing so, he was able to explain the spectrum of radiation emitted by
hydrogen atoms in gas-discharge tubes.

Based on the work of Planck and Einstein, Bohr made the revolutionary assumption that
certain properties of the electron in a hydrogen atom – including energy, can have only
certain specific values. That is to say, these properties are quantized. Bohr proposed the
following three postulates for his model.

1. The hydrogen atom has only certain allowable energy levels, called stationary states.
Each of these states is associated with a fixed circular orbit of the electron around the
nucleus.
2. The atom does not radiate energy while in one of its stationary states. That is, even
though it violates the ideas of classical physics, the atom does not change energy while
the electron moves within an orbit.
The electron moves to another stationary state (orbit) only by absorbing or emitting a
photon whose energy equals the difference in the energy between the two states.

The subscripts f and i represent the final and the initial states, respectively. The Bohr
radius, denoted by ao (ao = 0.0529 nm) can be calculated using the formula

where n is a positive integer which is called quantum number. r is the radius of the orbit
and is given by:

where εo is the vacuum dielectric constant (εo = 8.854 × 10–12 C V–1 m–1).

A spectral line results from the emission of a photon of specific energy (and therefore, of
specific frequency), when the electron moves from a higher energy state to a lower one.
An atomic spectrum appears as lines rather than as a continuum because the atom’s
energy has only certain discrete energy levels or states.
In Bohr’s model, the quantum number n (n = 1, 2, 3 ...) is associated with the radius of
the electron’s orbit, which is directly related to the atom’s energy. The lower the quantum
number, the smaller is the radius of the orbit and the lower is the energy level of the
atom. When the electron is in the orbit closest to the nucleus (n = 1), the atom is in its
lowest (first) energy level, which is called the ground state. By absorbing a photon whose
energy equals the difference between the first and second energy levels, the electron can
move to the next orbit. This second energy level (second stationary state) and all higher
levels are called excited states. The hydrogen atom in the second energy level (first
excited state) can return to the ground state by emitting a photon of a particular
frequency:

Where Eg and Ee represent the ground and the excited energy states, respectively. When a
sample of atomic hydrogen absorbs energy, different hydrogen atoms absorb different
amounts. Even though each atom has only one electron, so many atoms are presented that
all the allowable energy levels (orbits) are populated by electrons. When an electron
drops from orbits with n > 3 (second excited state), the infrared series of spectral lines is
produced i.e, Paschen Series. The visible series arises from the photons emitted when an
electron drops to the n = 2 orbit i.e, Balmer Series (first excited state), and the ultraviolet
series arises when these higher energy electrons drop to the n = 1 orbit (ground state).
Figure 3. Representation of the observed spectral lines of the hydrogen atom.

Since a larger orbit radius means a higher atomic energy level, the farther the electron
drops, the greater is the energy (higher v, shorter λ) of the emitted photon. The spectral
lines of hydrogen become closer and closer together in the short wavelength (high
energy) region of each series because the difference in energy associated with the jump
from initial state ( ni ) to the final state (nf ) becomes smaller and smaller as the distance
from the nucleus increases.
Having made this basic assumption, Bohr was then able to use classical physics to
calculate properties of the hydrogen atom. In particular, he derived an equation for the
electron energy (En ). Each specified energy value (E1, E2, E3...) is called an energy level
of the atom. A very useful result from Bohr’s work is an equation for calculating the
energy levels of an atom,
 –18
Where A is the constant, has a value of, A = 2.18 × 10 J. The number n is an integer
called the principal quantum number (n = 1, 2, 3, ...). Z is the charge of the nucleus. The
negative sign in the equation appears because it is defined as zero energy when the
electron is completely moved form the nucleus, i.e. En = 0 when n =, so, En < 0 for any
smaller n.
A can be expressed in terms of Rydberg constant as A = hCR , where R is Rydberg
constant with a value of 1.0967 x107M-1 . For the H atom, Z = 1, so we have

Therefore, the energy of the ground state n = 1 is –2.18 × 10–18 J. This equation is easily
adapted to find the energy difference between any two levels:

• When an electric discharge is passed through a sample of hydrogen, the H2

molecules dissociate into atoms, and the electron in a particular H atom is excited.
• These states are transient and the electron falls back to a lower energy state by
emitting energy as it does so.
• The consequence is the observation of spectral lines in the emission spectrum of
hydrogen.
• The spectrum (a small part of which is shown in below Figure consists of groups
of discrete lines.
• In1885, Balmer pointed out that the wave number of the spectral lines observed in
the visible region of the atomic spectrum of hydrogen is given by V= R ( 1/22-
1/n2)
where, R = Rydberg constant for =1.097 x 107 m-1 = 1.097 x 105 cm-1
and n is an integer 3, 4, 5 . . .
 • Wave number = reciprocal of wavelength(v=1/λ); convenient (non-SI) units are
‘reciprocal centimetres’, cm-1

Figure 4. Spectrum of hydrogen

Note that: the line become close together as the wave length decrease until the
continuum reached
Similar equation will be developed to hold for the other series in
the H- spectrum

Example 2. Calculate the wave number if an electron transit from1st energy level to 5th
energy level in Uv, visible and infrared region

Example 3. Calculate wave number for third transition in:

A. Lyman series B. Balmer series C. Paschen series D. Bracket

series

Example 4. Calculate wave length for 6th transition in Balmer series

• Solution:

• given:- ni= 1 nf= 5

i. V= R(1/12-1/n2)-----------uv region

= 1.097 x 107 m-1 (1/12-1/52)= 1.05X 107 m-1

ii. V= R(1/22-1/n2) ---------- visible region

V=1.097 x 107 m-1 (1/4-1/52) = 1.05X 107 m-1

= 0.21 X1.097 x 107 m-1 = 2.3 x 106 m-1

iii. V= R(1/32-1/n2) ---------- visible region

V=1.097 x 107 m-1 (1/9-1/52)= 1.097 x 107 m-1 x 0.071

=0.077887x 107 m-1 = 7.789x 105 m-1

1.6. The quantum mechanical model of the atom

Bohr’s model was very important because it introduced the idea of quantized energy
states for electrons in atoms. This feature is incorporated in our current model of the
atom, the quantum mechanical model of the atom.
The sophisticated mathematical description of atomic structure based on the wave
properties of subatomic particles is called wave mechanics or quantum mechanics.
Principally, the Austrian physicist Erwin Scrödinger developed a model of the hydrogen
atom based on the wave nature of the electron in the late 1920s.
Mathematical equations describing the nature of electron waves in atoms are fundamental
to the modern picture of the atom. The wave equations that are acceptable solutions to the
Schrödinger equation are called wave functions. To obtain one of these acceptable
solutions, we must assign integral values called quantum numbers to three quantities in
the wave equation. This requirement is similar to the requirement for an integral value of
n in the Bohr equation for the hydrogen atom.
In contrast to the precise planetary orbits of the Bohr atom, the wave mechanics picture
of the hydrogen atom is less certain. Instead of determining the exact location of the
electron, we can only speak of the probability of the electron being found in certain
regions of the atom. Or, if we adopt the view that the electron is just a cloud of negative
electric charge, we can only speak of the charge densities in various parts of the atom.
quantum mechanical model of the atom is called Modern Atomic Theory (Electron Cloud
Mode) It states that “electrons are always moving around the nucleus in a "cloud" of
energy levels”. It also states that the small particle atom is consists of three parts such as,
proton, neutron and electron.
 The electrons are found around the nucleus whereas; protons and neutrons are
found in the central part of atom which is called nucleus.
 Describes that, the mass of one atom is the summation of the number of proton
and the number of neutrons. A= p + n
 NB: the mass of an atom is occupying by the nucleus whereas the large space of
the atom is occupying by electrons.

1.6.1. Quantum Model of Atom

 Bohr's theory worked well for hydrogen atom (one electron).
 Complications such as elliptical rather than circular orbits were introduced in an
attempt to fit the data to Bohr’s theory.
 The Bohr Atomic Theory and its modifications that forced the theorists to work
hard to explain the spectroscopists’ observations.
 In spite of their efforts, the Bohr Atomic Theory eventually proved
unsatisfactory; the energy levels are valid only for the hydrogen atom.
 An important characteristic of the electron, its wave nature, still needed to be
considered.
 According to the de Broglie equation, proposed in the 1920s, all moving
particles have wave properties described by the equation

 where λ=wavelength of the particle, h = Planck's constant, m = mass of the

particle, v = velocity of the particle.
 Particles massive enough to be visible have very short wavelengths, too small to
be measured.
 Electrons, on the other hand, have wave properties because of their very small
mass.
  Electrons moving in circles around the nucleus (in Bohr’s Atomic theory), can be
thought of as forming standing waves that can be described by the de Broglie
equation.
 However, we no longer believe that it is possible to describe the motion of an
electron in an atom so precisely.
 This is a consequence of another fundamental principle of modern physics,

1.6.2. The Heisenberg's Principle

Heisenberg formulated what is known as the Heisenberg uncertainty principle, which
states that it is impossible to know simultaneously both the momentum and the position
of a particle with certainty. Mathematically,
 Heisenberg's uncertainty principle, which states that there is a relationship
between uncertainties in the location and momentum of an electron moving in
the x direction:

Where:
 Δx= uncertainty in the position of the electron
 Δp= uncertainty in the momentum of the electron.
 Working with Heisenberg’s Principle, Schrödinger developed a compromise which
calculates both the energy of an electron and the probability of finding an electron
at any point in the molecule.
1.6.3. Quantum Numbers
An atomic orbital is specified first by three quantum numbers that are associated
respectively, with the orbital's size (energy), shape, orientation and, later, independent of
these three quantum numbers, the electron spins. The first three sets of quantum numbers
have a hierarchical relationship: the size-related number limits the shape related number,
the shape-related number in turn limits the orientation-related number.
Three among the four quantum numbers characterize the orbitals in the atom. That is,
they describe the orbital or the space the electron is supposed to occupy. The fourth
quantum number is used to describe the spin of the electrons that occupy the orbitals.
 The four quantum numbers are:
 1. The principal quantum number (n) is a positive integer having values n = 1, 2, 3, ...
.
 It gives the following information:
 (i) Relative size of the orbital or the relative distance of the electron from the nucleus.
Size of orbital increases with the increase of principal quantum number n.
 (ii) Energy of the orbital. Higher the n value, greater is the energy. For example:
when the electron occupies an orbital with n = 1, the hydrogen
 atom is in its ground state and has lower energy than when the electron occupies an
orbital with n = 2 (first excited state).
 (iii) Maximum number of electrons present in any shell (given by the formula 2n2).
2. The azimuthal quantum number (l) is also known as angular momentum or subsidiary
quantum number. It is an integer having values from 0 to (n – 1). For an orbital with n =
1, l can have a value only of 0. For orbitals with n = 2, l can have a value of 0 or 1; for
those with n = 3, l can be 0, 1 or 2; etc. So, the number of possible l values equals the
value of n. For a given value of n, the maximum possible value of l is (n – 1). The
azimuthal quantum number gives the following information: Number of subshell present
within any shell.
(ii) It describes the shape of the orbital and is sometimes also called the orbital-shape
quantum number.
3. The magnetic quantum number (ml) is also known as the orbital-orientation quantum
number. It is an integer having values from –l through 0 to +l. The possible values of an
orbital's magnetic quantum number are set by its angular momentum quantum number
(that is, l determines ml). An orbital with l = 0 can have only ml = 0. However an orbital
with l = 1, can have ml value of –1, 0, or + 1; thus there are three possible orbitals with l
= 1 each with its own spatial orientation. The number of possible ml values or orbitals for
a given l value is (2l + 1). It prescribes the orientation of the orbital in the three-
dimensional space about the nucleus.
4. The electron spin quantum number (ms ) has only two possible values, +½
(represented by the arrow, ) and – ½ (represented by the arrow ). The name electron
spin quantum suggests that electrons have a spinning motion. However, there is no way
to attach a precise physical reality to electron spin. The quantum numbers specify the
energy states of the atom.
• The atom's energy levels or shells are given by the n value.
• The atom's sublevels or subshells are given by the n and l values. Each level contains
sublevels that designate the shape of the orbital.
• The atom's orbitals are specified by the n, l and ml values. Thus, the three quantum
numbers that describe an orbital express its size (energy), shape and spatial orientation.
Each sublevel is designated by a letter:
l = 0, is an s sublevel
l = 1, is a p sublevel
l = 2, is a d sublevel
l = 3, is a f sublevel
The letters s, p, d, and f are derived from the names of spectroscopic lines: s, sharp; p,
principal; d, diffuse; and f, fundamental. Sublevels are named by joining the n value and
the letter designation. For example, the sublevel (subshell) with n = 2, l = 0 is called the
2s sublevel; the only orbital in this sublevel has n = 2, l = 0 and ml = 0. A sublevel with n
= 3, l = 1, is a 3p sublevel. It has three possible orbitals: one with n = 3, l = 1 and ml = –1
; another with n = 3, l = 1 and ml = 0 and the third n = 3, l = 1, and ml = +1.
For a given principal quantum number, n, the total number of orbitals is determined as:
Number of orbitals = n2 in a shell.
Similarly, the number of orbitals in each subshell is determined as:
 Number of orbitals in a subshell = 2l+1.
Example 5
1. What values of the angular momentum quantum number (l) and magnetic quantum
number (ml ) are allowed for a principal quantum number (n) of 3? How many orbitals
are allowed for n = 3?
2. Give the name, magnetic quantum numbers, and numbers of orbitals for each sublevel
with the following quantum numbers:
a. n = 3, l = 2 c. n = 5, l = 1
b. n = 2, l = 0 d. n = 4, l = 3
3. What is wrong with each of the following quantum number designations and/or
sublevel names?

Exercise 2
1. Give the sublevel notation for each of the following sets of quantum numbers.
a. n = 3, l = 2 c. n = 4, l = 1
b. n = 2, l = 0 d. n = 4, l = 3
2. Indicate whether each of the following is a permissible set of quantum numbers. If the
set is not permissible, state why it is not.
a. n = 3, l = 1, ml = +2 b. n = 4, l = 3, ml = –3 c. n = 3, l = 2, ml = –2
d. n = 0, l = 0, ml = 0 e. n = 3, l = 3, ml = –3
3. Consider the electronic configuration of an atom:
a. What are the n, l and ml quantum numbers corresponding to the 3s orbital?
b. List all the possible quantum number values for an orbital in the 5f sub shell.
c. In which specific subshell will an electron be present if the quantum numbers n = 3, l =
1, and ml = –1?
4. Which of the quantum numbers relates to the electron only? Which relate (s) to the
orbital?
1.7. The electromagnetic radiation and atomic spectra

1.7.1. Electromagnetic radiation

In 1873, James Clerk Maxwell proposed that light consists of electromagnetic waves.
According to his theory, an electromagnetic wave has an electric field component and a
magnetic field component. Further, his theory accurately describes how energy, in the
form of radiation, propagates through space as electric and magnetic fields.
Electromagnetic radiation is the emission and transmission of energy in the form of
electromagnetic waves. The wave properties of electromagnetic radiation are described
by two interdependent variables, frequency and wavelength. Wavelength (λ, Greek
lambda) is the distance between any point on a wave and the corresponding point on the
next wave; that is, the distance the wave travels during one cycle.

Figure 5. Electromagnetic wave

In Figure 5, two waves with different wave lengths (λ) and thus different frequencies (ν)
are shown. Wavelength is commonly expressed in meters, but since chemists often deal
with very short wavelengths the nanometre, picometer and the angstrom are also used.
Frequency (ν Greek nu) is the number of cycles that pass a given point in space per
second, expressed in units of s–1 or hertz (Hz). The speed of the electromagnetic wave
(light), c (distance travelled per unit time, in meters per second), is the product of its
frequency (cycles per second) and itswavelength (metres per cycle),

In vacuum, light travels at a speed of 2.99792458 × 108 m s–1 (3.00 × 108 m s–1 to three
significant figures).
The speed of an electromagnetic wave depends on the nature of the medium through
which the wave is travelling. The speed of an electromagnetic wave in medium (c) is the
product of its wavelength and its frequency.

Another characteristic of a wave is its amplitude, the height of the crest (or depth of the
trough) of the wave. The amplitude of an electromagnetic wave is a measure of the
strength of its electric and magnetic fields. Thus, amplitude is related to the intensity of
the radiation, which we perceive as brightness in the case of visible light.

Figure 6. Amplitude (intensity) of waves.

Light of a particular shade of red, for instance, always has the same frequency and
wavelength, but it can be dim (low amplitude) or bright (high amplitude) Figure 6.
Visible light occupies a small portion of the continuum of radiant energy which is known
as the electromagnetic spectrum (Figure 7). The electromagnetic waves in the different
spectral region travel at the same speed but differ in frequency and wavelength.
The long wavelength, low-frequency portion of the spectrum comprises the microwave
and radio wave regions. The infrared (IR) region overlaps the microwave region on one
end and the visible region on the other.
We perceive different wave lengths (or frequencies) of visible light as different colours,
from red (λ= 750 nm) to violet (λ= 400 nm). Light of a single wavelength is called
monochromatic (Greek “one colour”), whereas light of many wavelengths is
polychromatic (Greek “many colours”). White light is polychromatic.
The region adjacent to visible light on the short-wavelength end consists of ultraviolet
(UV) radiation. Still shorter wavelengths (higher frequencies) make up the X-ray and
gamma ray (γ-ray) regions.
Thus, a TV signal, Infrared (IR) light, and a gamma ray emitted by a radioactive element
differ principally in frequency and wavelength.
Figure 7. Regions of the electromagnetic spectrum.
These wave and particle nature of light is linked by energy since both particle and wave
have energy.
 E=hv , the relationship between energy and frequency
 E= hc/λ, the relationship between energy and wavelength
  E= hvc, the relationship between energy and wave number
The wavelength(λ): is the distance between any two adjacent identical points of
the wave.
The frequency (v) is the number of wave crests passing a given point per unit time; and
is usually expressed in cycles per second or, more commonly, simply as 1/s with “cycles”
understood. It is the number of cycles or oscillations per time.
Wave number: is the reciprocal of wavelength.
Period (T): the time required to complete one complete cycle.
Black body: is a body which absorbs all radiate energy from the sun.
Photon: is a packet of light energy which carries hv amount of quantized energy.
Quantized energy: the motion of the electron is not free. The electron is bound to the
atom by the attractive force of the nucleus and consequently quantum mechanics predicts
that the total energy of the electron is quantized. White light is a combination of light of
many different wavelengths.
When passed through a prism, it is spread into its constituent wavelengths, resulting in a
band spectrum.
A band spectrum resembles a rainbow and contains many different wavelengths of light.
When light from a gas discharge tube is passed through a prism, the result is a line
spectrum.
In contrast to a band spectrum, a line spectrum contains only certain discreet wavelengths
of light.
Example 7. The yellow light given off by a sodium lamp has a wavelength of 589 nm.
What is the frequency of this radiation? 2. A dental hygienist uses X-ray (λ = 1.00 Å) to
take a series of dental radiographs while the patient listens to an FM radio station (λ =
325 cm) and looks out the window at the blue sky (λ = 473 nm). What is the frequency
(in s–1) of the electromagnetic radiation for each source?
Solution:
1. Rearranging Equation 2.1, and ν = co/λ. We insert the value for c and ? and then
convert nm to m. This gives us

2. Because we are provided with the wavelengths, we can find the frequencies from
Equation 2.1. Since co has units of m s–1, we first convert the entire wavelength to
metres.
For X-ray
Exercise 4. Some diamonds appear yellow because they contain nitrogenous compounds
that absorb purple light of frequency 7.23×1014 s–1. Calculate the wavelength (in nm) of
the absorbed light. 2. The FM station broadcasts traditional music at 102 MHz on your
radio. Units for FM frequencies are given in megahertz (MHz). Find the wavelength of
these radio waves in meters (m), nanometers (nm), and angstrom (Å).

1.8. Electronic Configuration and Periodic Table

Electron configuration is the way of filling\distribution of electrons among the given
orbitals. In writing electronic configurations or electron filling of orbitals, we follow the
Aufbau’s principle, Hund’s rule, and the Pauli Exclusion Principle. These are:

A.The Aufbau’s Principle (building up principle): explains the order of filling

electrons in various orbitals of an atom. Filling begins with the orbitals in the lowest
energy, or most stable, shells and Thus, the 1s orbital fills first, then the 2s, followed by
the 2p and the 3s orbitals.
Figure 8. Increasing order of filling orbitals.

B. Hund's rule: It’s Principle: Equal energy orbitals (degenerate orbitals) are each
occupied by a single electron before the second electrons of opposite spin enters the
orbital. In other words, each of the three 2p orbitals (2px, 2py and 2pz) will hold a single
electron before any of them receives a second electron. It states that each degenerate
orbital, (e.g.2px, 2py, and 2pz) must first receive one electron before any of the orbitals
can receive a second electron.

C. Pauli Exclusion Principle: It’s Principle. No two electrons can have the same four
quantum numbers. i.e. they must differ in at least one of the four quantum numbers.
- each orbital contains a maximum of two electrons.

 These two electrons must have opposite values for the spin, which is
generally indicated by showing the electrons as arrows pointing up () or
down ().
Excise 5. write the ground electron configuration of based on example given
The reason for these exceptions to the aufbau principle is not completely understood, but
it seems that the half-filled 3d subshell of chromium (3d5) and the fully filled 3d subshell
of copper (3d10) lends a special stability to the electron configurations. Apparently,
having a half-filled 4s subshell and a half-filled 3d subshell gives a lower energy state for
a Cr atom than having a filled 4s subshell.
Exercise 6
1. Write the electron configuration for the Co3+, Cl– , Al+, Cr, As–, and Cu.
2. Write the electron configuration and the orbital diagram of the first excited state of
sodium. (Hint: The outermost electron is excited).
3. What is the electron capacity of the nth energy level? What is the capacity of the fourth
energy level?

1.9. Periodic table and Periodic Properties of Elements

The periodic law (Modern periodic table) tells us that if we arrange the elements in
order of increasing atomic number, we periodically encounter elements that have similar
chemical and physical properties. Elements in a group have similar chemical and physical
properties, and those within a period have properties that change progressively across the
table.
1.9.1. Classification of the Elements
Representative or main group elements: These consist of all s- and p-block elements.
The chemical properties of the representative elements are determined by the number of
valence electrons in their atoms.
Transition elements: These are d-block elements. There are four series of transitional
elements, 3d, 4d, 5d and 6d depending on the energy levels of d-orbitals.
Inner transition elements: These are the f-block elements. There are two series of f-block
elements, 4f and 5f series called lanthanides and actinides, respectively. The periodic
table is unable to include the inner transition elements in its main frame. They have been
allotted the same single place in the periodic table though their electronic configurations
are not identical. Besides, the variation in their properties is not much.1s
The general properties of metals and nonmetals are distinct. Physical and chemical
properties that distinguish metals from nonmetals are summarized in Table:
S blok elements
IA
IIA p- blok elments
d-blok metals

4f Lanthanide series
5f Acthenide series

Figure 9. the general periodic presentation

s-blockmetals Nobel gas
d-block Lanthanide
p-block metals
Actinide series Metaloid

Nonmetals Hydrogen
Figure 10. the modern periodic table

The vertical columns are referred to as groups or families, and the horizontal rows are
called periods. Elements in a group have similar chemical and physical properties, and
those within a period have properties that change progressively across the table. Several
groups of elements have common names that are used so frequently they should be
learned. The Group IA elements, except H, are referred to as alkali metals, and the
Group IIA elements are called the alkaline earth metals. The Group VIIA elements are
called halogens, which means “salt formers,” and the Group VIIIA elements are called
noble(or rare) gases.

1.9.2. Periodic Variation in Physical Properties

a) Atomic Radii
As we move across the periodic table, atoms become smaller due to increasing effective
nuclear charges. As we move from left to right across a period in the periodic table,
atomic radii of representative elements decrease as a proton is added to the nucleus and
an electron is added to a particular shell.

The effective nuclear charge, Zeff experienced by an electron in an outer shell is less than
the actual nuclear charge, Z. This is because the attraction of outer-shell electrons by the
nucleus is partly counterbalanced by the repulsion of these outer-shell electrons by
electrons in inner shells. This concept of a screening, or shielding, effect helps us
understand many periodic trends in atomic properties. Consider an atom of lithium; it has
two electrons in a filled shell, 1s2, and one electron in the 2s orbital, 2s1. The electron in
the 2s orbital is fairly effectively screened from the nucleus by the two electrons in the
filled 1s orbital, so the 2s electron does not “feel” the full 3 charge of the nucleus. The
effective nuclear charge, Zeff, experienced by the electron in the 2s orbital:

Example;11Na, Zeff =11-10 = +1, 12Mg, Zeff =12-10= +2, 13Al, Zeff =13-10= +3
b) Ionic Radii
Ionic Radii of metals always greater than the corresponding cations because they lost
electrons from shells which reduced electron-electron repulsions. Example;Mg (1.6 Ǻ) >
Mg2+(0.85 Ǻ). Simple negatively charged ions (anions) are always larger than the neutral
atoms from which they are formed because there is addition of extra electron to the shell
which increased electron. Example: Cl- (1.67 Ǻ) > Cl (1.1 Ǻ). Both sizes of cations and
anions decrease from left to right across a period whereas both cation and anion sizes
increase going down a group. Within an isoelectronic series, radii decrease with
increasing atomic number because of increasing nuclear charge.
Isoelectronic ions: ions with the same number of core electrons. Na+, Mg+2, Al+3, F-, O-
2
, N-3 all contain 10 electrons; all have the same electron configuration as Ne but, in
terms of size,
N-3>O-2>F->Na+>Mg+2>Al+3; the ion with the greater number of protons in an
isoelectronic series will be the smallest due to the greater nuclear charge pulling the
electrons in closer.
In general,

Arrange the following ions in order of increasing ionic radii: (a) Ca2+, K+, Al3+ +, (b) Se2-,
Br-, Te2- Answers: a) Al3+< Ca2+ <K , b)Br-< Se2- < Te2-

c) Ionization Energy
Ionization (potential) Energy: the minimum amount of energy required to remove an
electron from ground state (loosely bounded) atom in gaseous state. It measures how
outermost electrons held by the nucleus.
X (g) + IE → X+ (g) + e-

The first ionization energy (IE1), also called first ionization potential, is the minimum
amount of energy required to remove the most loosely bound electron from an isolated
gaseous atom to form an ion with a 1+charge. Elements with low ionization energies (IE)
lose electrons easily to form cations.

General trends in first ionization energies of A group elements with position in the
periodic table. Exceptions occur at Groups IIIA and VIA. The first ionization energies for
the Group IIIA elements (B, Al, Ga, In, Tl) are exceptions to the general horizontal
trends. They are lower than those of the IIA elements in the same periods because the
IIIA elements have only a single electron in their outermost p orbitals. Less energy is
required to remove the first p electron than the second s electron from the outermost
shell, because the p orbital is at a higher energy (less stable) than an s orbital within the
same shell (n value) Effective nuclear charge, Zeff, increases going from left to right
across a period. The increase in effective nuclear charge causes the outermost electrons to
be held more tightly, making them harder to remove. The first ionization energies of the
Group IIA elements (Be, Mg, Ca, Sr, Ba) are significantly higher than those of the Group
IA elements in the same periods. This is because the Group IIA elements have higher Zeff
values and smaller atomic radii. Thus, their outermost electrons are held more tightly
than those of the neighboring IA metals. It is harder to remove an electron from a pair in
the filled outermost s orbitals of the Group IIA elements than to remove the single
electron from the half-filled outermost s orbitals of the Group IA elements. The general
left-to-right increase in IE1for each period is interrupted by a dip between Groups VA
(N, P, As, Sb, Bi) and VIA elements (O, S, Se, Te, Po). Presumably, this behavior is
because the fourth np electron in the Group VIA elements is paired with another electron
in the same orbital, so it experiences greater repulsion than it would in an orbital by itself.
This increased repulsion apparently outweighs the increase in Zeff, so the fourth np
electron in an outer shell (Group VIA elements) is somewhat easier to remove (lower
ionization energy) than is the third np electron in an outer shell (Group VA elements).
The first ionization energies therefore generally increase from left to right across the
periodic table. The order of ionization energies decreased as 3rd > 2nd >1st because it is
too difficult to remove an electron from cation specie.

In general ionization energy increase across the period since Zeff increase by one; the size
of atoms decrease which results the outermost electrons bound tightly with the nucleus.
d) Electron Affinity (EA):

The electron affinity (EA) of an element may be defined as the amount of energy released
when an electron is added to an isolated gaseous atom to form a negative charge.

X (g) + e- → X-(g) + EA

This process can be either endothermic or exothermic, depending on the element. You
can see that many of these elements have negative values of EA, which means that
energy is released when the gaseous atom accepts an electron. However, for some
elements, energy is required for the atom to become negatively charged and the value of
their EA is positive.

Electron Affinity (EA): Energy is always required to bring a negative charge (electron)
closer to another negative charge (anion). So the addition of a second electron to anion to
form an ion with a -2 charge is always endothermic. Thus, electron affinities of anions
are always positive.
X-(g) + e- → X2-(g) + EA
Most elements have no affinity for an additional electron and thus have an electron
affinity (EA) equal to zero. We can represent the electron affinities of helium and
chlorine as

Halogens have ns2-np5 electron cofiguration which are most electron affinities to form
noble gas cofiguration.
 increase except noble gas

decrease

e) Electro-negativity (EN):

The electro-negativity (EN) of an element is a measure of the relative tendency of an

atom to attract electrons towards it when it is chemically combined with another atom.
Elements with high electro-negativities (nonmetals) often gain electrons to form anions.
Example; F (4.0) > O (3.5) > Cl=N (3.0). Elements with low electro-negativities (metals)
often lose electron to form cations.

For the representative elements, electro-negativities usually increase from left to right
across periods except noble gas and decrease from top to bottom within groups

increase except noble gas

decrease
.
Arrange the following elements in order of increasing electro-negativity? B, Na, F, O

answer, Na< B< O< F

f) Metallic character:

Metallic character increases from top to bottom and decreases from left to right with
respect to position in the periodic table. Nonmetallic character decreases from top to
bottom and increases from left to right in the periodic table.

General trends in metallic character of A group elements with position in the periodic
table.
In modern periodic table, elements are arranged as the function of their atomic number.
Main group elements are elements contain s-block metals and p-block metal, metalloid
and nonmetals. This periodic table constructed with the columns (family) and the
horizontal row (periods). Periodicity is the variation of physical properties in periodic
table such as atomic size, electronegativity, ionization energy, metallic character and
others as summarized below.

Advantages of Periodic Classification of the Elements

Some of the advantages of periodic classification of elements are: 1. The classification of
elements is based on the atomic number, which is a fundamental property of an element.
2. The reason for placing isotopes at one place is justified as the classification is on the
basis of atomic number. 3. It explains the periodicity of the properties of the elements and
relates them to their electronic configurations. 4. The position of the elements that were
misfits on the basis of mass number (anomalous pairs like argon and potassium) could be
justified on the basis of atomic number. 5. The lanthanides and actinides are placed
separately at the bottom of the periodic table. 6. The table is simple, systematic and easy
way for remembering the properties of various elements as it is based on the electronic
configuration.
Unit summary
 Cathode rays (electrons) are produced when electricity passes through evacuated
tubes. X-rays form when cathode rays strike matter.
 Radioactivity is the emission of radiation by unstable nuclei, and the most
common types of radiations are alpha (α) particles, beta (β) particles and gamma
(γ) rays. Alpha particles are helium nuclei; beta particles are electrons; and
gamma rays are high frequency electromagnetic radiation similar to X-rays.
 Rutherford's atomic model is that of a very small positively charged nucleus and
extra-nuclear electrons. The nucleus consists of protons and neutrons and contains
practically all the mass of an atom. Atomic masses and relative abundances of the
isotopes of an element can be established by mass spectrometry. The atomic mass
of the element is the average of these mass numbers based on their percentage
abundances Electromagnetic radiation is the transmission of electric and magnetic
 fields as a wave motion. The waves are characterized by their velocity in a
medium: c = νλ. A light source that emits an essentially unbroken series of
wavelength components has a continuous spectrum. Only a discrete set of
wavelength components is present in the emission spectrum of an atom.
 Einstein's explanation of the photoelectric effect views light as packets of energy
called photons. The energy of the photon (Eph) is given by the expression E = hν
where h is Planck's constant.
 Bohr's theory requires the electron in a hydrogen atom to be in one of a discrete
set of energy levels. The fall of an electron from a higher to a lower energy level
releases a discrete amount of energy as a photon of light with a characteristic
frequency.
 Bohr's theory accounts for the observed atomic spectrum of hydrogen atom.
 The electron in a hydrogen atom can be viewed as a matter-wave enveloping the
nucleus. The matter-wave is represented by a wave equation, and solutions of the
wave equation are wave functions. Each wavefunction is characterized by the
value of four quantum numbers: the principal quantum number, n; the angular
momentum quantum number l; the magnetic quantum number, ml; and the spin
quantum number, ms. Wave functions with acceptable values of the three are
called atomic orbitals. An orbital describes a region in an atom that has a high
probability of containing an electron or a high electron change density. Orbitals
with the same value of n are in the same principal energy level or principal shell.
Those with the same value of n and of l are in the same sublevel or subshell. The
shapes associated with orbitals depend on the value of l. Thus, the s orbital (l = 0)
is spherical and the p orbital (l = 1) is dumbbell-shaped.
 The n, l and ml quantum numbers define an orbital, but a fourth quantum number
is also required to characterize an electron in an orbital - the spin quantum
number, ms. This quantum number may have either of two values: +½ or –½.
 The wave mechanical treatment of the hydrogen atom can be extended to multi-
electron atoms, but with this essential difference: principal energy levels are (i)
lower than those of the hydrogen atom and (ii) split, that is, having different
energies for the different subshells.
 Electron configuration refers to the distribution of electrons among orbitals in an
atom. Introduced here are the subshell notations (or "s,p,d, f") and the orbital
diagram. Key ideas required to write a probable electron configuration are: (i)
electrons tend to occupy the lowest energy orbitals available; (ii) no two electrons
in an atom can have all four quantum numbers alike; and (iii) where ever possible,
electrons occupy orbitals singly rather than in pairs.
 The Aufbau principle describes a hypothetical process of building up one atom
from the atom of preceding atomic number. With this principle and the idea cited
above, it is possible to predict probable electron configurations for many of the
elements. In the Aufbau process, electrons are added to the s or p subshell of
highest principal quantum number in the representative or main group elements.
In transition elements, electrons go into the d subshell of the second last shell, and
in the inner transition elements, into the f subshell of the third last shell.
 Elements with similar valence-shell electron configurations fall in the same
group of the periodic table. For A-group elements, the group number corresponds
to the number of electrons in the principal shell of highest quantum number. The
period number is the same as the highest number of principal shell containing
electrons (the outer shell). The division of the periodic table into s, p, d and f
blocks greatly assists in the assignment of probable electron configurations.
 Certain atomic properties vary periodically, when atoms are considered in terms
of increasing atomic number. The properties and trends considered in this unit are
those of atomic radius, ionic radius, ionization energy and electron affinity.
Values of these atomic properties strongly influence physical and chemical
properties of the elements.
Review questions
Part I: Multiple Choice Type Questions
1. The number of neutrons in an atom of 226 88Ra is:
a. 88 b. 82 c. 138 d. 314
2. Which of the following are usually found in the nucleus of an atom?
a. Protons and neutrons only b. Protons, neutrons and electrons
c. Neutrons only d. Electrons and neutrons only
3. An atom has an atomic number of 31 and a mass numbers of 70. How many
electrons will it have in its valence shell?
a. 5 b. 4 c. 3 d. 2
4. Which of the following would produce a line spectrum rather than a continuous
spectrum?
a. Sunlight b. Excited hydrogen atom
c. A normal filament light bulb d. A yellow (sodium) street light
5. Among the following, which colour corresponds to light of the highest frequency?
a. Green b. Red c. Yellow d. Blue
6. Which one of the following is not a valid electronic configuration?
a. 2, 8, 8, 2 b. 2, 8, 9, 1 c. 2, 6 d. 2, 8, 4
7. Which of the given elements will have the electronic configuration?
1s22s22p63s23p64s2.
a. Neon atoms b. Chlorine atoms c. Magnesium ion d. Calcium atoms
8. Which of the following elements has the lowest first ionization energy?
a. Potassium b. Sodium c. Calcium d. Argon
9. How many 3d electrons are present in the ground state of chromium atom?
a. 9 b. 4 c. 6 d. 5
10. The first ionization energy of aluminium is slightly lower than that of magnesium
because:
a. magnesium has a higher nuclear charge
b. the outer electron in aluminium is in a p-orbital not an s-orbital
c. in aluminium the electron is being lost from a doubly filled orbital
d. the radius of the aluminium atom is greater than that of the magnesium
atom
11. Which of the following atoms would have the highest fourth ionization energy?
a. P b. Si c. N d. C
12. How many unpaired electrons are there in the Cr3+ ion?
a. 6 b. 3 c. 1 d. 0
13. Which of the following species would require the highest energy for the removal of
one electron?
a. Mg2+ b. Na+ c. Ne d. F–
14. Which of the following has the lowest electronegativity?
a. Carbon b. Magnesium c. Beryllium d. Boron
15. Which of the following has the smallest radius?
a. Na b. Na+ c. Mg d. Mg2+
Part II: Answer the following questions:
16. Identify the following subatomic particles:
17. The number of these in the nucleus is equal to the atomic number.
18. The particle that is gained or lost when ions are formed.
19. The particle that is not found in the nucleus.
20. The particle that has no electrical charge.
21. The particle that has a much lower mass than the others.
22. Calculate the number of protons, neutrons and electrons in the following:

23. Carbon has atomic number 6. It comprises three isotopes, the first with 6 neutrons,
the second with 7 neutrons, the third with 8 neutrons.
24. Calculate the mass numbers of the three isotopes and represent them in the form of
xCy
25. Explain what is meant by “isotope”
63 65
26. In naturally occurring copper isotopes, 29Cu contributes 69.09% and 29Cu,

30.91%. Calculate the relative atomic mass of copper. (Accurate mass determined;
63
29Cu = 62.9298 mu, 6529Cu = 64.9278 mu)
27. Two particles X and Y have the following composition: X: 17 protons, 18 neutrons,
17 electrons Y: 17protons, 18 neutrons, 18 electrons
28. What is the relationship between these particles?
29. Will these two particles have similar chemical properties? Explain why?
30. Arrange the following in order of increasing ionization energy: Li, Na, Ne, N, O
31. Explain the following:
32. The first ionization energy of beryllium is greater than that of boron.
33. The first ionization energy of oxygen is less than that of nitrogen.
34. The first ionization energy of lithium is greater than that of sodium.
35. The electron configuration of a particular metal cation M3+ is [Ar] 3d 2.
36. Identify the corresponding metal.
37. Write the electron configuration of the metal atom.
38. Arrange the following in order of increasing atomic radius Mg, Cs, Ca, Al, Ba.
39. Explain briefly, why potassium always occurs as a +1 ion in its compounds and
calcium as a +2 ion.
40. Arrange the atoms (ions) in each of the following groups in order of increasing size
based on their location in the periodic table.
a. Mg2+, O2–, Na+, F –, Al3+ b. Ne, N3–, F –, Na+, C 4–
c. F, Be, C, B, Li d. K+, S2–, As3–, Cl–, Ca2+
41. Excited sodium atoms emit light with a wavelength of 589 nm. Calculate the:
a. frequency of the light, and b. energy of one of these photons
in joules
42. A hydrogen atom is excited to the n = 8 energy level. It emits a photon of light as it
falls to the n = 2 energy level. Calculate the:
a. wavelength of light emitted, and b. frequency of the light
emitted
43. The electron of a hydrogen atom is in the n = 3 level. What is its energy?
44. Calculate the wavelength of the light emitted when an electron falls from n = 3 to
the n = 1 state in hydrogen atom.
45. The photon emitted by a cyclotron has a velocity of 1.50 × 103 m s–1. What is the
wavelength of this photon? Given that the mass of photon = 1.676 × 10–27 kg and
Planck’s constant = 6.62 × 10–34 J.s.
46. Write the number and the letter for the orbital that corresponds to the following
pairs of n and l quantum numbers:
a. n = 3, l = 1 b. n = 4, l = 0 c. n = 3, l = 2 d. n = 5,
l=3
47. Write the electron configurations for the following atoms and ions:
a. Fe3+ b. V c. Cr3+ d. Al3+
48. Identify the transition element (s) from the following:
a. 40Zr b. 88Ra c. 56Fe d. 36Kr
2. Chemical Bonding and Structure (10 hours)
2.1. Introduction
Octet Rule
You have studied in your earlier classes that noble gases have very stable electron
arrangements such as 2; 2, 8; 2, 8, 8 and their outer shells are fully saturated. The first
three are shown in Figure 1 and explains why noble gases are so reluctant to form
compounds with other elements.

Figure 1. First three noble gases.

The noble gases have very stable electron configuration, as reflected by their high
ionization energies, low electron affinity and general lack of reactivity. Because all noble
gases (except He) have eight valence electrons, many atoms undergoing reaction also
attain eight valence electrons. The ns2np6 electron configuration of the valence shell of all
noble gas (except Helium) atoms, is commonly called an octet of electrons. The octet
rule, a useful generalization that applies to all types of bonding which states that when
atoms bond, they lose, gain or share electrons to attain the electronic configuration
ns2np6 of the nearest noble gas. Nearly every main-group monoatomic ion has a filled
outer level of electrons (either two or eight), the same number as in the nearest noble gas.
Note! Octet rule states that during the formation of a chemical compound, each atom has
an octet (8) electrons in its highest occupied energy level by gaining, losing, or sharing
electrons.
Bonding: Bonding is the joining of at least two atoms to form a molecule or compound.
The electrons in the valence shell are the active portion of an atom during bonding.
A chemical bond is a mutual electrical attraction between the nuclei and valence
electrons of different atoms that binds the atoms together.
Why atoms combined to form compounds, clusters, molecules/ polyatomic ions?
 to attain octet (eight) rules
 to have lower energy (to be stable)
 to have net attractive force
e R e
A A
+ R +
A A
e R e

Where, A= attraction force, R= repulsion force

Chemical bond is formed when A>R:
Example: H + H→ H2 (g) but, He + He → no reaction since R>A.

Most atoms are chemically bonded to each other? As independent particles, they are at
relatively high potential energy. However, nature favors arrangements in which potential
energy is minimized. This means that most atoms are less stable existing by themselves
than when they are combined. By bonding with each other, atoms decrease in potential
energy, thereby creating more stable arrangements of matter.

HOW DO ATOMS COMBINED TO FORM COMPOUNDS / POLYATOMIC IONS ?

In chemical bonds, atoms are combined to form compounds, clusters, molecules/
polyatomic ions by:
 Transferring /losing of their valence electrons.
 Gaining of extra valence electrons
 Sharing of valence electrons
 Coordinating valence electrons

2.2. TYPES OF CHEMICAL B ONDING

When atoms bond, their valence electrons are redistributed in ways that make the atoms
more stable. The way in which the electrons are redistributed determines the type of
bonding. There are different types of chemical bonds that results by transfer or sharing of
electrons. These are:

There are three fundamental categories of bonds

 Ionic bond = Electropositive element + electronegative elements
 Na+Cl → NaCl
 Covalent bond = Electronegative element + electronegative elements
 F+F → F
2

 Metallic bond = Electropositive element + Electropositive element

 Li+Li →Li
2

 Elements can be classified as:

 Electropositive elements, whose atoms give up one or more electrons
 Electronegative element which take up electrons
Elements which neither tends to lose or gain electrons

2.2.1. Ionic Bond/Electrovalence bond

Ionic bond is a type of chemical bond that results from the electrostatic force of attraction
between two oppositely charged ions. It is formed due to transfer of electrons from one
atom (mostly metals) to another (mostly nonmetals) by losing and gaining of
electrons. The atom that loses electrons will form a cation and the atom that gains
electrons will form an anion. These oppositely charged ions come closer to each other
due to electrostatic force of attraction and thus form an ionic bond. In general, an ionic
bond is formed between two atoms that contain greater electronegativity difference and it
is between a metal atom and a nonmetal atom. Example: NaCl, LiF, MgCl2.

A. Favorable condition for ionic bond formation

 Low ionization energy of the metals

 High electro-affinity of the nonmetals
 High lattice energy
 High electro-negativity difference between a metal and a nonmetal atom

Note! Ionic compounds are usually formed when metal cations bond with non-metal
anions. The only common exception is ammonium ion which is not a metal, but it forms
ionic compounds

2.2.2. Covalent bond:

Covalent bond is formed between two atoms due to the sharing of electron pair(s). It is
formed between two atoms that contain the same or small difference in electronegativity.
Bonding between two atoms of the same element is completely covalent. In covalent
bond, single, double, triple and quadruple bonds are formed. There are different types of
covalent bonds like polar and non-polar and coordinate covalent bonds.
1. A polar-covalent bond is a covalent bond in which the bonded atoms have an
unequal attraction for the shared electrons.
Example: hydrogen-chlorine bond (H+-Clδ⁻).
 Polar covalent bonding occurs because one atom has a stronger affinity for electrons
than the other.
 When electrons are shared but shared unequally.
2. Non-polar covalent bond is a covalent bond in which the bonding electrons are
shared equally by the bonded atoms, resulting in a balanced distribution of
electrical charge. Example; hydrogen-hydrogen bond (H: H), Cl2. The electrons are
shared equally.
3. Coordinate Covalent bond: A coordinate covalent bond (also called a dative bond) is
formed when one atom donates both of the electrons to form a single covalent bond.
These electrons originate from the donor atom as an unshared pair.
Example:

H F H
F

F B + N H F B N H

H F H
F
Dative bond

2.2.2.1. Lewis structures

 The Lewis theory was the first explanation of a covalent bond in terms of
electrons that was generally accepted.
 He uses dot (.) to represent valance electron. Atoms continue to form bonds until
they have made up on octet of electrons. This is called the octet rule’.
 In forming compounds, atoms gain, lose, or share electrons to give a stable
electron configuration characterized by 8 valence electrons.
Hints on Lewis Dot Structures
 Octet rule‖ is the most useful guideline.
 Carbon forms 4 bonds.
 Hydrogen typically forms one bond to other atoms.
 When multiple bonds are formed, they are usually between C, N, O / S.
 Nonmetals can form single, double, and triple bonds, but not quadruple
bonds.
 Always account for single bonds and lone pairs before forming multiple
bonds.
 Look for resonance structures.
Steps in writing Lewis structure
 1) Calculate the total number of valance electrons for molecules by adding the
number of valance electrons for each species that participate in the formation of
molecules/compounds.
2) Determine the number of electron pairs by dividing the total number of valance
electrons by two to get number of paired electrons.
3) Predict the central atoms: Place the least electronegative element at the centre,
except for H which is always an outer atom.
4) Distribute the remaining electrons to the attached atoms till they acquired octet
rules.
5) If there are still remaining electrons distribute them as lone pairs to the central
atom.
6) If octet rules not obeyed for the central atom, it forms multiple bonds.
Let us determine the Lewis structures of CHO2−, NO+, and OF2 as examples in
following this procedure:
1. Determine the total number of valence (outer shell) electrons in the molecule or ion.
For a molecule, we add the number of valence electrons on each atom in the molecule:
SiH4
Si: 4 valences electrons/atom×1 atom = 4
+H: 1 valence electron/atom×4atoms = 4
= 8 valence electrons
For a negative ion, such as SiH4, CHO2−, we add the number of valence electrons on the
atoms to the number of negative charges on the ion (one electron is gained for each single
negative charge): CHO2−
C: 4 valences electrons/atom×1 atom=4

H: 1 valence electron/atom×1 atom=1

O: 6 valence electrons/atom×2 atoms=12

+1 additional electron
=18 valence electrons
For a positive ion, such as NO+, we add the number of valence electrons on the atoms in
the ion and then subtract the number of positive charges on the ion (one electron is lost
for each single positive charge) from the total number of valence electrons:
NO+
N: 5 valences electrons/atom×1 atom=5
O: 6 valences electron/atom×1 atom=6
+ −1 electron (positive charge = -1
=10 valence electrons
Since OF2 is a neutral molecule, we simply add the number of valence electrons:

OF2
O: 6 valences electrons/atom×1 atom=6
F: 7 valence electrons/atom×2 atoms = 14
= 20 valence electrons
2. Draw a skeleton structure of the molecule or ion, arranging the atoms around a
central atom and connecting each atom to the central atom with a single (one
electron pair) bond. (Note that we denote ions with brackets around the structure,
indicating the charge outside the brackets:)

When several arrangements of atoms are possible, as for CHO2−, we must use
experimental evidence to choose the correct one. In general, the less electronegative
elements are more likely to be central atoms. In CHO2−, the less electronegative carbon
atom occupies the central position with the oxygen and hydrogen atoms surrounding it.
Other examples include P in POCl3, S in SO2, and Cl in ClO4−. An exception is that
hydrogen is almost never a central atom. As the most electronegative element, fluorine
also cannot be a central atom.
3. Distribute the remaining electrons as lone pairs on the terminal atoms (except
hydrogen) to complete their valence shells with an octet of electrons. There are no
remaining electrons on SiH4, so it is unchanged:
4. Place all remaining electrons on the central atom.
1. For SiH4, CHO2−, and NO+, there are no remaining electrons; we already placed all of
the electrons determined in Step 1.
2. For OF2, we had 16 electrons remaining in Step 3, and we placed 12, leaving 4 to be
placed on the central atom:

5. Rearrange the electrons of the outer atoms to make multiple bonds with the central
atom in order to obtain octets wherever possible.
SiH4: Si already has an octet, so nothing needs to be done.
CHO2−: We have distributed the valence electrons as lone pairs on the oxygen
atoms, but the carbon atom lacks an octet:

NO+: For this ion, we added eight valence electrons, but neither atom has an octet. We
cannot add any more electrons since we have already used the total that we found in Step
1, so we must move electrons to form a multiple bond:

This still does not produce an octet, so we must move another pair, forming a triple bond:

In OF2, each atom has an octet as drawn, so nothing changes

2.2.2.2. Resonance Structures

We have assumed up to this point that there is one correct Lewis structure. Now we turn
our attention to Resonance structures: the structure of the same relative placement of
atoms but different locations of bonding and non-bonding electron pairs. Systems which
have more than one Lewis structure are Resonance Structures. Resonance is
delocalization of electrons through double bond. NO3- is a classic example of resonance:

2.2.2.3. Formal Charge: The hypothetical charge on an atom in the molecules. It helps
to check whether the Lewis structure is stable or not. The smallest formal charge is the
more stable in Lewis structure regardless numerical sign. A more negative formal charge
should reside on an atom with a larger EN value.
1 
FC  # Ve. -  # BE # LPE 
2 
FC = formal charge; #Ve. = Number of valence electrons

#BE = bonding electrons; #LPE = lone pair electrons

Example: NCO- has three possible resonance forms.

Forms B and C have negative formal charges on N and O. These forms are more
important than Form A. Form C has a negative charge on O which is more
electronegative than N. Therefore, Form C contributes the most to the resonance hybrid
and stable.

Limitations: Lewis theory was good for s-block and p-block elements but not for d-
block elements. Except in simple cases, Lewis structure can predict neither the 3D shape
of the species (bond angles) nor the relative internuclear distances (bond lengths). Lewis
theory cannot write one correct structure for many molecules where resonance is
important. Lewis theory often does not predict the correct magnetic behavior of
molecules. Oxygen, O2, is paramagnetic, though the Lewis structure predicts it is
diamagnetic.
2.2.2.4. Exceptions to the Octet Rule:
 Many covalent molecules have central atoms that do not have eight electrons in
their Lewis structures. These molecules fall into three categories:
 Odd-electron molecules have an odd number of valence electrons, and
therefore have an unpaired electron.
 Electron-deficient molecules have a central atom that has fewer electrons
than needed for a noble gas configuration.
 Hypervalent molecules have a central atom that has more electrons than
needed for a noble gas configuration.
a) Electron-Deficient Molecules: gaseous molecules containing either Be or B as the
central atom; have fewer than 8 electrons around the Be or B (4 e- around Be and 6
e- around B) (BF3).
b) Odd-Electron Molecules: have an odd number of valence electrons; examples
include free radicals, which contain a lone (unpaired) electron and are
paramagnetic (use formal charges to locate the lone electron) (NO2). Example: CH3,
OH, H, NO2 etc.
c) Expanded Valence Shells: for molecules that have more than 8 electrons around
the central atom; use empty outer d orbitals; occurs only with a central atom from
Period 3 or higher (SF6, PCl5). E.g: PCl5, SF6, H2SO4, H3PO4
2.3 Valence Shell Electron Pair Repulsion Theory and Molecular Geometry
2.2.2.5. Molecular structure and polarity

VSEPR Theory
Valence shell electron-pair repulsion theory (VSEPR theory) enables us to predict
the molecular structure, including approximate bond angles around a central atom,
of a molecule from an examination of the number of bonds and lone electron pairs
in its Lewis structure.
The VSEPR model assumes that electron pairs in the valence shell of a central
atom will adopt an arrangement that minimizes repulsions between these electron
pairs by maximizing the distance between them.
The electrons in the valence shell of a central atom form either bonding pairs of
electrons, located primarily between bonded atoms, or lone pairs.
The electrostatic repulsion of these electrons is reduced when the various regions
of high electron density assume positions as far from each other as possible.
VSEPR theory predicts the arrangement of electron pairs around each central
atom and, usually, the correct arrangement of atoms in a molecule.
We should understand, however, that the theory only considers electron-pair
repulsions.
Other interactions, such as nuclear-nuclear repulsions and nuclear-electron
attractions, are also involved in the final arrangement that atoms adopt in a
particular molecular structure.

2.2.2.6. Electron-pair Geometry versus Molecular Structure

 It is important to note that electron-pair geometry around a central atom is not

the same thing as its molecular structure.
 Molecular structure describes the location of the atoms, not the electrons.
 We differentiate between these two situations by naming the geometry that
includes all electron pairs the electron-pair geometry.
 The structure that includes only the placement of the atoms in the molecule is
called the molecular structure.
  The electron-pair geometries will be the same as the molecular structures when
there are no lone electron pairs around the central atom, but they will be
different when there are lone pairs present on the central atom.
 For example, the methane molecule, CH4, which is the major component of
natural gas, has four bonding pairs of electrons around the central carbon atom;
the electron-pair geometry is tetrahedral, as is the molecular structure
 On the other hand, the ammonia molecule, NH3, also has four electron pairs
associated with the nitrogen atom, and thus has tetrahedral electron-pair
geometry.
VSEPR theory predicts these distortions by establishing an order of repulsions
and an order of the amount of space occupied by different kinds of electron pairs.
The order of electron-pair repulsions from greatest to least repulsion is:
Lone pair-lone pair>lone pair-bonding pair>bonding pair-bonding pair
 This order of repulsions determines the amount of space occupied by different
regions of electrons.
 A lone pair of electrons occupies a larger region of space than the electrons in a
triple bond; in turn, electrons in a triple bond occupy more space than those in a
double bond, and so on.
 The order of sizes from largest to smallest is:
Lone pair>triple bond>double bond>single bond

Predicting Electron Pair Geometry and Molecular Structure

The following procedure uses VSEPR theory to determine the electron pair geometries
and the molecular structures:
1) Write the Lewis structure of the molecule or polyatomic ion.
2) Count the number of regions of electron density (lone pairs and bonds) around the
central atom. A single, double, or triple bond counts as one region of electron
density.
3) Identify the electron-pair geometry based on the number of regions of electron
density: linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral.
4) Use the number of lone pairs to determine the molecular structure. If more than
one arrangement of lone pairs and chemical bonds is possible, choose the one that
 will minimize repulsions, remembering that lone pairs occupy more space than
multiple bonds, which occupy more space than single bonds. In trigonal
bipyramidal arrangements, repulsion is minimized when every lone pair is in an
equatorial position. In an octahedral arrangement with two lone pairs, repulsion is
minimized when the lone pairs are on opposite sides of the central atom.
The following examples illustrate the use of VSEPR theory to predict the molecular
structure of molecules or ions that have no lone pairs of electrons. In this case, the
molecular structure is identical to the electron pair geometry.

Example 1: Predicting Electron-pair Geometry and Molecular Structure: CO2 and BCl3
Predict the electron-pair geometry and molecular structure for each of the following:
(a) Carbon dioxide, CO2, a molecule produced by the combustion of fossil fuels
(b) Boron trichloride, BCl3, an important industrial chemical
Solution
(a) We write the Lewis structure of CO2 as:

This shows us two regions of high electron density around the carbon atom-each double
bond counts as one region, and there are no lone pairs on the carbon atom. Using VSEPR
theory, we predict that the two regions of electron density arrange themselves on opposite
sides of the central atom with a bond angle of 180°. The electron-pair geometry and
molecular structure are identical, and CO2 molecules are linear.
(b) We write the Lewis structure of BCl3 as:

The electron-pair geometry and molecular structure of BCl3 are both trigonal planar. Note
that the VSEPR geometry indicates the correct bond angles (120°), unlike the Lewis
structure shown above.
2.4 Bonding Theories
2.2.2.7. Molecular structure and dipole moment
 Polar covalent bonds connect two atoms with differing electronegativity, leaving
one atom with a partial positive charge (δ+) and the other atom with a partial
 negative charge (δ–), as the electrons are pulled toward the more electronegative
atom.
 This separation of charge gives rise to a bond dipole moment.
 The magnitude of a bond dipole moment is represented by the Greek letter mu
(µ) and is given by the formula shown here, where Q is the magnitude of the
partial charges (determined by the electronegativity difference) and r is the
distance between the charges:

 This bond moment can be represented as a vector, a quantity having both

direction and magnitude
 Dipole vectors are shown as arrows pointing along the bond from the less
electronegative atom toward the more electronegative atom.
 A small plus sign is drawn on the less electronegative end to indicate the partially
positive end of the bond.
 The length of the arrow is proportional to the magnitude of the electronegativity
difference between the two atoms.
To summarize, to be polar a molecule must:
 Contain at least one polar covalent bond.
 Have a molecular structure such that the sum of the vectors of each bond
dipole moment does not cancel.

The trans isomer has no dipole moment (µ= 0 D) because the C – Cl bond polarities
balance each other. In contrast, the cis-isomer is polar (µ = 1.90 D) because the bond
dipoles partially reinforce each other, with the molecular dipole pointing between the Cl
atoms.
 2.2.2.8. Intermolecular Forces in Covalent Compounds

Dipole-Dipole forces: When polar molecules are brought near one another, their partial
charges act as tiny electric fields that orient them and give rise to dipole-dipole forces;
the partially positive end of one molecule attracts the partially negative end of another
For compounds of approximately the same mass and size, greater is the dipole moment,
greater is the dipole-dipole forces between their molecules, so greater is the energy
required to separate their particles.

For instance, both methyl chloride (CH3Cl) and ethanol (CH3CHO) have comparable
mass and size, but CH3Cl has a smaller dipole moment than CH3CHO; therefore, less
energy is needed to overcome the dipole-dipole forces between its molecules and it boils
at a lower temperature. Dipole-dipole forces give polar cis-1,2-dichloroethene a higher
boiling point than nonpolar trans-1, 2- dichloroethene.
Hydrogen Bonding: Hydrogen bonding is a special type of dipole-dipole forces that arise
between molecules that have a hydrogen atom bound to a smaller sized, most
electronegative atoms. These atoms are: Fluorine, Oxygen and Nitrogen

The partially positive (δ+) H of one molecule is attracted to the partially negative (δ –)
lone pair on the F, O or N of the same or another molecule. As a result, hydrogen bonds
result.

Dispersion or London Forces: Intermolecular forces in which non-polar atoms or

molecules interact by inducing dipoles in each other are known as dispersion or London
Forces. These classes of intermolecular forces cause substances like CO2, Cl2, noble
gases, etc., to condense and solidify.
An attractive force must be acting between these non-polar molecules and atoms, or they
would remain gaseous under any conditions. In fact, bond dipoles exert some weak
attraction but the intermolecular forces are mainly responsible for the condensed state of
non-polar substance. These forces are known as London forces or dispersion forces,
named after Fritz London, the physicist who explained the quantum mechanical basis of
the attractive forces. Dispersion forces are very weak and are caused by a sudden shift of
electron density to one side of the nucleus than the other

2.2.3. Metallic Bond

It is the attraction between metal atoms in a metallic crystal. It is formed between
electropositive metal atoms of same or different elements. The valence electrons of pure
metals are not strongly associated with particular atoms. This is a function of their low
ionization energy. Electrons in metals are said to be delocalized (not found in one
specific region, such as between two particular atoms). Since they are not confined to a
specific area, electrons act like a flowing “sea”, moving about the positively charged
cores of the metal atoms. Delocalization can be used to explain conductivity,
malleability, and ductility. In general, the greater the number of electrons per atom that
participate in metallic bond the stronger the metallic bond. E.g. the metal atoms Na, Cu,
Ag, and Fe etc. are bound to each other in their crystals by metallic bonds.
2.3. CHEMICAL BONDING THEORIES
2.3.1 Valence Bond Theory (VBT)
The basic principle of valence bond theory is that a covalent bond forms when orbitals of
two atoms overlap and the overlap region, which is between the nuclei, is occupied by a
pair of electrons


In the HCl bond, for example, the 1s orbital of hydrogen overlaps the half-filled 3p
orbital of Cl along the axis of that orbital.
Similarly, in the Cl–Cl bond of Cl2, the two 3p orbitals interact end-to-end, that is along
the orbital axes, to attain maximum overlap
Overlap of Atomic Orbitals (Sigma and Pi Bonds): Various type of atomic orbital
overlap leads to covalent bond formation. Three simple basic ones are s-s, s-p and p-p
1. s-s overlap in which half-filled s orbitals overlap,
2. s-p overlap where half-filled s orbital of one atom overlaps with one of the p orbital
having one electron only and
3. p-p overlap in which two half-filled p orbitals overlap.

Figure 2. Atomic orbital overlap.

(A) s-s overlap, (B) s-p overlap, (C) p-p overlap along the orbital axis and
(D) p-p sideways overlap.
Because of the directional nature of p orbital, the overlap may take place in two ways: (i)
the half-filled p orbitals along the line joining the two nuclei. This is called as head on,
end-to-end, end on or linear overlap and (ii) the half-filled p orbitals overlap along the
line perpendicular to the nuclear axis. This overlapping of p orbitals is parallel to each
other, hence is called lateral or sideways overlapping. These two modes give rise to the
two types of covalent bonds, namely, sigma (σ) and pi (π) bond respectively. Valence
bond theory is used to describe the two types here, but they are essential features of
molecular orbital theory as well. s - s and s - p will always overlap along the nuclear axis,
hence results only in sigma bonds. A side - to - side or parallel overlap forms another
type of covalent bond called a pi (π) bond. It has two regions of electron density. One
above and one below the sigma bond axis. One π -bond holds two electrons that occupy
both regions of the bond. A double-bond always consists of one σand one πbond. The
extent of overlap influences bond strength, however, many factors, such as lone - pair
repulsions, bond polarities, and other electrostatic contributions affect overlap and the
relative strength of σ and π bonds between other pairs of atoms.
Hybridization of Orbitals

Hybridization
 The combination of atomic orbital is called hybridisation
 The word ‘hybridization’ means ‘mixing
 Hybrid orbitals may be formed by mixing the characters of atomic orbitals
that are close in energy
 The number of bond formed by an atom is the same as the number of
unpaired electron that combined in ground state.

sp hybrid orbitals
Combination (mixing) of one s and one p orbital is called sp hybridization and the
resultant orbitals are called sp hybrid orbitals.
VB theory explains this by proposing that mixing two non-equivalent orbitals of a central
atom one s and one p, gives rise to two equivalent sp hybrid orbitals that lie 180° apart.
Example: In the triatomic molecule BeCl2 the 2s and one of the 2p orbitals of the Be atom
are hybridized into sp hybrid orbitals. The remaining two 2p orbitals are left hybridized
and unoccupied in the orbital diagram.
sp2 hybrid orbitals
example in BCl3, the boron atom has four orbitals but only three electrons in its valence
shell. In most boron compounds the hybridization scheme combines one 2s and two 2p
orbitals into three sp2 hybrid orbitals.

The valence bond method predicts that BCl3 is a trigonal planar

molecule with 120° Cl–B–Cl bond angle
sp3 hybridization
Carbon, the central atom in a molecule of methane CH4 has only two unpaired electrons
in the ground state. rule. Carbon in its excited state can form four bonds. The one 2s and
three 2p-orbitals of carbon are mixed to produce four new orbitals that are equivalent to
each other in energy and in shape, but pointing in different directions,
The four sp3 hybrid orbitals point to the corners of a regular tetrahedron, and make a
bond angle of 109.50° in CH4
We might also expect to use sp3 hybridization not only for structures of the type AX4
type (as in CH4), but also for AX3E type (as in NH3) and AX2E2 type (as in H2O).

sp3d hybridization
PCl5, the central atom phosphorus has only
three unpaired electrons in its ground state. Electrons must be unpaired to provide the
correct number of unpaired electrons for bond formation

sp3d2 Hybridization
Another structure featuring an expanded octet is SF6. Here, six hybrid orbitals are
required in order to describe bonding.
In hybridization schemes, one hybrid orbital is produced for every single atomic orbital
involved. In a molecule, each of the hybrid orbitals of the central atom acquires an
electron pair, either a bond-pair or a lone-pair. And the hybrid orbitals have the same
orientation as the electron-set arrangement predicted by VSEPR theory.
Lewis structure of ethene (ethylene) is depicted as

All the C–H bonds in C2H4 are formed by the overlap of sp2 hybrid orbitals of the C
atoms with 1s orbitals of the H atom, A double covalent bond consists of one σ and one π
bond.
Consider the Ethyne (acetylene) molecule, C2H2.

The molecule is linear, with 180° H–C–C bond angles

in C2H2, as in all triple bonds, one bond is a σ-bond and two are π- bonds.

2.3.2. Molecular Orbital Theory (MOT)

 Deals with combination of atomic orbitals
 Combination will be; - s—s, s—p, p—p, p—d, d—d combination
Bonding and Anti-Bonding Molecular Orbitals
In place of atomic orbitals of the separated atoms, the molecular orbitals for the united
atoms are obtained, and these are of two types. One type, a bonding molecular orbital,
places a high electron charge density in between the two nuclei and the other type, an
anti-bonding molecular orbital, places a high electron charge density away from the
region between the two nuclei. Electrons in bonding orbitals contribute to bond formation
and electrons in anti-bonding orbitals detract from bond formation.
In place of atomic orbitals of the separated atoms, the molecular orbitals for the united
atoms are obtained, and these are of two types. One type, a bonding molecular orbital,
places a high electron charge density in between the two nuclei and the other type, an
anti-bonding molecular orbital, places a high electron charge density away from the
region between the two nuclei. Electrons in bonding orbitals contribute to bond formation
and electrons in anti-bonding orbitals detract from bond formation.
When two atomic orbitals overlap end-to-end, they form two sigma-molecular orbitals
(MOs). Consider the H2 molecule, which has two H atoms and therefore two 1s
AOs. The two 1s atomic orbitals combine (see Figure 3) to produce two σMOs,
which differ in energy and location. One of the σMOs is a bonding orbital, denoted
σ1s, the other is an anti-bonding orbital denoted σ*1s.
The relative energy levels of these two MOs are different. The σ1s MO has a lower
energy than the original 1s AOs, while the σ*1s MO has a higher energy.

Electron Configuration of Diatomic Molecules

Figure 3 Electron Configuration of Hydrogen Molecules

 Electron configuration in MO is given by:
б1s, б*1s, б2s, б*2s, б2p ∏2p ∏*2p б*2p
x , y y x

∏2p ∏*2p
z z

 This order show exception for lighter elements such as B, C & N (14 electron and
less) as:
б1s, б*1s, б2s, б*2s, б ∏2p б2p ∏*2p б*2p
y, x, y x
 ∏2p ∏*2p
z z
2 2
Eg. He molecule consists 4 electron б1s , б*1s
2
2 2 2
Li molecule consists 6 electron б1s , б*1s б2s
2
2 2 2 2 2 2 1
O molecule consists 16 es. б1s , б*1s , б2s , б*2s , б2p ∏2p ∏*2p б*2p
2 x, y , y, x
2 1
∏2p , ∏*2p
z z

O is paramagnetic in nature
2

Bond Order
The term bond order is used to indicate whether a covalent bond is single (bond order =
1), double (bond order = 2) or triple (bond order = 3).
• The existence of molecules determined by bond order which is given by:
BO= BMO- ABMO

Eg. In O , 10-6 = 2 , the molecule can exist

2

2
 A molecule exists only if the bond order is positive.
 If bond order is zero or negative, the molecule does not exist.
Magnetic Properties
A species with unpaired electrons exhibits paramagnetic property. The species is attracted
by an external magnetic field. A species in which all the electrons are paired, exhibits
diamagnetism. Such species are not attracted (and, in fact, are slightly repelled) by a
magnetic field.
Exercise 1
1. Use the molecular orbital theory and derive the electron configuration of the
following molecules. Identify those which exist and do not exist. a He2 b Be2 c
B2 d C2 e N2
2. Draw a molecular orbital energy diagrams for: a. C2– b. .C2 c. C2+
2.4. TYPES OF CRYSTALS
A crystal is a piece of a solid substance that has plane surface, sharp edges, and a regular
geometric shape A structural unit of a crystalline solid has a characteristic repetitive
pattern. The crystal types and their basic units are (i) ionic (electrostatic attraction of
ions), (ii) Molecular (electrostatic attraction of dipoles in molecules) (a) Polar (dipole-
dipole and Hbonding) and (b) Non-polar (London forces) (iii) Network (covalently
bonded atoms); and (iv) metallic (positive nuclei in electron sea).There are four important
classes of crystalline solids.
A. Ionic Crystals
The fundamental units of an ionic solids are positive and negative ions. As a result, the
inter-particle forces (ionic bonds) are much stronger than the van der Waals forces in
molecular solids. Ionic solids are table salt (NaCl), saltpeter (KNO3), washing soda
(Na2CO3.10H2O), and black board chalk (CaCO3).
B. Molecular Crystals
Various combinations of dipole-dipole, dispersion and hydrogen-bonding forces are
operative in molecular solids, which accounts for their wide range of physical properties.
Dispersion forces are the principal forces acting in non-polar substances, so melting
points generally increase with molar mass. Among polar molecules, dipole dipole forces
and where ever possible, hydrogen-bonding dominate. Nevertheless, intermolecular
forces are still relatively weak, so the melting points are much lower than ionic, metallic
and network covalent solids. The fundamental unit of a molecular solid is the molecule.
Such solids are common among organic compounds and simple inorganic compounds.
Molecular crystals are usually transparent, brittle, and break easily when stressed. They
are usually non-conductors of heat and electricity and usually have low melting points
molecular crystalline solids include sugar, dry ice (solid carbon dioxide), glucose and
aspirin.
C. Covalent Network Crystals
Exercise 2
 Match the substances from list A with the best suited characteristics given in
list B and justify your answer

Unit summary
 Ionic compounds are crystalline solids at room temperature. The fundamental units
of ionic solid are positive and negative ions. Crystalline ionic solids are usually
brittle and non-conductors of electricity, although molten crystals may be good
conductors. They usually have high melting and boiling points.
 Ionic compounds are nonvolatile.
 Ionic compounds are usually soluble in inorganic solvents (water is the most
common solvent for ionic compounds) but insoluble in organic solvents like
benzene, ethanol and carbon tetrachloride.
 Lewis symbols of representative elements are related to their location in the periodic
table. The lattice energy and enthalpy of formation of an ionic compound, together
with other atomic and molecular properties, can be related in a thermochemcial cycle
called Born-Haber cycle.
 A covalent bond is created by the sharing of an electron-pair between atoms.
 In a Lewis structure representing covalent bonds, electron-pairs are either bonding-
pairs or lone-pairs.
 In a covalent bond between atoms of different electronegativity, electrons are
displaced toward the atom with the higher electronegativity. In terms of
electronegativity differences chemical bonds vary over the range: non polar to polar
covalent to ionic.
 In some cases of covalent bonding, one atom appears to provide both electrons in the
bonding pair; the bond is known as coordinate-covalent bond.
 Bonded atoms may share more than one pair of electrons between them, giving rise
to multiple covalent bonding.
 In the phenomenon of resonance, two or more Lewis structures have the same
skeletal structure but different bonding arrangements. The best description of the
resonance structure (resonance hybrid) is obtained by combining plausible structures
(contributing structures).
 Exceptions to the octet rule are found in odd-electron molecules and molecular
fragments called free radicals. A few structures appear to have too few electrons to
complete all the octets. Some structures appear to have too many. In the latter case, a
central atom may employ an "expanded" octet with five or six electron-pairs.
 Valance shell electron pair repulsion theory (VSEPR) predicts the geometrical
structures of molecules and polyatomic ions based on mutual repulsions among
valance shell electron groups.
 Multiple bonds, whether they consist of four electrons (double bond) or six
electrons (triple bond), are treated as one electron set.
 The separation of the centers of positive and negative charge in a polar covalent
bond creates a bond moment. Whether a molecule as a whole is polar, that is,
whether there is a resultant dipole moment, is established by bond moments and
molecular geometry.
 A symmetrical distribution of identical bond moments about a central atom results in
a cancellation of all bond moments, with the result that the molecule is nonpolar, as
in the case of CCl4.
 In the valence bond method (VB) a covalent bond is viewed as the overlap of
atomic orbitals of the bonded atoms in a region between the atomic nuclei.
 Molecular geometry is determined by the spatial orientations of the atomic orbitals
involved in bonding.
 The VB method often requires that bonding atomic orbitals be hybridized in order to
rationalize known structures of molecules. A hybridized orbital is some combination
of s, p and d orbitals, such as sp, sp2, sp3, sp3d and sp3d2. The geometric distribution
of hybridized orbitals in the valence bond method is the same as the electron set
geometry proposed by VSEPR theory.
 Hybrid orbitals overlap in the usual way (end-to-end) and form σ bonds.
Unhybridized p orbitals overlap in a side-by-side manner and give rise to π bonds. A
double bond consists of one σ bond and one σ bond; a triple bond, one σ bond and
two π bonds.
 Acceptable solutions to wave equations written for the electrons in a molecule are
called molecular orbitals (MO). The two main types of MOs are bonding molecular
orbitals, which concentrate electron charge density between atoms or just above and
below the imaginary line joining the two nuclei and antibonding molecular orbitals,
which concentrate electron charge densities away from the intermolecular bonding
region.
 Electrons can be assigned to molecular orbitals by the scheme similar to the aufbau
process. MO theory provides more satisfactory descriptions for certain structures
than does the VB method, for example, some odd - electrons species and the O2
molecule.
Review questions
1. Explain the formation of bonds in the following pairs of elements:
a. potassium and chlorine b. magnesium and oxygen and c. sodium and
oxygen.
2. Which of the following elements will form an ionic bond with chlorine and why?
Calcium, Carbon, Oxygen and Silicon
3. Why ionic bond is also known as electrovalent bond?
4. How many types of chemical bonding you are familiar with?
5. State and explain the formation of ionic, covalent and metallic bonds. Use diagrams
wherever required.
6. List four important characteristics of ionic compounds.
7. What observable properties can you use to distinguish one kind of bond from
another?
8. Write the formulas and names of the compounds formed from the following ionic
interactions: (use periodic table)
a. The 2+ ion and 1– ion is both isoelectronic with the atoms of a chemically
unreactive period 4 element.
b. The 2+ ion and the 2– ion are both isoelectronic with the period 3 noble
gas.
c. The ions formed are the largest and smallest ionisable atoms in period 2.
9. In each of the following ionic compounds identify the main group to which X
belongs:
a. XF2 b. MgX c. X2O3 d. Na2X
–1
10. For lithium, the enthalpy of sublimation is +161 kJ mol , and the first ionization
energy is +520 kJ mol–1. The dissociation energy of fluorine is +154 kJ mol–1, and the
electron affinity of fluorine is –328 kJ mol–1. The lattice energy of LiF is –1047 kJ
mol–1. Calculate the overall enthalpy change for the reaction?
11. Li(s) + ½F2(g) →LiF(s) ΔH° =?
12. The enthalpy of formation of caesium chloride is
Cs(s) + ½Cl2(g) →CsCl(s) ΔH° = – 44.28 kJ mol–1
The enthalpy of sublimation of caesium is
Cs(s) →Cs(g) ΔH° = + 77.6 kJ mol–1
Use these data, with other data from other sources, to calculate the lattice energy
of CsCl(s)
13. Using the following data:
Enthalpy of sublimation of Ca = +178.2 kJ mol–1
Enthalpy of dissociation of Cl2 = +243.4 kJ mol–1
Enthalpy of formation of CaCl2 = –795.8 kJ mol–1
 First and second Ionization energies for Ca are +590 kJ mol–1 and +1145 kJ mol–1
respectively.
The electron affinity of Cl = –348.7 kJ mol–1
Determine the lattice energy of CaCl2
14. Determine the total number of valence electrons for the following species:
a. CO2 b. SO24– c. NH4+ d. N2O4
15. Write a plausible Lewis structure of:
16. a. nitrogen trichloride, NCl3 b. chlorate ion, ClO3-
17. c. phosphonium ion, PH4+ d. phosgene, COCl2
2–
18. Draw a Lewis structure for CO3 , SF4 and HCOOH (formic acid).
19. Which of the following atoms cannot serve as a central atom in a Lewis structure O,
He, F, H, P? Explain.
20. Write a plausible Lewis structure for carbonyl sulphide, COS.
21. In which of the following substances do hydrogen bonds occur? Explain with the help
of diagrams.
a. CH4 b. CH3CH2OH

22. Identify the dominant intermolecular force that is present in each of the following
substance and select the substance with the higher boiling point in each pair:
a. CH3OH or CH3CH2OH c. MgCl2 or PCl3
b. Hexane or cyclohexane d. CH3NH2 or CH3F
23. Which type of intermolecular force is dominant in the following substances?
a. ICl b. H2O
c. F2 d. HBr
24. Compare intermolecular forces with that of intramolecular bonding.
25. How do carbon and silicon differ with regard to the types of orbitals available for
hybridization? Explain
26. Are these statements true or false? Correct any that is false.
a. Two σ bonds comprise a double bond.
 b. A triple-bond consists of one π bond and two σ bonds.
c. Bonds formed from atomic s orbitals are always σ bonds.
d. A π -bond consists of two pairs of electrons
e. End-to-end overlap results in a bond with electron density above and below
the bond axis.
27. Name the three types of hybrid orbitals that may be formed by an atom with only s
and p orbitals in its valence shell. Draw the shapes of the hybrid orbitals so produced.
28. Describe a hybridization scheme for the central atom and molecular geometry of the
triiodide ion, I3–.
29. Describe a hybridization scheme for the central atom S and the molecule geometry of
a SO3 and b SO2
30. Describe a hybridization scheme for the central atom and the molecular geometry of
CO2.
31. Discuss the bonding in nitrate ion, predict the ideal bond angles, bond length, shape
of the ion, the number of sigma and pi bonds.
2-

32. Use the MOT to explain bond order and magnetic properties of: A. O2 B. O2
C. O2+ D. NO E. CO F. C
2

33. What is the bond order for CN–, CN and CN+?

34. Which homonuclear diatomic molecules of the second period elements,
35. besides O2, should be paramagnetic?
Multiple Choice Questions
1. "Two atoms each provide two electrons that are shared by the two atoms" This is a
description of a: a. triple covalent bond c. double covalent bond
b. coordinate covalent bond d. single covalent bond
2. A non-metal usually forms two covalent bonds in its compounds. How many
electrons will it have in its valence shell? a. 2 b. 6 c. 4 d. 8
3. The noble gases do not usually form chemical compounds because:
(a) they have very stable nuclei.
(b) the bonds between their atoms are very strong.
(c) they already have paired valence shell electrons.
(d) they are not polar.
4. Which of the following compounds contain both ionic and covalent bonds?
a. CO2 b. Cl2O c. Na2CO3 d. BaCl2
5. Which of the following molecules would you expect to be non-planar?
a. CH4 b. BCl3 c. XeF4 d. HCHO
6. If a molecule has a trigonal pyramidal shape, how many non-bonding pairs of
electrons are present in the valence shell of the central atom?
a. 1 b. 2 c. 3 d. 4
7. Of the following, the most polar bond is: a. P – Cl b. Si – F c. S – O d. C – N
8. Carbon and chlorine form a series of compounds: CH3Cl, CH2Cl2, CHCl3, CCl4.
Which of these will be polar?
a. CCl4 only b. CH3Cl and CHCl3 only
c. CH3Cl, CH2Cl2 and CHCl3 only d.CH3Cl, CH2Cl2, CHCl3 and
CCl4
9. The carbon atoms in ethane (C2H6), ethene (C2H4) and ethyne (C2H2) provide,
respectively, examples of the three common types of hybridization corresponding to:
a. sp, sp2, sp3 b. sp3, sp2, sp c. sp, sp3, sp2 d. sp3, sp, sp2
10. Which of the following correctly describes a π bond?
a. It is formed by the end-to-end interaction of p-orbitals and has a low
electron density on the internuclear axis.
b. It is formed by the parallel interaction of p-orbitals and has a high electron
density just above and below the internuclear axis.
c. It is formed by the interaction of s-orbitals and has a low electron density
on the internuclear axis.
d. It is formed by the interaction of s-orbitals and has a high electron density
on the internuclear axis.
11. Which of the following species cannot be adequately described by a single Lewis
structure?
a. OH– b. C2H2 c. NH4+ d. HCO3–
12. In which of the following compounds would inter-molecular hydrogen-bonding
occur?
a. HCHO b. PH3 c. CH3OH d. COCl2
13. Which of the following molecules would you expect to have the highest boiling
point?
a. CH3–CH2–CH2–CH2–CH2–CH3
b. CH3–CH(CH3)–CH2–CH2–CH3
c. CH3–CH2–CH(CH3)–CH2–CH3
d. CH3–C(CH3)2–CH2–CH3
14. Which of the following molecules possesses the strongest forces between the
molecules?
a. N2 b. CH3–CH3 c. H2CO d.
CH4
15. Which one of the following usually produces the weakest interaction between
particles of similar molar mass?
a. Van der Waals forces ……… …. c. Covalent bonds
b. Dipole-dipole forces …………. d. Hydrogen-bonding
16. The hydrogen bonding is not significant in:
a. DNA b. protein c. ice d. poleythene
17. A gas is likely to:
a. have its atoms held together by metallic bonds
b. have a giant covalent structure.
c. be a compound of a metal.
d. have a molecular covalent structure.
18. If an element in group IIA of the periodic table formed a compound with an element
in group VIIA of the periodic table, the compound formed is likely to:
a. be a crystalline solid c. have a low boiling point
b. dissolve in non-polar solvents d. conduct electricity in the solid
state
19. Which of the following substances would you expect to have the lowest boiling
point? a. CsCl b. SrSO4 c. Sc2O3 d. AsCl3
20. Which of the following would not conduct an electric current?
a. solid mercury b. Aqueous sodium chloride
c. liquid sodium chloride d. solid sodium chloride
 Part II: Answer the following questions:
21. Name the following compounds:
a. KBr b. Al I3 c. Li3N d. BeO e. BaS
22. Predict the shape for the following molecules:
a. SiF4 b. PCl3 c. H2S d. NF3 e. CCl4 f. PCl5
23. Draw Lewis structures and predict the shapes, giving approximate bond angles
for: CH3+, CH3–, .CH3 (methyl radical), HCN, ICl2, NH2 –, NO2, ClO2.
24. Classify the given molecules as polar or non-polar:
a. SiF4 b. NF3 c. PCl3 d. H2S e. CCl4 f. CO2
25. Which atom (if any), in each of the following bonds would you expect to carry a
partial negative charge?
a. H–N b. S–S c. O–P d. C–F e. B–O f. I–F
26. Draw the Lewis structure for the given molecules and state whether the molecule is
polar or non-polar:
a. PF4– b. ICl4– c. N2F2 (2 forms possible)
27. Boron trifluoride can react with a fluoride ion to give the tetrahedral BF4–ion.
What type of hybridization would you expect the boron to have in BF4–?
28. Carbon and oxygen can bond either by a single bond (as in CH3–OH), a double bond
(as in O=C=O), or a triple bond (as in HC≡CH).
a. Describe these three types of bonds in terms of σ-bonds and π-bonds.
b. How would you expect the length of the carbon-oxygen bond to vary in
the three given examples?
29. Explain the following in terms of the intermolecular forces that exist in them:
a. Water is a liquid at room temperature and atmospheric pressure, but
hydrogen sulphide is a gas.
b. At room temperature and atmospheric pressure chlorine is a gas, bromine
is a liquid and iodine is a solid.
c. Pentan-1-ol boils at 137°C, whereas pentan-3-ol boils at 116°C.
d. The boiling point of sulphur dioxide is 24°C higher than that of chlorine.
30. Of the two possible valence-bond formulas for formic acid (HCO2H), which one is
the more reasonable structure?
31. Draw the correct valence-bond formula for hydrazine (N2H4).
32. Draw the resonance hybrid of SO2-.
33. Based on the VB approach, predict the shape of a ClO3- ion.
34. The molecular orbital description of NO is similar to that of O2 except that the
energies of the oxygen atomic orbitals are slightly lower than those of nitrogen
atomic orbitals and NO molecule contains different number electrons than O2.
35. Draw an energy diagram to show how the atomic orbitals of N and O combine to
form molecular orbitals of NO.
36. How many unpaired electrons are present in the NO molecule?
37. What is the bond order in NO molecule?
38. Is NO paramagnetic or diamagnetic?
39. List the following in order of increasing O – O bond length: O2, O2+, O2–
40. Show the direction of the bond moment (dipole moment) in each of the following
bonds: N – F, N – H, N – Si, N–O
Chapter 3

Physical states of matter

3.1 Introduction

Matter is defined as anything that occupies space and has mass. It can exist in the form of
gas, liquid and solid. The simplest example is the water we use in our daily life. The three
physical states of water are:

• Steam, water in the form of gas.

• Water, in the form of liquid.

• Ice, water in the form of solid.

The physical state of a given sample of matter depends on the temperature and pressure.
Changing these conditions or variables may change the behaviour of the substances as
solids, liquids, gases or plasma

 A solid is rigid and possesses a definite shape.

 A liquid flows and takes the shape of its container, except that it forms a flat or
slightly curved upper surface when acted upon by gravity. Both liquid and solid
samples have volumes that are very nearly independent of pressure.
 A gas takes both the shape and volume of its container.  A fourth state of matter,
plasma occurs naturally in the interiors of stars.
 Plasma is a gaseous state of matter that contains appreciable numbers of
electrically charged particles. Example of plasma state: Lightning, Comet tail,
Solar wind, Stars (including the Sun), Interstellar gas clouds, Welding arcs,
fluorescent lights, Static electricity, Rocket exhaust etc.

3.2 The kinetic theory of matter

 The states of matter in which substances are chemically the same but
physically different are explained by the kinetic theory of matter
 gives an explanation of the nature of the motion and the heat energy
  according to the theory, every substance consists of a very large number of
very small particles called ions, atoms and molecules
 The motion of the particles increases with a rise in temperature
 kinetic theory of matter is based on:
 All matter is composed of particles which are in constant motion.
 The particles possess kinetic energy and potential energy.
 The difference between the three states of matter is due to their energy
contents and the motion of the particles.

3.3 Properties of Matter

properties of gases can be summarized.

 Gases have no definite shape and definite volume.

 Gases can be easily compressed.
 Gases have low densities compared with liquids and solids.
 Gases exert pressure in all directions.
 Gases easily flow and diffuse through one another.
Properties of Liquids

 Have a definite volume, but have no definite shape.

 Have higher densities than gases.
 Are slightly compressible.
 Are fluids.

Properties of Solids

 Solids have a definite shape and a definite volume.

 Solids generally have higher densities than gases and liquids.
 Solids are extremely difficult to compress.
 Solids are not fluids.
3.4 The gaseous state
3.4.1 The kinetic molecular theory of gases
 The particles in an ideal gas are very widely spaced and constant random motion.
  The pressure of a gas is the result of continuous collisions between the particles
and the walls of their container

Assumptions of the kinetic molecular theory of gases

1. The particles are in a state of constant, continuous, rapid, random motion and,
therefore, possess kinetic energy. The motion is constantly interrupted by collisions
with molecules or with the container. The pressure of a gas is the effect of these
molecular impacts.

2. The volume of the particles is negligible compared to the total volume of the gas.
Gases are composed of separate, tiny invisible particles called molecules. Since
these molecules are so far apart, the total volume of the molecules is extremely
small compared with the total volume of the gas. Therefore, under ordinary
conditions, the gas consists chiefly of empty space. This assumption explains why
gases are so easily compressed and why they can mix so readily.

3. The attractive forces between the particles are negligible. There are no forces of
attraction or repulsion between gas particles.

4. The average kinetic energy of gas particles depends on the temperature of the gas.
At any particular moment, the molecules in a gas have different velocities. The
mathematical formula for kinetic energy is K.E. = ½ mν2, where m is mass and ν is
velocity of gas molecules. Because the molecules have different velocities, they
have different kinetic energies. However, it is assumed that the average kinetic
energy of the molecules is directly proportional to the absolute (Kelvin)
temperature of the gas.
3.4.2 The Gas Laws

Express mathematical relationships between the volume, temperature, pressure, and

quantity of a gas.
Pressure: it is the force applied per unit area.

Pressure =
Force
Area
  Can be expressed in atmosphere, Pascal, torr, millimetre of mercury. Its SI unit is
Pascal (Pa), and defined as one Newton per square metre. 1 Pa = 1 N/m2 and

1 atm = 760 mmHg = 76 cmHg = 760 torr = 101325 Pa = 101.325 kPa

Volume: space taken up by a body. ItsSI unit is the cubic metre (m3) and also expressed
in cubic centimetre (cm3) and cubic decimetre (dm3). Other common units of volume are
millilitre (mL) and litre (L).

1 cm3 = (1×10–2 m)3 = 1×10–6 m3

1 dm3 = (1×10–1 m)3 = 1×10–3 m3 = 1 L

Temperature: degree of hotness or coldness of a body. Degree Fahrenheit (°F), degree

Celsius (°C) and Kelvin (K) are temperature scales commonly used. Necessary inter-
conversions are:

K = °C + 273

°F = 59 C° + 32
1. Boyle’s Law
 analyze the relationship between the volume and pressure of a fixed amount of a
gas at constant temperature.
 law states that the volume of a fixed mass of gas is inversely proportional to the
pressure at a constant temperature. mathematically given as
1
V α (at constant T and n)
𝑝

From which follows,

V = k/p or PV = k where k is a constant at a specific temperature for a given

sample of gas.

If P1 and V1 represent the initial conditions; and P2 and V2 represent the new or final
conditions, Boyle’s law can be written as:
P1V1 = P2V2;
 1200

900

600

300

0
0.5 1.0 2.0 3.0 4.0 5.0 6.0
Pressure (atm)

Figure 3.1 Volume versus pressure graph for a gas at constant temperature and mass.

2. Charles’ law

States that the volume of a fixed mass of gas at constant pressure varies directly with the
Kelvin temperature. Mathematically,

V α T at constant P and n:

V = kT,
The value of T is the Kelvin temperature, and k is a constant. The value of k depends only
on the quantity of gas and the pressure. The ratio V/T for any set of volumetemperature
values always equals the same k. Charles' law can be applied directly to volume-
temperature problems using the relationship:

T2V1 = T1V2

where V1 and T1 represent the initial condition; V2 and T2 represent the new condition.

Figure 3.2 A plot of volume versus the Kelvin temperature.

 Note that the ratio V/T is constant for every plot of the curve.

3. The Combined Gas Law

Expresses the relationship between pressure, volume, and temperature of a fixed amount
of gas.

A sample of a gas often undergoes changes in temperature, pressure, and volume. When
this happens, the three variables must be dealt with at the same time.

Derivation of the combined gas law:

Boyle’s law: V α 1/P

Charles’ law: V α T Then, V

T/P (combined)
V = kT/P (where k is a constant)

It follows,
𝑃1𝑉1 𝑃2𝑉2
= k and =k
𝑇1 𝑇2

Since in each case k is constant, the combined gas law equation is given as follows:

𝑃1𝑉1 𝑃2𝑉2
= =k
𝑇1 𝑇2

Where P1, V1 and T1 are the initial pressure, volume and temperature; P2, V2 and
T2 are the final pressure, volume and temperature of the gas respectively.
Example

If a 50 cm3 sample of gas exerts a pressure of 60.0 kPa at 35°C, what volume will it
occupy at STP?

Solution:
Given V1=50cm3, P1 = 60kpa, T1 =35OC=308K V2 =?
at STP, P2 =1atm = 101325pa = 101.325kpa, T2 = 0oc=273K

𝑃1𝑉1 𝑃2𝑉2 60𝑘𝑝𝑎𝑥50𝑐𝑚3 𝑉𝑥101.325𝑘𝑝𝑎

= = = ,
𝑇1 𝑇2 308𝑘 273𝑘

60𝑘𝑝𝑎∗50𝑐𝑚3∗273𝑘
V2 = V2= 26.243cm3
𝑇308𝑘∗101.325𝑘𝑝
 4. Avogadro’s law: states that the volume of the gas is directly proportional to the
number of mole of gas, when the tempreture and pressure are held constant.

Mathematically, V α n; where V is the volume and n is number of moles.

Common example of avogadro’s law is the deflation of automobile tires.

5. The Ideal Gas Equation

 a hypothetical gas that obeys the gas laws. Real gases only obey the ideal
gas laws closely at high temperature and low pressure.
 The ideal gas law is a combination of Boyle’s law, Charles’ law and
Avogadro’s law.
Boyle's law: V α 1/p (at constant T and n)
P
Charles' law: V α T (at constant P and n)
Avogadro's law: V α n (at constant P and T)
This relationship indicates how the volume of gas depends on pressure,
temperature and number of moles.
nT 𝑛𝑇
Vα or V= R where R, is a proportionality constant called the gas
V 𝑃

constant.

PV = nRT
the ideal gas equation
Thus, the ideal gas equation describes the relationship among the four
variables P, V, T and n. An ideal gas is a gas whose pressure-volume-
temperature behavior can be completely explained by the ideal gas
equation. At STP, the values of R can be calculated from the ideal gas
equation.
6. Graham's Law of Diffusion
States that at constant temperature and pressure, the rate of diffusion of a gas, r, is
inversely proportional to the square root of its density, d, or molar mass, M.
Mathematically it can be expressed as:

where r is the rate of diffusion, d is the density and M is the molecular mass of the gas.
For two gases (Gas 1 and Gas 2), their rates of diffusion can be given as:

Rearranging these relationships gives the following expression

Where r1, d1 and M1 represent the rate of diffusion, density and molecular mass of gas 1. r 2, d2 and M2
represent the rate of diffusion, density and molecular mass of gas 2.

The rate at which a gas diffuses is also inversely proportional to the time taken.
Mathematically,

𝟏
r∝ 𝐭

If two different gases (gas 1 and gas 2) under the same conditions of temperature and
pressure diffuse through a porous container, then the time required to diffuse for the two
gases can be given by the following formula:

where t1 and t2 are the time taken, r1 and r2 are the rates, M1 and M2 are the molecular
masses of Gas 1 and Gas 2 respectively

3.5 The liquid state

As in a gas, particles in a liquid are in constant motion. However, the particles in a liquid
are closer together than those in a gas. The attractive forces between particles in a liquid
are more effective than between particles in a gas. This attraction between liquid
particles is caused by the intermolecular forces (dipole-dipole forces, London dispersion
forces, and hydrogen bonding). Liquids are more ordered than gases because of the
stronger intermolecular forces and the lower mobility of liquid particles. Accordingly,
liquid particles are not bound together in fixed positions.

Surface tension of liquids

Surface tension is property of a liquid surface displayed by its acting as if it were a

stretched elastic membrane. This phenomenon can be observed in the nearly spherical
shape of small drops of liquids and of soap bubbles. Because of this property,
certain insects can stand on the surface of water. A razor blade also can be supported by
the surface tension of water. The razor blade is not floating: if pushed through the
surface, it sinks through the water.

𝟏𝒙 𝑭
┌= ┌-surface tension, F-force, L- length
𝟐 𝑳

Surface tension depends mainly upon the forces of attraction between the particles
within the given liquid and also upon the gas, solid, or liquid in contact with it.
The molecules in a drop of water, for example, attract each other weakly. Water
molecules well inside the drop may be thought of as being attracted equally in all
directions by the surrounding molecules. However, if surface molecules could be
displaced slightly outward from the surface, they would be attracted back by the nearby
molecules. The energy responsible for the phenomenon of surface tension may be
thought of as approximately equivalent to the work or energy required to remove the
surface layer of molecules in a unit area. Surface tension may be expressed, therefore, in
units of energy per unit area. Water has a surface tension of 0.073 joule per square metre
at 20 °C. In comparison, organic liquids, such as benzene and alcohols, have lower
surface tensions, whereas mercury has a higher surface tension. An increase
in temperature lowers the net force of attraction among molecules and hence decreases
surface tension
Surface tension is the energy, or work, required to increase the surface area of a liquid
due to intermolecular forces. Since these intermolecular forces vary depending on the
nature of the liquid (e.g. water vs. gasoline) or solutes in the liquid (e.g. surfactants like
detergent), each solution exhibits differing surface tension properties

Molecular Perspective

In a sample of water, there are two types of molecules. Those that are on the outside,
exterior, and those that are on the inside, interior. The interior molecules are attracted to
all the molecules around them, while the exterior molecules are attracted to only the
other surface molecules and to those below the surface. This makes it so that the energy
state of the molecules on the interior is much lower than that of the molecules on the
exterior. Because of this, the molecules try to maintain a minimum surface area, thus
allowing more molecules to have a lower energy state. This is what creates what is
referred to as surface tension.

The water molecules attract one another due to the water's polar property. The hydrogen
ends, which are positive in comparison to the negative ends of the oxygen cause water to
"stick" together. This is why there is surface tension and takes a certain amount of
energy to break these intermolecular bonds. Same goes for other liquids, even
hydrophobic liquids such as oil. There are forces between the liquid such as Van der
Waals forces that are responsible for the intermolecular forces found within the liquid. It
will then take a certain amount of energy to break these forces, and the surface tension.
Water is one liquid known to have a very high surface tension value and is difficult to
overcome.

Cohesive and Adhesive Forces

There are several other important concepts that are related to surface tension. The first of
these is the idea of cohesive and adhesive Forces. Cohesive forces are those that hold the
body of a liquid together with minimum surface area and adhesive forces are those that
try to make a body of a liquid spread out. So if the cohesive forces are stronger than the
adhesive forces, the body of water will maintain its shape, but if the opposite is true than
the liquid will be spread out, maximizing its surface area. Any substance that you can
add to a liquid that allows a liquid to increase its surface area is called a wetting agent.

Viscosity of liquids

Resistance of a fluid (liquid or gas) to a change in shape, or movement of neighboring

portions relative to one another. Viscosity denotes opposition to flow. The reciprocal of
the viscosity is called the fluidity, a measure of the ease of flow. Molasses, for example,
has a greater viscosity than water. Because part of a fluid that is forced to move carries
along to some extent adjacent parts, viscosity may be thought of as
internal friction between the molecules; such friction opposes the development
of velocity differences within a fluid. Viscosity is a major factor in determining the
forces that must be overcome when fluids are used in lubrication and transported in
pipelines. It controls the liquid flow in such processes as spraying, injection molding,
and surface coating. For many fluids the tangential, or shearing, stress that causes flow
is directly proportional to the rate of shear strain, or rate of deformation, those results. In
other words, the shear stress divided by the rate of shear strain is constant for a given
fluid at a fixed temperature. This constant is called the dynamic, or absolute, viscosity
and often simply the viscosity

The unit of viscosity, accordingly, is newton-second per square metre, which is usually
expressed as pascal-second in SI units.

The viscosity of liquids decreases rapidly with an increase in temperature. Thus, upon
heating, liquids flow more easily. For example, the viscosities of water at 27 °C and at
77 °C are 8.5 ×10−4 and 3.6 × 10−4 pascal-second, respectively.

Inter molecular forces

The forces holding liquids together are called intermolecular forces. It is weaker than
intramolecular forces (e.g. ionic, metallic, or covalent bonds)
The strength of these intermolecular forces is directly related to the melting/boiling
points, enthalpy of fusion, enthalpy of vaporization, and solubility of the substances. The
strength of these intermolecular forces is directly related to the melting/boiling points,
enthalpy of fusion, enthalpy of vaporization, and solubility of the substances.

Factors Affecting Dispersion Forces

1. Number of electrons in an atom (more electrons, more dispersion force)

2. Size of atom or molecule/molecular weight
3. Shape of molecules with similar masses (more compact, less dispersion force).
The shape of the molecule affects the dispersion forces: long, thin molecules
(like n-pentane tend to have stronger dispersion forces than short, round ones
(like neopentane). This is due to the increased surface area in n-pentane.
Types of Attractive Forces

There are several types of attractive intermolecular forces:

Dipole-dipole forces, London dispersion forces, Hydrogen bonding, and Induced-dipole

forces. The first three forces are also collectively called van der Waals forces. All
molecular and intermolecular attractive forces are electrostatic in nature. That is, they
involve attractions between positive and negative species. The strengths of these
attractive forces vary widely, though usually the intermolecular forces between small
molecules are weak compared to the intramolecular forces that bond atoms together
within a molecule. All of the intermolecular forces that hold a liquid together are
called cohesive forces.

1. Dipole-Dipole Forces
Polar molecules that have permanent dipoles are attracted to each other via
electrostatic attractions. •The partial positive, δ+, end of one is attracted to the
partial negative, δ−, end of the other and vice-versa. •These forces are only
important when the molecules are close to each other.

Figure 3.3 dipole forces

2. London Dispersion Forces

All molecules, including those without dipole moments, exert forces on each
other. We know this because all substances, even the noble gases, change from
liquid to solid state under different conditions. It act between all atoms and
molecules. They are the only forces that exist between noble gas atoms and non-
polar molecules.
Because dispersion forces result from temporary redistribution of the electrons
causing induced dipoledipole interactions, their strength increases with the
number of electrons in the interacting atoms or molecules. Hence, dispersion
forces increase with atomic number or molar mass. This trend can be seen by
 comparing the boiling points of gases (helium, He, and argon, Ar), (hydrogen,
H2, and oxygen, O2 ), and (chlorine, Cl2, and bromine, Br2).
3. Hydrogen Bonding
A special form of dipole-dipole attraction, which enhances dipole-dipole
attraction. It is a particular type of intermolecular force arising when a hydrogen
atom is bonded to highly electronegative elements, fluorine, oxygen and
nitrogen.

Figure 3.4 hydrogen bonding

Hydrogen bonding in ice results in an ordered, open arrangement. Density of ice at 0 °C

is 0.917 g/mL vs liquid water at 0 °C is 1.00 g/mL

3.6 The solid state

3.6.1 classification of solids

Based on their crystal structures, solids can be classified into the following categories:

1. Crystalline solids(true solids)

The solids featuring highly ordered arrangements of their particles (atoms, ions, and
molecules) in microscopic structures are called crystalline solids. These ordered
microscopic structures make up a crystal lattice that accounts for the structure of the solid
at any given point. Examples of crystalline solids include salt (sodium chloride),
diamond, and sodium nitrate.
Crystalline solids consist of atoms, ions, and molecules arranged in a strongly ordered
microscopic arrangement in consistent and repeated three-dimensional structures,
forming a crystal lattice that stretches in any direction. What are the 7 types of crystals?
Seven crystal structures are available in total: triclinic, monoclinic, orthorhombic,
tetragonal, trigonal, hexagonal, and cubic.

Geometry:
Crystalline Solids – Particles are arranged in a repeating pattern. They have a regular and
ordered arrangement resulting in a definite shape.

Amorphous Solids – Particles are arranged randomly. They do not have an ordered
arrangement resulting in irregular shapes

Melting Points

Crystalline Solids – They have a sharp melting point

Amorphous Solids – They haven’t sharp melting points. Tends to soften gradually over a
temperature range

Heat of Fusion: (The change in enthalpy when a substance is heated to change its state
from solid to liquid.

Crystalline Solids – They have a definite heat of fusion.

Amorphous Solids – They do not have definite heat of fusion

Cleavage Property

Crystalline Solids – When cutting with a sharp edge, the two new halves will have
smooth surfaces
Amorphous Solids- When cutting with sharp edge, the two halves will have irregular
surfaces

Rigidity:
Crystalline Solids – They are rigid solids and applying mild forces will not distort its
shape.
Amorphous Solids – They are not rigid, so mild effects may change the shape

2. Amorphous solids (Pseudo – Solids or super-cooled liquids)

The solids in which the particles are not arranged in any specific order or the solids that
lack the overall order of a crystal lattice are called amorphous solids.

The term ‘amorphous’, when broken down into its Greek roots, can be roughly translated
to “without form”. Many polymers are amorphous solids. Other examples of such solids
include glass, gels, and nanostructured materials.

In compounds like coal and charcoal, carbon also occurs as an amorphous, or

“shapeless,” form. Allotropes are called varying variants of the same substance.

However, crystalline solids can be further classified into molecular, ionic, metallic, and
covalent solids. A brief introduction to the classification of solids is provided in this
article. Classification of Solids – Crystalline and Amorphous Solids

Crystalline solids are made of stone, wood, paper and cloth. Such solids consist of atoms
arranged in a particular fashion. The transition to liquid, called melting, is sharp and
transparent as crystalline solids are heated. Amorphous solids are made of rubber, glass,
and sulphur.

3.7 Phase transition in solids

A phase is any part of a system that has uniform composition and properties. A state of
matter represents a phase. Most solid substances undergo two changes of state when
heated. A solid change to a liquid at the melting point, and the liquid changes to vapor at
the boiling point. To understand this consider the heating curve for a substance given in
Figure 3.5. A heating curve is a plot of temperature verses the uniform addition of heat.
This can be illustrated for a hypothetical substance, in which the temperature of the
substance is on the vertical axis and the passage of time during which heat is added to the
substance is on the horizontal axis. Figure 3.5 shows the changes in the temperature and
phases of a pure substance as it is heated, beginning with a solid and continuing to the
gaseous state as described.

Figure 3.5 Heating curve.

Initially, the substance exists in the solid state, and the addition of heat increases its
temperature. When the solid is heated, its temperature rises (A to B) until it reaches the
melting point (point B), and the temperature remains constant (B to C) until all the solid
is converted to a liquid (point C). The added heat energy is used to break the
intermolecular forces, thus disrupting the solid structure. At point C phase change is
completed. Once melting is completed, heating of the liquid raises its temperature (C to
D) until the boiling point is reached at point D. In region (D to E) the addition of heat is
utilized to break the intermolecular forces of the liquid to change it to a gas.

Unit four
4. SOLUTIONS
4.1 Introduction
Solution is a homogeneous mixture in which no settling occurs and in which solute
particles are at the molecular or ionic state of subdivision. Solutions have no fixed
composition. However, the components of solution should be distributed uniformly
throughout the entire solution A homogeneous mixture is a mixture, in which the
composition of the mixture is the same throughout, i.e., it has no visible boundaries
because the components are mixed as individual atoms, ions and molecules. Example:
air, sugar solution, salt solution, alloys, soft drinks (pepsi, coca cola, etc.), gasoline and
so on.
There are two components of a solution: solute and solvent. A solute is a component of a
solution present in a smaller amount than the solvent. A solvent is either a components of
solution that is present in a large amount or the component that determines the physical
state of the solution.

4.2 Types of solutions

A typical solution consists of at least two components;
 solute: substance being dissolved; present in smaller or lesser amount
 solvent: substance doing the dissolving; present in larger amount
Solutes and solvents may be of any form of matter: solid, liquid or gas. Therefore,
solutions may exist in any of the three states of matter. In general, the physical state of
the solvent determines the physical state of the solution.
PROJECT
Do you know how jewellery gold is made?
Your teacher will arrange a visit to the nearby goldsmith. Ask the following questions during
your visit to the goldsmith:
a. How is jewellery made from gold?
b. What are the possible components, other than gold present in the jewellery?
c. Which component is present in large quantity?
d. What is the advantage of mixing gold with other materials?
e. Can you separate the components present in the jewellery?
f. How purity of gold assessed and what is the unit of measuring the purity?
g. Write a report about what you have observed during your visit and present it to the
class.

Exercise 4.1
Instruction: Choose the correct answer from the alternatives given and circle the
letter of your choice.
1. What is the solvent in 70 % alcohol solution? (UEE, 2008)
A) Water B) Alcohol C) Sugar D) Kerosene
2. Which of the following types of solutions are possible? (UEE, 2002)
I. solid dissolves in a liquid II. Gas dissolved in a liquid
III. Gas dissolved in a gas IV. Solid dissolved in a solid
A) I and II B) I, II, III and IV C) I, II and IV D) I
3. Which of the following is NOT a solution? (UEE, 2004)
A) Milk B) Brass C) Whisky D) Coca cola drink
4. What percent is gold in 18 karat gold? (UEE, 2003)
A) 18 B) 09 C) 25 D) 75
5. What kind of solution forms when gasoline evaporates in air? (UEE, 2008)
A) Gas in gas solution B) Gas in liquid solution
C) Liquid in liquid solution D) Liquid in gas solution
6. Give two examples of:
a. gas-gas solution
b. gas-liquid solution
c. liquid-liquid solution
d. solid-liquid solution

4.3 The solution process

4.3.1 Liquid solutions and inter-particle force of attractions
What are the major factors that affect the solubility of substances in liquid solvents?
When a solute dissolve in a solvent, particles of the solute disperse throughout the
solvent. The solute particles occupy positions that are normally taken by solvent
molecules i.e., particles of the solute mix with particles of the solvent. The tendency to
mix is affected by the relative strengths of three types of interaction. These are:
 Solvent-solvent interaction
 Solute-solute interaction
 Solvent-solute interaction
For simplicity, we can imagine the solution process taking place in three distinct steps.
Figure 4.1 A molecular view of the solution process, portrayed as taking place in three
steps

Step 1: is the separation of solvent molecules from each other to “make room” for the
incoming solute particles and
Step 2: is the separation of solute molecules. These steps require energy input to break
attractive intermolecular forces; therefore, they are endothermic.
Step 3: the solvent and solute molecules mix. This process can be exothermic or
endothermic.
Strong forces of attraction between solute particles (solute-solute interactions) or between
solvent particles (solvent-solvent interactions) tend to keep like particles close together
and reduce the solubility of solute in the solvent. On the other hand, strong attraction
forces between solute and solvent particles (solute-solvent interactions) makes dissolving
easier and helps to keep particles in solution.
Activity: Ethanol mixes with water but oil does not. Why?
The saying “like dissolves like” is helpful in predicting the solubility of a substance in a
given solvent.
Solubility is a measure of how much solute will dissolve in a solvent at a specific
temperature.
“Like dissolves like” principle states that two substances with intermolecular forces of
similar type and magnitude are likely to be soluble in each other. This is because the
forces between the solute molecules are replaced by the forces between the solute and
solvent molecules.
For example: carbon tetrachloride and benzene
Water is the most prominent solvent. This is because it is common and it dissolves a large
number of ionic and polar substances. Water molecules are strongly associated with each
other through hydrogen bonding. Compounds that form hydrogen bond such as certain
alcohols tend to be soluble in water. Examples: methanol, ethanol and 1, 2-ethylene
glycol

If a compound has both polar and non-polar components, it may dissolve in both polar
and non-polar solvents.
For example: Acetic acid, CH3COOH, is a liquid that forms hydrogen bonds with water.
It is fully miscible with water. Acetic acid also dissolves in non-polar solvents, such as
benzene and carbon tetrachloride, because the CH3 component is non-polar.
 Fluids that mix with or dissolve in each other in all proportions are said to be
miscible fluids.
 If two fluids do not mix but, rather, form two layers, they are said to be immiscible
fluids.
For example: acetone is miscible in water where as gasoline, which is a mixture of
hydrocarbons, is immiscible with water.

Exercise 4.2
Instruction: Choose the correct answer from the alternatives given and circle the
letter of your choice.
1. Which of the following is most likely to be miscible in carbon tetrachloride, CCl4?
(UEE 2001)
A) Na2CO3 B) CH3F C) (CH2)4O2 D) I2
2. What would be the solubility of HOCH2(CH2)6CH2OH compared to
CH3(CH2)6CH2OH? (UEE, 2005)
A. A) Less soluble in water
B. More soluble in water
C. The same solubility in water
D. More soluble in a non-polar solvent such as dichloro ethane
3. The dissolution of water in octane (C8H18) is prevented by (UEE, 2007)
A) Dipole-dipole attraction between octane molecules
B) Hydrogen bonding between water molecules
C) London dispersion forces between octane molecules
D) Repulsion between like charged water and octane molecule
4. Gasoline and water do not mix because gasoline is (UEE, 2000)
A. Less dense than water C. Less viscous than water
B. Volatile D. Nonpolar
5. Which of the following compounds is likely to be more soluble in water: C4H9OH or
C4H9SH? Explain.
6. Would boric acid, B(OH)3, be more soluble in ethanol, C2H5OH, or in benzene,
C6H6?
7. Would naphthalene, C10H8, be more soluble in ethanol, C2H5OH, or in benzene,
C6H6?
8. Arrange the following substances in order of increasing solubility in hexane, C6H14:
CH2OHCH2OH, C10H22, H2O.
9. Indicate which of the following is more soluble in ethanol, C2H5OH: acetic acid,
CH3COOH, or stearic acid, C17H35COOH.
Solutions of Solids in Liquids
Activity 4.1
Take two containers and add water to both. Put table salt in the first container and table sugar inthe
second container and stir.
a. What did you observe?
b. Did both solids dissolve in water?
Repeat the same procedure using oil instead of water.
c. What did you observe?
d. Did both solids dissolve in oil?

Solids that are made up of polar molecules or ions are not soluble in non-polar solvents.
The interaction between the polar solid and the non-polar liquid is weak compared to the
interaction within the solid. Therefore, the solute and solvent particles do not mix. For
example sugar which has molecules that are likely bound together by hydrogen bonding
is not soluble in solvents like oil. Sugar does dissolve in water though.

Why does sugar dissolve in water, but not in oil?

Sugar dissolves in water because the water molecules attract the sugar molecules in the
same way that the sugar molecules attract each other. Both of these substances are polar.
The Sugar does not dissolve in the oil because the sugar is polar but the oil is non-polar.
The interaction between these two substances is weaker than the interaction within sugar
molecules.
Solids containing crystals held together by London forces can dissolves somewhat in
non-polar solvents.
However, they do not dissolve in polar solvents. For the same reason non-polar liquids do
not dissolve in polar solvents.
Water as well as other polar molecules is attracted to ions.
The electrostatic attraction between an ion and a molecule with a dipole is called an ion-
dipole attraction. These are important in the dissolution of ionic compounds in water.
When ionic compounds dissolve in water, the ions in the solid separate because they are
surrounded by water molecules. The water molecules penetrate the compounds and
surround the individual components.
This reduces the strong inter-ionic forces that bind the ions together. This allows them to
move into the solution as hydrated ions. Thus, when a solute particle becomes surrounded
by solvent molecules, we say that it is solvated. When the solvent is water we call the
process hydration. This image shows a potassium chloride (KCl) molecule being solvated
or hydrated.

Solvation is the process of surrounding a solute particle with solvent particles. When the
solvent is water the process is called hydration.

4.3.2 The Rate of Dissolution

Dissolution is the process of dissolving a solute in a solvent to form a solution.
Solute + solvent ⟶ solution
The rate of dissolution is the speed with which a solute goes into solution. This largely
depends upon the inter-particle forces and to a lesser extent, on conditions such as the
surface area of the solid solute, the temperature and the pressure of the system.
Inter-particle force is the extent to which solvent molecules interact with particles of the
solute. When the solvent-solute interactions are stronger than those between solute-solute
and solvent- solvent particles, the dissolution process becomes easier.
Increasing surface area will increase the rate of dissolution because it increases the
number of solute particles in contact with the solvent.

4.3.3 Energy Changes in the Solution Process

Heat of Solution
The heat of solution of an ionic compound in water is the sum of the lattice energy of the
compound and the heat of hydration. The relative magnitudes of these two quantities
determine whether the solution process is endothermic or exothermic.
ΔHsoln value may be exothermic or endothermic, depending on the substance involved.
For instance, the ΔHsoln value for: CaCl2 (s) is -81.3 kJ/mol and NH4NO3(s) is +25.7
kJ/mol

Hydration of Ionic Solids in Water

Activity 4.2
Discuss the following questions and write a report. Present your report to the class.
1. How do you compare the hydration of the ions of Group I metals with that of Group II
metals? Give reasons.
2. Compare the solubility of Mg(OH)2 with Ba(OH)2, MgSO4 with BaSO4. Justify your
answers in terms of lattice energy and hydration energy.

Hydration of ions favors the dissolution of an ionic solid in water. The ions in an ionic
crystal are very strongly attracted to one another. Therefore, the solubility of an ionic
solid depends not only on the energy of the hydration of ions, but also on lattice energy,
which are the energy holding ions together in the crystal lattice.
 The energy required to completely separate one mole of a solid ionic compound into
gaseous ions is called lattice energy, ΔHLattice energy. It is always a positive quantity.

 The enthalpy change associated with the hydration process is called the heat of
hydration,
ΔHhydr. It is always a negative quantity.

Lattice energy works against the solution process, so an ionic solid with relatively large
lattice energy is usually insoluble. Lattice energies depend on the charge on the ions and
also the distance between the centers of the neighboring positive and negative ions.
 As the charge of the ions increases, the lattice energy increases. For example, the
energy of hydration is greater for Mg2+ than for Na+
MgO has greater lattice energy than NaCl
 As the size of the ions increases, the lattice energy decreases
Note: Overall as we go down in a group, ionic radius increases and lattice energy
decrease. In a period from left to right as the charge on ion increases, lattice energy
increases.
 The magnitude of lattice energies depends predominantly on the ionic charges
because ionic radii vary over only a limited range.
The energy of hydration also depends on:
 ionic radius, the energy of hydration is greatest for a small ion
 charge on the ion, the energy of hydration increases with the charge on the ion
Solved problems
1. Arrange the following ionic compounds in order of increasing lattice energy: NaF,
CsI, and CaO.
Solution: CsI < NaF < CaO
Consider the dissolution of sodium chloride (NaCl) in water.

Therefore, when 1 mole of NaCl dissolves in water 4 kJ of heat will be absorbed

from the surroundings.
2. What is the relationship between solvation and hydration?
3. Solvation can be the same as hydration if the solvent used for the solution process
is water. If the solvent is not water, we use the term solvation to describe the
process in which solvent molecules surround solute particles.
4. Sodium carbonate is soluble in water, while calcium carbonate is insoluble. Why?
Answer: Because the lattice energy of calcium carbonate CaCO3 is higher than
that of Na2CO3. Thus, CaCO3 is less soluble than Na2CO3.
5. Which ion in each pair has the larger heat of hydration?
 Na+ or Cs+ b) O2- or F- c) Na+ or Cl-
Exercise 4.3
Instruction: Choose the correct answer from the alternatives given and circle the
letter of your choice.
1. What is the most important type of solute-solvent interaction in a solution of
CaCl2 in water? (UEE 2003 E.C)
A) Dipole-dipole B) Dispersion forces
C) Hydrogen bonding D) Ion-dipole
2. Which of the following is the most important type of solute-solvent interaction in
a solution of n-butanol in water? (UEE, 2005)
A) Dispersion B) Ion-dipole C) Dipole-dipole D)
Hydrogen bonding
3. What type of solute-solvent interactions should be the most important in a
solution of iodine in carbon tetrachloride? (UEE, 2007)
A) London forces B) Ionic bonding C) Ion-dipole forces D)
Dipole-dipole forces
4. Which of the following ionic compounds has the greatest lattice energy? (UEE,
2007)
A) LiF B) LiCl C) LiBr D) LiI
5. The exothermic step in the dissolution of a salt in water is (EHEECE, 2003)
a. Separation of the salt into ions in gaseous state
b. Separation of the water molecules
c. Hydration (salvation) of the ions from the salt
d. Separation of the salt into ions in gaseous state and separation of the water
6. Indicate the type of solute-solvent interaction that should be most important in
each of the following solutions:
a. AgNO3 in water
b. Ammonia in water
c. Bromine in CCl4
d. KBr in water
e. NaOH in water
 f. Sugar in water
g. Toluene in cyclohexane
7. The rate of dissolution largely depends on inter-particle forces. What other factors
influence the rate of dissolution?
8. Which of the following ions would be expected to have the greater energy of
hydration?
a) Mg2+ or Al3+ b) F- and Cl-
4.4. SOLUBILITY
Solubility is defined as the maximum amount of a solute that will dissolve in a given
amount of solvent at a specified temperature. Solubility is a characteristic property of a
specific solute–solvent combination, and different substances have greatly differing
solubility. Usually, it is expressed as the amount of solute in gram that dissolves in 100
ml of water.
Example
a) The solubility of NaCl is 39.12 g /100 mL of water at 100°C
b) The solubility of AgCl is 0.0021 g /100 mL water at 100°C

Factors Affecting Solubility

"The solubility of a solute is the number of grams of the solute necessary to saturate
100gm of the solvent at that particular temperature. There are three main factors that
control solubility of the solute.
(1) Temperature
(2) Nature of solute or solvent
(3) Pressure
Effect of Temperature on Solubility
The solubility of most molecular and ionic solids increases with temperature, although
some are almost unchanged, and some decrease.
In endothermic process solubility increases with the increase in temperature and vice
versa
 For a solute with ΔHsoln > 0:
 solute + solvent + heat → Saturated solution
 Solubility increases with temperature. For instance, solubility of potassium nitrate
increases with the increase in temperature
In exothermic process solubility decreases with the increase in temperature
 For a solute with ΔHsoln < 0:
Solute + Solvent → Saturated solution + heat
 Solubility decreases with temperature.
For example solubility of sodium sulphate decreases with the increase in temperature, fig.
4.2

Fig 4.2. Solubility vs temperature of some salts

Solubility of Gases in Water:

Gases are more soluble in cold solvent than in hot solvent (see fig 4.3). Gases become
less soluble at higher temperatures.
Example.
 Soft drinks become “flat” as they warm up and lose carbon dioxide.
 Aquatic life is affected by decreasing amounts of dissolved oxygen as a result of
thermal pollution.
 Temperture (0C)
Fig 4.3. Solubility vs temperature of some gas
In general,
 for most exothermic solution processes, solubility decreases with an increase in
temperature
 for most endothermic solution processes, solubility increases with an increase in
temperature
Nature of Solute and Solvent
Solubility of a solute in a given solvent depends purely on the nature of both the solute
and solvent. A polar solute dissolves in a polar solvent whereas a polar solute has low
solubility in non-polar solvent.
Liquids that mix with each other in all proportion are said to be completely miscible with
each other. Liquids that mix with water in all proportions are usually polar substances or
substances that form hydrogen bonds. For such liquids, the dipole-dipole attraction or
hydrogen bonding of the solute molecules with the solvent molecules are at least as
strong as those between molecules in the pure solute or in the pure solvent. Hence the
two kinds of molecules mix easily.
Two liquids that do not mix are called immiscible. When two immiscible liquids are
poured in to the same container, two layers are formed. Immiscibility occurs when either
the solvent or solute is non polar and the other is polar or vice versa.
Gasoline, carbon disulfide, benzene, Carbon tetrachloride, and many other non-polar
liquids are immiscible with water (polar liquid).There is no effective attraction between
the molecules of such non-polar liquids and polar water molecules.
 Non-polar liquids may be miscible with each other because of the absence of
appreciable tendency of attraction between molecules of similar kind or molecules of
another non polar liquid.

Polar molecules are soluble in polar solvent and non-polar molecules in non-
polar solvent (like dissolves like).

There are still partially miscible liquids. Liquids, such as ether and water or bromine and
water, are said to be partially miscible. The two partially miscible liquids usually form
two layers when mixed. The general rule in solubility is that “like dissolves like”.
Hydrogen bonds- substances with O—H and N—H bonds are often soluble in water
because of H-bonding
Dipole-dipole forces -polar solutes interact well with polar solvents through attraction of
partial charges.
Ion - induced dipole forces- responsible for the attraction between Fe2+ and O2
molecules in the bloodstream.
Dipole - induced dipole forces are responsible for the solvation of gases (nonpolar) in
water (polar).
London (dispersion) forces -the principal attractive force in solutions of nonpolar
substances (e.g. petroleum).

Activity: Decide whether the following liquids are miscible, immiscible or partially miscible

Mixture Conditions

Methanol + Water

Hexane + water

Hexane+ Benzene

Iodine + Carbontetrachloride

Iodine
Effect of pressure on solubility

Pressure has little effect on the solubility of solids and liquids, but has a large effect on
the solubility of gases. An increase in pressure increases solubility of a gas in a liquid. At
a given pressure, there is equilibrium between the gas which is dissolved in the solution
and the gas in the vapor phase. If the pressure increases, more gas dissolves to reduce the
“extra” pressure; the new equilibrium is established with more gas dissolved. . For
example, carbon dioxide is filled in soft drink (such as coca cola, Pepsi, 7up, etc.) under
pressure

When you open the bottle of carbonated beverage, the pressure decreases and some gases
escape from the solution. This escape of bubbles of a gas from a liquid is known as
effervescence. It occurs due to decrease in pressure.
All gases become soluble in a liquid at a given temperature when the partial pressure of
the gas over the solution is increased. The solubility of gas is governed by Henry’s law.
Henry’s law states that the solubility of a gas (Sgas, in mol/L) is directly proportional to
the partial pressure of the gas (Pgas, in atm) over the solution.

Sgas = kH Pgas

Where, kH is Henry’s law constant and it is unique to every gas, at a given temperature,
with units of mol L-1 atm-1.
This expression can be rearranged to deal with changing solubility and pressure

S1 S 2
  kH
P1 P2
 FIGURE 4.4 The effect of pressure on gas solubility

Table 4.1 Henry’s law constants for some gases in water at 250C

Gas KH (M/atm)
O2 1.3x10-3
N2 6.1x10-4
CO2 3.4x10-2
NH3 5.8x102
He 3.7x10-4

Example:
Calculate the solubility of oxygen in water at 25oC for a partial pressure of 0.22 atm.
Solution: - From table 2.3, the value of solubility constant for oxygen is 1.3x10-
3mol//L.atm. Then, the Solubility of oxygen is
S O 2 = kH.PO2

= 1.3x10-3mol/L.atmx0.22 atm.
= 2.86x10-4 mol/L
Example:
1. The partial pressure of carbon dioxide gas inside a bottle of mineral water is 4 atm
at 25°C. What is the solubility of CO2? The Henry’s law constant for CO2
dissolved in water is 3.3×10−2 mol/Latm at 25°C.
Solution: CCO2=KH PCO2=3.3×10−2molL atm⁄ ×4 atm = 0.1molL/L
 2. Calculate the molar concentration of oxygen in water at 25°C for a partial
pressure of 0.22 atm.
Solution: CO2=KH PO2=1.23×10−3 molL atm/ × 0.22 atm = 2.706×10−4molL⁄

Exercise 4.4
Instruction: Choose the correct answer from the alternatives given and circle the
letter of your choice.

1. Which law relates the concentration of a dissolved gas, Cg, to its partial pressure?
(UEE 2004)
A) Henry’s law B) Raoult’s law
C) Boyle’s law D) Ideal gas law
2. Which of the following is NOT correct? (UEE, 2003)
a. Addition of a strongly solvated solute decreases the solubility of a gas
in liquid
b.For solid dissolved in liquid that is exothermic, increase in temperature
increases the solubility
c. The solubility of a substance is its concentration in the saturated
solution
d.For solution to occur solvent-solute attractions must overcome solute-
solute and solvent-solvent attractions
3. The solubility of oxygen gas in wter at 25°C and 1.0 atm pressure of oxygen is 0.04
g/L. the solubility of oxygen in water at 3.0 atm and 25°C is _____ g/L. (UEE, 2007)
A) 0.014 B) 0.31 C) 0.041 D)
0.123
4. Which of the following does NOT affect the solubility of a gas dissolved in a
liquid?(2008)
A. Nature of solute and solvent C. Pressure
B. Temperature D. Rate at which the gas
dissolves
4.5. SOLUBILITY AS AN EQUILIBRIUM PROCESS
 Activity 4.3
Form a group and perform the following activity.
i. Pour 50 mL water in a beaker.
ii. Add some crystal of Na2SO4 into the water using spatula and stir until it
dissolves.
What do you call this type of solution?
iii. Continue adding more and more Na2SO4 stirring to dissolve.
What do you observe after addition of large amount of solute?
iv. Filter the undissolved solute. Collect the filtrate or the solution.
Why does the solute remain undissolved?
What is the name of such a solution?
v. Add some more solute to the filtrate and stir.
Does the additional solute dissolve?
Discuss your findings with the rest of the class.

Since solutions are mixtures of two or more substances they are not governed by the law
of constant composition. Hence, one can make solutions of different concentrations from
the same solute and solvent. So, solutions can be unsaturated, saturated or supersaturated.
When an ionic solid dissolves, ions leave the solid and become dispersed in the solvent.
Some dissolved ions collide occasionally with the undissolved solute and recrystallize.
As long as the rate of dissolving is greater than the rate of recrystallizing, the
concentration of ions rises. Eventually, given enough solid, ions are dissolving at the
same rate as ions in the solution are recrystallizing. At this point, even though the
dissolving and recrystallizing continue, there is no further change in the concentration
with time. The system has reached equilibrium; that is, excess undissolved solute is in
equilibrium with the dissolved solute:
Solute (Undissolved) Solute (dissolved)

This solution is called saturated. Therefore, Chemists also characterize solutions by their
capacity to dissolve a solute as
 Saturated
  Unsaturated and
 Supersaturated
 A saturated solution contains the maximum amount of a solute that will dissolve in a
given solvent at a specific temperature; the dissolved and undissolved solutes are in
dynamic equilibrium. Or a solution that is in equilibrium with respect to a given
dissolved substance.
 An unsaturated solution contains less solute than it has the capacity to dissolve. Or a
solution not in equilibrium with respect to a given dissolved substance and in which
more of the substance can dissolve.
 A supersaturated solution is a solution that contains more dissolved substance than a
saturated solution does.
 Honey is an example of naturally occurring supersaturated solution. It contains
glucose, as a solute, and water, as a solvent. If honey is left to stand, the glucose
eventually crystallizes.
Supersaturated solutions are not in equilibrium with the solid substance. If a small crystal
of ionic solid is added to a supersaturated solution, the excess immediately crystallizes
out.
 Crystallization is the process in which dissolved solute comes out of solution and
forms crystals. Adding small crystals to a supersaturated solution to crystallize out the
excess is called seeding.

Activity 4.4
Prepare a saturated solution of sodium thiosulphate (Na2S2O4) in beaker containing 50 mL water.
Heat the solution so that extra amount of solute dissolves. Filter and divide the solution into two
parts in two different beakers. Keep both the solutions to cool slowly undisturbed. After cooling stir
one of the solutions and observe what happens. Drop a small grain of crystal in the second one and
observe the result.
a. What is the name of such a solution that is made by cooling concentrated solution?
b. What happens when the solution was stirred?
c. What about the solution to which a crystal grain is dropped? Why?
d. What did you understand from the activity?
Exercise 4.4

Instruction: Choose the correct answer from the alternatives given and circle the
letter of your choice.
5. The solubility of sodium selenate, Na2S2O4, is 84 g/100 g of water at 35°C. If a
solution is obtained by dissolving 92 g of Na2S2O4 in 200 g of water at 35°C, what do
you call this solution? (UEE, 2004)
A. Diluted B. Saturated C. Unsaturated D.
Supersaturated

6. A tiny crystal of solid sodium acetate is added to three aqueous solutions of sodium
acetate. Classify each original solution as saturated, unsaturated, or supersaturated.
a) The added sodium acetate just sits there.
b) The added sodium acetate causes more solid sodium acetate to form.
c) The added sodium acetate dissolves.

7. You are given three bottles containing aqueous solutions of X one saturated, one
unsaturated and one supersaturated. How can you identify which solution is which?

4.6 WAYS OF EXPRESSING CONCENTRATIONS OF SOLUTIONS

The concentration of a solution is the amount of solute present in a given quantity of
solution. Chemists use several different concentration units, each of which has
advantages as well as limitations. Example, mass percentage, ppm and ppb, mole
fraction, molarity, normality and molality.

4.6.1 Mass Percentage, ppm and ppb of Solute

Percent by mass
The percent by mass, also called percent by weight or weight percent, is the ratio of the mass
of a solute to the mass of the solution, multiplied by 100.
mass of solute
% by mass of solute = X 100%
mass of solution

Example 1:
 A solution is made by dissolving 13.5 g of glucose, C6H12O6, in 0.100 kg of water.
What is the mass percentage of solute in this solution?
Solution
mass of glucose
% by mass of glucose = mass of solution X 100%
13.5g
% by mass of solute = 13.5g+ 100g X 100%

=11.9%
(b) How would you prepare 425 g of an aqueous solution containing
2.40% by mass of sodium acetate, NaC2H3O2?
Solution: The mass of sodium acetate (solute) in 425 g of solution is
mass of sodium acetate
% by mass of NaC2H3O2 = X 100%
mass of solution
% by mass of NaC2H3O2 x mass of solution
Mass of NaC2H3O2 = 100%
2.40% x 425g
Mass of NaC2H3O2 = 100%

=10.2g
Therefore, the quantity of water in the solution is
Mass of H2O=mass of solution − mass of NaC2H3O2=425 g−10.2 g=415 g
You would prepare the solution by dissolving 10.2 g of sodium acetate in 415 g of water.

Parts per million (ppm) and parts per billion (ppb)

When the mass of solute in the solution is very small, a widely used expression is parts
per million.
mass of slute
ppm of solute= X 10-6
mass of solution

ppm= mass percentage of solute×104

Example 2:
A 2.5 g sample of ground water was found to contain 5.4 µg of Zn2+. What is the
concentration of Zn2+, in parts per million?
Given:
Mass of sample = 2.5 g
Mass of Zn2+ = 5.4 µg = 5.4 × 10-6g
ppm of Zn2+ =?
Solution:
mass of slute
ppm of Zn2+= X 10-6
mass of solution
5.4 𝑥
ppm of Zn2+= X 10-6g
2.5 g

=2.16 ppm
1. Express 5% (m/m) solution of NaCl in ppm.
Given:
% by mass of NaCl = 5% 4
ppm of NaCl = ?
Solution:
ppm = mass percentage of solute × 104
= 5 × 104 ppm
 For solutions that are even more dilute, parts per billion (ppb) is used. A
concentration of 1 ppb represents 1 g of solute per billion (109) grams of solution.
Exercise: 4.5
1. A sample of 0.892 g of potassium chloride (KCl) is dissolved in 54.6 g of water.
What is the percent by mass of KCl in this solution?
2. A solution is prepared by mixing 1.00 g of ethanol with 100.0 g of water. Calculate
the mass percent of ethanol in this solution.
3. A 135 g sample of seawater is evaporated to dryness, leaving 4.73 g of solid residue
(the salts formerly dissolved in the seawater). Calculate the mass percent of solute
present in the original seawater.
4. A commercial bleaching solution contains 3.62 mass % sodium hypochlorite, NaOCl.
What is the mass of NaOCl in a bottle containing 2.50 kg of bleaching solution?
5. What mass of solution containing 5.00% potassium iodide, KI, by mass contains 258
mg KI? (Ans. 5.16 g)
6. If 150 g of orange juice contains 120 mg of ascorbic acid (Vitamin C), what is the
concentration of ascorbic acid, expressed in ppm?
7. Express the concentration of ascorbic acid given in (10) in ppb.
8. Find the concentration of calcium (in ppm) in a 3.50 g pill that contains 40.5 mg of
Ca. (Ans.1.16 × 104 ppm Ca)
9. The label on a 0.750 L bottle of Italian Chianti indicates “11.5% alcohol by volume.”
How many liters of alcohol does the wine contain? (Ans. 0.0862 L)
10. Seawater contains 0.0079 g Sr2+ per kilogram of water. What is the concentration of
Sr2+ measured in ppm?
Mole Fraction (X)
The mole fraction (X) of a solute is the ratio of the number of moles of solute divided by
the total number of moles of a solution (moles of solute + moles of solvent).
Number of moles of solute
Mole fraction for a non-electrolyte (X) = Number of moles of solute + Number of moles of solvent

mole of solute nsolute

x 
mole of solution nsolute  nsolvent

𝑚𝑜𝑙𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑚𝑜𝑙𝑒 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡

Xsolute = 𝑚𝑜𝑙𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒+𝑚𝑜𝑙𝑒 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 ; Xsolvent = 𝑚𝑜𝑙𝑒 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡+𝑚𝑜𝑙𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒

Xsolute + Xsolvent = 1
Example:
1. Calculate the mole fraction of ethanol when 23.0g of ethanol is mixed with 18 .0g of
water to make aqueous solution of ethanol.
Solution: Molar mass of ethanol (C2H5OH) is 46g/mol; that of water is 18g/mol.
23g
Mole of C2H5OH = = 0.5mol
46g / mol
18g
Mole of H2O = = 1mol
18g / mol
Total mole = mole of C2H5OH + mole of H2O = 0.5mol + 1mol =1.5mol

mole of ethanol 0.5mole

Mole fraction of ethanol (C2H5OH) = = = 0.3
Total mol 1.5mol

 Mole fraction is unit less.

 The total sum of mole fractions is unity (one)

2. A sample of rubbing alcohol contains 142 g of isopropyl alcohol (C3H7OH) and 58.0
g of water. What are the mole fractions of alcohol and water?
 Solution
First find the number of moles of C3H7OH and the number of moles of H2O:
𝑚𝑎𝑠𝑠 𝑜𝑓 C3H7OH
Moles of C3H7OH =
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 C3H7OH

= 142 g/60.06g/mol = 2.36 mole

𝑚𝑎𝑠𝑠 𝑜𝑓 water
Moles of water = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 water

= 18.09g /18.02 g/mole = 3.22 mole

Calculating mole fractions:
Number of moles of isopropyl alcohol
Mole fraction of isopropyl alcohol (X) = moles of isopropyl alchol+ moles of water

= 2.36 mole/(2.36 mole + 3.22 mole)

= 0.423
Number of moles of water
Mole fraction of isopropyl alcohol (XH2O) = moles of isopropyl alchol+ moles of water

= 3.22 mole/(2.36 mole + 3.22 mole)

= 0.577
3. A solution was prepared by adding 48 g of methanol into 81 g of water (H 2O). What
is the mole fraction of methanol in this solution? (UEE, 2005)
(Answer 0.25)

Molarity (M)
Molarity or molar concentration is the number of moles of solute in 1 L of solution.
Molarity is determined by the equation

Thus moles of solute needed to prepare a solution of desired concentration can

theoretically be calculated by rearranging the above equation as:

Example 1: A solution is prepared by dissolving 1.26g of AgNO3 in a 250 mL

volumetric flask and diluting to volume.
a) Calculate the molarity of silver nitrate solution.
b) How many millimoles of AgNO3 were dissolved?
 b) millimoles = 0.0297 mmol/mL x 250 mL = 7.42 mmol.
Example 2: How do you prepare 500 mL of 0.75M solution of Cu(NO3)2?
Solution: We have to calculate the number of moles or grams of Cu(NO3)2 needed for
dissolving and prepare the required solution.
Volume of the solution = 500 mL, Molarity = 0.75M, Amount of Cu(NO3)2 = ?
Moles of Cu(NO3)2 = molarity x volume of solution
= 0.75M x 0.5L
= 0.375 mol
Mass of Cu(NO3)2 = 0.375 mol x 187.5g/mol = 70.31g
Thus 500 mL of 0.75M of Cu(NO3)2 solution can be prepared by dissolving about
70.31g of the salt In sufficient water and the diluting it to 500 mL

NORMALITY (N)
The normality of a solution (N) is the number of equivalents of the solute contained in
one liter of solution. A one normal solution contains one equivalent per liter.

Example
1. What is the normality of a solution made by dissolving 5 g of H2SO4 in enough water
to make 200 ml of solution?
Solution
Since H2SO4 has two ionizable hydrogen then, assuming total dissociation, the equivalent
weight of H2SO4 is:
Formula weight of H 2 SO4
Equivalent weight of H2SO4 =
Total positivevalence
98g / mol
= = 49g/eq
2mol / eq
Given mass
Normality =
Equivalentwt. x volume
 5g
Normality = x 0.2 L
49 g / eq
Normality = 0.51N H2SO4
2. Calculate the normality of a solution of HCl containing 2.2g of Hydrogen chloride
in 200mL.
Solution: Volume of the solution = 200mL = 0.20L
Number of equivalents of HCl = actual mass/equivalent wt of HCl
= 2.2g/36.5g/equivalent
= 0.06 equivalent
Normality = no of equivalents/Liters of solution
= 0.06equivalents/0.20L
= 0.3 equivalents/L = 0.3N
3. Iodine (I2) is an oxidizing agent that in reactions with reducing agents is reduced to
iodine ion (I-). How many grams of I2 would you weigh out to prepare 100mL of a
0.100N I2 solution?
Solution: Each molecule of I2 consumes 2 electrons. I2 + 2e- → 2I-
Equivalent wt of I2 = molar mass of I2 /number of equivalents per mole
= 254g/mol /2eq/mol = 127g/eq
From Normality = (actual mass/ eq. wt)/liters of solution
Actual mass = Normality x eq. wt x liter of solution
= 0.100N x127g/eq x 0.100L = 1.27g
Molality (m)
The molality of a solution is the number of moles of solute per kilogram of solvent
contained in a solution.
Number of moles of solute
Molality (m) = Mass (Kg)of solvent

Or
mass of solute
Molality = Molar mass of solute × Mass (Kg) of solventg)of solvent

A solution is prepared by dissolving 17.2 g of ethylene glycol (C2H6O2) in 0.500 kg of

water. If the final volume of the solution is 515 mL, calculate concen-tration of the
solution in: a) molarity b) molality c) percent by mass d) mole fraction, and e) mole
percent.

Solution:
Given: Mass of solute = 17.2 g
Mass of solvent = 0.500Kg
Volume of solution = 515 mL = 0.515 L
Molar mass of (C2H6O2) = 62 g/mole
Required: Molarity?, molality?, mass percent? Mole fraction? Mole percent?
a) We can calculate molarity using equation
actual mass ( g )
Molarity 
Molar mass ( g / mol ) x volume of solution ( L)
17.2 g
M 
62 g / mol x 0.515 L
 0.539mol / LorM
b) For molality we will use equation 2.3.8 or 2.3.9. Equation 2.3.8 is a two step
process but let’s use the second one for direct substitution.

m (g)
m
M ( g / mol ) x Kg ( solvent)
17.2 g

62 g / mol x 0.500 Kg
 0.554 molal

c) To calculate mass percent, we use equation

mass of solute
mass percent  x100%
mass of soltution
17.20
 x100%
17.70
 3.33%
d) To determine the mole fraction of the solute we will use equation

Since the solution contains only two components, solute (glycerol) and solvent (water),
we must first determine the number of moles of these two components.
 actual mass ( g )
mole of C 2 H 6 O 
Molar mass ( g / mol )
17.2 g
n
62 ( g / mol )
 0.28 mol
actual mass ( g )
mole of H 2 O 
Molar mass ( g / mol )
500 g
n
18 ( g / mol )
 27.78 mol

Thus equation 2.3.9 will reduce to:

mole of C 2 H 6 O2
xC2 H 6O2 
mole of C 2 H 6 O2  mole of H 2 O
0.28

0.28  27.78
 0.00998

e) The percentage of mole fraction of the solute in the solution is 0.998%.

Exercise 4.6
Instruction: Choose the correct answer from the alternatives given and circle the
letter of your choice.
1. How many moles are there in 159 g of alanine, C3H7NO2? (UEE, 2002)
A) 0.560 B) 0.992 C) 1.78 D) 3.31
2. What is the molarity of a solution obtained by dissolving 0.01 moles of NaCl in 500
mL of solution? (UEE 2002)
A) 0.01 M B) 0.005 M C) 0.02 M D) 0.10 M
3. How many grams of Ca(OH)2 are contained in 1500 mL of 0.0250 M Ca(OH)2
solution? (UEE, 2002)
A) 1.85 g B) 2.78 g C) 3.17 g D) 4.25 g
4. What is the molarity of a solution containing 10 g of sulfuric acid in 500 ml of
solution? (UEE, 2004)
A) 0.02 B) 0.03 C) 0.12 D) 0.20
5. How many moles of sodium hydroxide are present in 2.5 L of 0.5 M aqueous
solution? (UEE, 2006)
A) 0.2 B) 0.5 C) 1.25 D) 12.5
6. What is the molarity of a 5 g hydrogen peroxide (H2O2) in 100 mL solution that is
used for hair bleaching? (UEE, 2007)
A) 0.015 M B) 0.15 M C) 1.5 M D) 3 M
7. What is the final concentration of Cl− ion when 250 mL of 0.20 M CaCl2 solution is
mixed with 250 mL of 0.40 M KCl solution? (UEE, 2007)
A) 1.60 M B) 0.40 M C) 0.20 M D) 0.60 M
8. The concentration of nitrate ion in a solution that contains 0.900 M aluminum nitrate
is (UEE, 2007)
A) 0.90 M B) 0.45 M C) 0.30 M D)
2.70 M
9. How many moles of H2SO4 are needed to prepare 5.0 L of a 2.0 M solution of
H2SO4? (UEE, 2008)
A) 2.5 B) 5.0 C) 20 D) 10
10. What is the molarity of a solution made by dissolving 10 g of glucose (C6H12O6) in
sufficient water to form 200 mL solution? (UEE, 2008)
A) 0.18 B) 0.251 C) 0.362 D) 0.278
11. Which one of the following solutions has the greatest concentration of Na+ ions?
(2000)
A) 0.026 M Na2CO3 B) 0.014 M NaCl C) 0.013 M Na3PO4 D) 0.032
M NaNO3
12. What is the normality of a 2.3 M sulphuric acid solution? (UEE, 2001)
A) 1.15 N B) 2.3 N C) 4.6 N D) 6.9 N
13. What is the normality of 1.0 M solution of Na2CO3? (UEE, 2002)
A) 1 N B) 0.5 N C) 2 N D) 3 N
14. What is the mole fraction of a solute in one molal aqueous solution? (UEE, 2003)
 A) 0.009 B) 0.018 C) 0.027 D) 0.036
15. An aqueous solution is 70 % nitric acid (HNO3) by mass. What is the concentration of
HNO3 expressed in molality? (UEE, 2006)
A) 0.55 gm B) 8.62 m C) 11.1 m D) 37 m
Workout
16. 5.85 g of sodium chloride (NaCl) is dissolved in 250 mL of solution. Calculate
a) the molarity of the solution
b) the mass percentage of the solute
17. How would 250 ml of 0.15 M KNO3 solution be prepared?
18. What is the molarity of each of the following solutions:
(a) 15.0 g Al2(SO4)3 in 0.250 mL solution,

(b) 5.25 g Mn(NO3)2.2 H2O in 175 mL of solution,

4.7 PREPARATION OF SOLUTIONS

A common task in school, medical, industrial and other chemical laboratories is the
preparation of solutions of known concentrations.
Solutions are usually prepared from solutes of liquids or solids. Occasionally they are
prepared from gases.

Short procedures for preparation of solutions from solids

 First, the solute is accurately weighed and transferred to a volumetric flask,
 Second water is added through a funnel,
 Next, the solid is slowly dissolved by gently swirling the flask,
 After all the solid has dissolved, more water is slowly added to bring the level of
solution exactly to the volume mark.
Activity
Dissolve 1g each of NaCl, CuSO4, and KMnO4 in different 50 mL measuring flasks. Note the
colour intensity of the solutions. Take 25 mL of solution from each flask and further dilute each
solution to 50 mL. Note the colour intensity again.
1. What did you observe?
2. Can you correlate the colour intensity with the concentration of the solution?
3. Does this analogy works for colorless solution also?
Discuss your observations with your classmates.

Preparation of solutions from liquids - Dilution

 Dilution is the procedure for preparing a less-concentrated solution from a more
concentrated one.
In carrying out a dilution process, Number of moles of solute before dilution = Number
of moles of solute after dilution MiVi=MfVf Or, more generally, CiVi=CfVf
Where Ci is initial concentration, Cf is final concentration, Vi is initial volume and Vf is
final volume
Example
1. Describe how you would prepare 750 mL of a 0.015 M KBr solution from solid KBr.
Solution
Mass of KBr = Molarity x Molar mass of KBr x Volume of solution
= 0.015 M x 119 g/mol x 0.75 L
= 1.34 g
Weigh out 1.34 g KBr, dissolve in water, dilute with sufficient water by stirring to give a
final volume of 750 mL.
2. Describe how you would prepare 500 mL of a 1.75 M H2SO4 solution, starting with
an 8.61 M stock solution of H2SO4
Solution:
MiVi = MfVf
Vf = MfVfMi
=1.75 M×500 mL8.61 M=102 mL
Dilute 102 mL of 8.61 M H2SO4 solution with sufficient water to give a final volume of
500 mL.
Activity
1. How do you prepare 60.0 mL of 0.2 M HNO3 from a stock solution of 4 M HNO3?
2. You have 100 mL of a 0.5 M HCl solution, and you want to dilute it to exactly 0.1 M.
How much water should you add?

Exercise 4.7
Instruction: Choose the correct answer from the alternatives given and circle the
letter of your choice.
1. How would you prepare 0.90 L of 1.2×10−2M NaCl aqueous solution? (2003 E.C)
A. Weigh 0.012 g of NaCl, dissolve in a small amount of water and dilute to 0.09 L
B. Weigh 0.63 g of NaCl, dissolve in a small amount of water and dilute to 0.09 L
C. Weigh 1.2 g of NaCl, dissolve in a small amount of water and dilute to 0.09 L
D. Weigh 1.4 g of NaCl, dissolve in a small amount of water and dilute to 0.09 L
2. How much water has to be evaporated from 250 mL of 1 M Ca(OH)2 to make it 3
M? (UEE, 2004)
A) 100 mL B) 150 mL C) 167 mL D) 200 mL
3. How many mL of water is required to dilute 50 mL of 3.5 M H2SO4 to 2.00 M
H2SO4? (UEE, 2004)
A) 37.5 B) 45 C) 75 D) 87.5
4. What is the quantity of water in mL, required to prepare 0.5 M of HCl from a
concentrated solution of 3.5 M in 50 mL? (UEE, 2005)
A) 50 B) 100 C) 300 D) 350
5. How many mL conc. HNO3 and how many mL of water are required to prepare 500
mL of 0.1 M HNO3 from a conc. 13 M HNO3?
A. 1 mL HNO3 and 496.15 mL of H2O
B. 3 mL HNO3 and 500 mL H2O
C. 3.85 mL HNO3 and 500 mL H2O
D. 3.85 mL HNO3 and 496.15 mL H2O
 6. A lab instructor is preparing 5.0 L of a 0.10 M Pb(NO3)2 (molecular mass = 331)
solution. What is the mass required/ (UEE, 2006)
A. 165.5 g Pb(NO3)2 and add 5.0 Kg of H2O
B. 165.5 g of Pb(NO3)2 and add H2O until the solution has a volume of 5.0 L
C. 31.1 g of Pb(NO3)2 and add H2O until the solution has a volume of 5.0 L
D. 31.1 g Pb(NO3)2 and add 5.0 L of H2O
7. If a student wishes to prepare approximately 100 ml of an aqueous solution of 6 M
HCl using 12 M HCl. Which procedure is correct? (UEE, 2007)
A. Adding 50 mL of 12 M HCl to 50 mL of water while stirring the mixture steadily.
B. Adding 25 mL of 12 M HCl to 50 mL of water while stirring the mixture steadily.
C. Adding 50 mL of water to 50 mL of 12 M HCl while stirring the mixture steadily.
D. Adding 25 mL of water to 50 mL of 12 M HCl while stirring the mixture steadily.
8. What is the mass of CuCl2.6H2O required to prepare 300 mL of 0.1 M aqueous
solution? (UEE 2000)
A) 1.19 g B) 2.19 g C) 7.28 g D) 11.9 g

Answer key
Exercise 4.1
1. B 2. B 3. A 4.D 5. A
Exercise 4.2
1. D 2. B 3. C 4. A
Exercise 4.3
1. C
2. a) saturated b) supersaturated c) unsaturated
3. By adding a crystal of X in each solution. The size of the crystal decreases, remains
the same and increases when added into unsaturated, saturated and super saturated
solutions, respectively.
Exercise 4. 4
1. A 2. B C. D 4. D
Exercise 4.6
1. D 2. C 3. B 4. D 5. C 6, C 7. B 8. D
 9. D 10.D 11. A 12. C 13. C 14.B 15. D
Exercise 4. 7
1. B 2. C 3. A 4. C 5. D 6. B 7. A 9. C
Chapter 6

Chemical equilibrium

6.1 introductions

A reversible reaction is a chemical reaction that results in an equilibrium mixture of

reactants and products.

H2(g) + I2(g) 2HI(g)

Chemical equilibrium is the state of the reaction when the macroscopic properties like
temperature, pressure, volume and concentration of the reaction do not change with time.

When a chemical reaction takes place, the reactants are consumed and the products are
formed. As a result the concentrations of reactants decrease with time and that of
products increase. If the reaction is reversible in nature and it takes place in a closed
vessel then the products decompose to give back reactants.

Fig 6.1 Change in the rate of reaction with time for the forward and
reverse reactions.

The chemical equilibrium is also dynamic in nature since the reaction does not stop at
equilibrium. The macroscopic properties remain constant because the rate of forward
reaction becomes equal to the rate of reverse reaction and hence no net reaction takes
place.

The general equation that represents a reversible reaction is

------(1)

For the forward reaction A and B are reactants and M and N are products. It is assumed
that the reaction is homogeneous in nature, that is all the reactants and the products are
present in the same state. The rate of forward reaction (rf ) depends upon the
concentrations of A and B and is given by the expression

rf = kf [A]a [B]b

The rate of reverse reaction is given by the expression

rb = kb [M]m [N]n

where [A], [B], [M] and [N] are the concentrations of A, B, M and N respectively. k f and
kb are rate constants for forward and reverse reactions, respectively.

At equilibrium, the rate of forward reaction is equal to the rate of reverse reaction.

rf = rb or kf [A]a [B]b = kb [M]m [N]n

kf and kb are constant. Therefore, the ratio of kf to kb is also a constant.

The ratio of kf and kb is represented by Keq. Since this constant represents the reaction at
equilibrium it is known as equilibrium constant. [M], [N], [A] and [B] are concentrations
at equilibrium. When the concentrations are expressed in molarities, the equilibrium
constant is represented by Kc .
When all the reactants and the products are in the gaseous state, their concentrations can
be written in terms of partial pressures. In such cases the Keq is denoted by Kp. The
expression for Kp is

Units of Kc from rxn (1) will be

Unit of Kp is decided by the unit of pressure. If the partial pressure is expressed in atm,
then the unit of Kp , will be:

Example:

Write the concentration-based expression for each of the following reactions:

a) 3O2(g)⇌2O3(g)

b) N2(g)+3H2(g)⇌2NH3(g)

[𝑂3 ]2 [𝑁𝐻3 ]2
Solution a) 𝐾𝑐 = [𝑂2 ]3
b) 𝐾𝑐 = [𝑁 3
2 ][𝐻2 ]

2SO2(g)+O2(g)⇌2SO3(g), the concentrations at equilibrium are [SO2] = 0.90 M, [O2] =

0.35 M, and [SO3] = 1.1 M. What is the value of the equilibrium constant, Kc?

Solution:

(𝑆𝑂3 )2
𝐾𝑐 = 𝑄𝑐 =
(𝑆𝑂2 )2 (𝑂2 )1
 (1.1)2
=
(0.9)2𝑥(𝑂2)

= 𝟒. 𝟐𝟕

6.2 Characteristics of chemical equilibrium

The status of a reversible reaction is conveniently assessed by evaluating its reaction

quotient (Q). For a reversible reaction described by

𝒎𝑨 + 𝒏𝑩+⇌ 𝒙𝑪 + 𝒚𝑫

The reaction quotient is derived directly from the stoichiometry of the balanced equation
as

[𝑪]𝒙 [𝑫]𝒚
𝑸𝒄 =
[𝑨]𝒎 [𝑩]𝒏

The constant value of Q exhibited by a system at equilibrium is called the equilibrium

constant, K: 𝑲 = 𝑸 , at equilibrium. A reaction exhibiting a large K will reach
equilibrium when most of the reactant has been converted to product, whereas a small K
indicates the reaction achieves equilibrium after very little reactant has been converted.
The magnitude of K does not indicate how rapidly or slowly equilibrium will be reached.

The value of equilibrium constant predicts the extent of reaction at equilibrium. It also
tells about the relative concentrations of products and reactants present at equilibrium,
that is, the position of equilibrium. Kc can have three types of values.

When Kc >1: This indicates that the value of numerator is greater than the denominator.
So the formation of products is favoured at equilibrium.

When Kc < 1: This indicates that the value of numerator is less than the denominator. So
the formation of products is not favoured at equilibrium.

When Kc =1: This indicates that the rate of forward reaction is equal to the rate of
backward reaction.
When Q is less than K, then the reaction will proceed in the forward direction and more
products will be formed till the equilibrium is reached. When Q is greater than K, then
the reaction will proceed in the reverse direction and more reactants will be formed till
the equilibrium is reached. When Q is equal to K, then the reaction has attained
equilibrium.

Calculating the Equilibrium Concentrations:

Not only can we estimate the extent of reaction from the Kc value, but also the expected
concentrations at equilibrium can be calculated from knowledge of the initial
concentrations and the Kc value. In these types of problem it will be very helpful to use
the following approach.

1. Express the equilibrium concentrations of all the species in terms of the initial
concentrations and an unknown x, which represents the change in concentration.

2. Substitute the equilibrium concentrations derived in part 1 into the equilibrium

constant expression, and solve for x. The equilibrium concentration is given by:
equilibrium concentration = initial concentration ± the change due to the reaction where
the + sign is used for a product, and the – sign for a reactant.

3. Use x to calculate the equilibrium concentration of all the species.

Example:

Calculate the pH of a 0.50 M HF solution at 25°C. The ionization of HF is given by

HF (aq) + H2O (l) H3O+(aq) + F–(aq) is 6.8x10-4

Solution:

The species that can affect the pH of the solution are HF, and the conjugate base F–, Let x
be the equilibrium concentration of H3O+ and F- ions in molarity (M). Thus,
 Rearranging this equation gives

x2 = (0.5)(6.8 × 10–4) = 3.4 × 10–4

x = ∫ 3.4× 10–4 = 1.8 × 10–2 M Thus, we have solved for x without using the quadratic
equation. At equilibrium, we have

[HF] = (0.50 – 0.018) M = 0.48 M, [H3O+] = 0.018 M [F–] = 0.018 M and the pH
of the solution is pH = –log(0.018) = 1.74

The equilibrium constant depends upon the following factors:

1. Temperature at which the experiment is performed.

2. The form of equations which describe the equilibrium.

For example, for the reaction N2(g) + 3H2(g) → 2NH3(g)

If the reaction is multiplied by 2,

2N2(g) + 6H2(g) → 4NH3(g) the expression for Kc becomes

If the reaction is reversed, 2NH3(g) → N2(g) + 3H2(g) the Kc becomes

The equilibrium constant does not depend upon the initial concentrations of the reactants.
Kc and K p are independent of pressure.

Homogeneous Equilibria

A homogeneous equilibrium is one in which all reactants and products are present in the
same phase. Homogeneous equilibria take place in solutions.

These solutions are most commonly either liquid or gaseous phases, as shown by the
examples below:

[𝑪𝟐 𝑯𝟐 𝑩𝒓𝟐 ]
𝑪𝟐 𝑯𝟐 (𝒂𝒒) + 𝟐𝑩𝒓𝟐 (𝒂𝒒) ⇌ 𝑪𝟐 𝑯𝟐 𝑩𝒓𝟐 (𝒂𝒒) 𝑲𝑪 =
[𝑪𝟐 𝑯𝟐 ][𝑩𝒓𝟐 ]𝟐

[𝑰𝟑 − ]
𝑰𝟐 (𝒂𝒒) + 𝑰− (𝒂𝒒) ⇌ 𝑰𝟑 − (𝒂𝒒) 𝑲𝑪 =
[𝑰𝟐 ][𝑰− ]

Water is not included in the reaction quotient in the second reaction. The
reason for this omission is related to the relative concentrations for liquids
and solids are equal to 1 and needn’t be included. The equilibria below all
involve gas-phase solutions:

[𝑪𝟐 𝑯𝟒 ][𝑯𝟐 ]
𝑪𝟐 𝑯𝟔 (𝒈) ⇌ 𝑪𝟐 𝑯𝟒 (𝒈) + (𝒈)𝑯𝟐 𝑲𝑪 =
[𝑪𝟐 𝑯𝟔 ]

For gas-phase solutions, the equilibrium constant may be expressed in terms

of either the molar concentrations (Kc) or partial pressures (Kp). A relation
between these two K values may be simply derived from the ideal gas
equation and the definition of molarity:

𝑷𝑽 = 𝒏𝑹𝑻
 𝒏
𝑷 = ( ) 𝑹𝑻 = 𝑴𝑹𝑻
𝑽

For the gas-phase reaction

𝒎𝑨 + 𝒏𝑩 ⇌ 𝒙𝑪 + 𝒚𝑫:

(𝑷𝑪 )𝒙 (𝑷𝑫 )𝒚
𝑲𝑷 =
(𝑷𝑨 )𝒎 (𝑷𝑩 )𝒏

([𝑪]𝒙𝑹𝑻)𝒙 ([𝑫]𝒙𝑹𝑻)𝒚
=
([𝑨]𝒙𝑹𝑻)𝒎 ([𝑩𝒙𝑹𝑻)𝒏

[𝑪]𝒙 [𝑫]𝒚 (𝑹𝑻)𝒙+𝒚

= 𝒙
[𝑨]𝒎 [𝑩]𝒏 (𝑹𝑻)𝒎+𝒏

= 𝑲𝑪 (𝑹𝑻)(𝒙+𝒚)−(𝒎+𝒏) = 𝑲𝑪 (𝑹𝑻)∆𝒏

And so, the relationship between 𝐾𝐶 and 𝐾𝑃 is

𝑲𝑷 = 𝑲𝑪 (𝑹𝑻)∆𝒏

Where 𝛥𝑛 is the difference in the molar amounts of product and reactant gases, in this
case:

𝜟𝒏 = (𝒙 + 𝒚) − (𝒎 + 𝒏)

Example:

Write the equations relating Kc to KP and calculate KP for each of the following reactions
Kc is equal to 0.28 at 900 °C:

𝑪𝑺𝟐 (𝒈) + 𝟒𝑯𝟐 (𝒈) ⇌ 𝑪𝑯𝟒 (𝒈) + 𝟐𝑯𝟐 𝑺(𝒈)

0.28
𝐾𝑃 = 𝐾𝐶 (𝑅𝑇)Δn = = 3.0𝑋10−5
[(0.0821)(1173)]2

Heterogeneous Equilibria
  A heterogeneous equilibrium involves reactants and products in two or more
different phases:

𝑷𝒃𝑪𝒍𝟐 (𝒔) ⇌ 𝑷𝒃𝟐+ (𝒂𝒒) + 𝟐𝑪𝒍− (𝒂𝒒) 𝑲𝑪 = [𝑷𝒃𝟐+ ][𝑪𝒍− ]𝟐

𝟏
𝑪𝒂𝑶(𝒔) + 𝑪𝑶𝟐 (𝒈) ⇌ 𝑪𝒂𝐂𝑶𝟑 (𝒔) 𝑲𝑪 =
[𝑪𝑶𝟐 ]

[𝑪𝑺𝟐 ]
𝑪(𝒔) + 𝟐𝑺(𝒈) ⇌ 𝑪𝑺𝟐 (𝒈) 𝑲𝑪 =
[𝑺]𝟐

6.3The Le Chatelier’s principle

Châtelier’s principle (T,P,V&C) if an equilibrium system is stressed, the system will

experience a shift in response to the stress that re-establishes equilibrium. It gives the
effect of any one or more of the reaction parameters namely, temperature, pressure or
concentration on equilibrium.

Le Chatelier’s principle describes the effect of the change in parameters on the position
of equilibrium. That is, it predicts whether the changes in reaction parameters will favor
the formation of reactants or products.

Effect of Change in Temperature on the Position of Equilibrium

If the temperature of the reaction at equilibrium is increased then the reaction will
proceed in that direction where heat is absorbed so as to undo the effect of heating.
Similarly if the temperature of the reaction at equilibrium is lowered then the reaction
will proceed in that direction where the heat is produced so that the equilibrium is again
attained.

𝒉𝒆𝒂𝒕 + 𝑵𝟐 𝑶𝟒 (𝒈) ⇌ 𝟐𝑵𝑶𝟐 (𝒈)

Raising the temperature of the system is causing to increasing the amount of a product,
and so the equilibrium will shift to the right. Lowering the system temperature will
likewise cause the equilibrium to shift left.
For exothermic processes, heat is viewed as a product of the reaction and so the opposite
temperature dependence is observed. A +B →C + heat, when the temperature is lowered,
more products are formed. That is, the decrease in temperature favors exothermic
reactions.

Effect of Change in Pressure on the Position of Equilibrium

If the pressure at equilibrium is increased then the reaction will proceed in that direction
where the pressure is reduced. Since the pressure depends upon the number of moles, on
increasing the pressure the reaction will precede in that direction where the number of
moles are reduced.

For a general reaction,

𝒂𝑨 + 𝒃𝑩 ⇌ 𝒎𝑴 + 𝒏𝑵

The effect of pressure is decided by n. 𝒏 = (𝒎 + 𝒏) − (𝒂 + 𝒃)

 If n > 0, that means the total moles of products is greater than the total moles of
reactants.

 Lowering of pressure will favor the reaction in reverse direction.

 If n < 0, that means the total moles of products is less than the total moles of
reactants.

 Increasing the pressure will favor the reaction in forward direction.

 If n = 0, then the change in pressure has no effect on the position of equilibrium.

Effect of Change in the Concentration on the Position of Equilibrium

At equilibrium, on increasing the amount of a substance, the reaction proceeds in that

direction where the substance is consumed. When the amount of the reactants is
increased the reaction proceeds in the forward direction and when the amount of
products is increased, the reaction proceeds in the backward direction. The forward
reaction is favored if the products formed are removed from the vessel.
 Effect of a Catalyst

Catalysts are substances that enable a reaction to proceed via a different mechanism with
an accelerated rate. The catalyzed reaction mechanism involves a lower energy transition
state than the uncatalyzed reaction, resulting in lower activation energy, Ea Consequently,
both forward and reverse reactions are accelerated, and equilibrium is achieved more
quickly but without a change in the equilibrium constant.

6.3.1. Calculation of equilibrium constant

Iodine molecules react reversibly with iodine ions to produce tri iodide ions.

𝐼2 (𝑎𝑞) + 𝐼 − (𝑎𝑞) ⇌ 𝐼3− (𝑎𝑞)

If a solution with the concentrations of 𝐼2 and 𝐼 − both equal to 1.000x10-3M

before reaction gives an equilibrium concentration of I2 of 6.61x10-4M, what is
the equilibrium constant for the reaction?

Solution: To calculate the equilibrium constants, constant are needed for all the
reactant and products:

[𝑰𝟑− ]
𝑲𝒄 =
[𝑰𝟐 ][𝑰− ]

 The initial concentrations of the reactants and the equilibrium concentration of the
product.
Use this information to derive terms for the equilibrium concentrations of
reactants, presenting all the information in the reaction table.

𝑰𝟐 (𝒂𝒒) + 𝑰− (𝒂𝒒) ⇌ 𝑰−
𝟑 (𝒂𝒒)

Initial concentration (M) 1.000x10-3 1.000x10-3 0

 Change (M) -x -x +x

Equilibrium conc.(M) 1.000x10-3-x 1.000x10-3-x X

At equilibrium the concentration of 𝐼2 is 6.61x10-4M so that

1.000x10-3-x =6.61x10-4

X = 1.000x10-3-6.61x10-4 = 3.39x10-4M

𝑰𝟐 (𝒂𝒒) + 𝑰− (𝒂𝒒) ⇌ 𝑰−
𝟑 (𝒂𝒒)

Initial concentration (M) 1.000x10-3 1x10-3 0

Change (M) -3.39x10-4 -3.39x10-4 +3.39x10-4

Equilibrium conc.(M) 6.61x10-4 6.61x10-4 3.39x10-4

Substitute the equilibrium concentrations into the K expression and solve:

[𝑰𝟑− ]
𝑲𝒄 =
[𝑰𝟐 ][𝑰− ]

3.39𝑥10−4 𝑀
= = 776
(6.61𝑥10−4 𝑀)(6.61𝑥10−4 𝑀)

6.4 IONIC EQUILIBRIA

6.4.1 Acid-Base Equilibria

6.4.1.1. Theories of Acids and Bases

I. Arrhenius (Classical) Acid-Base Definition

An acid is a substance that dissociates in water to yield a H3O+ and a base is a

–
substance that dissociates in water to yield OH
Neutralization is the reaction of an H+ (H3O+) ion from the

Eg. Acids: HCl, H2SO4, HNO3 etc

Bases: NaOH, KOH, Mg(OH)2 etc.

 Three limitations:

1. Applied only to substances in aqueous solutions.

2. Restricted to substances that produce H+ and OH– ions

3. Does not explain why some compounds containing hydrogen such as CH4 dissolve
in water and do not give acidic solutions

II. Brønsted-Lowry Acid-Base Definition

 An acid is a proton donor, any species that donates an H+ ion.

 It must contain H in its formula.

Eg. HCl, HNO3, H2PO4- etc.

 All Arrhenius acids are Brønsted-Lowry acids.

 A base is a proton acceptor, any species that accepts an H+ ion.

 It must contain a lone pair of electrons to bind the H+ ion.

Eg. NH3, CO32-, F -, OH – etc.

 Brønsted-Lowry bases are not Arrhenius bases, but all Arrhenius bases contain
the Brønsted-Lowry base OH-.

 Therefore in the Brønsted-Lowry perspective, one species donates a proton and

another species accepts it and hence, an acid-base reaction is a proton transfer
process.

 Based on the number of proton donated acids can be classified into 3:

i. Monoprotic acid: yields one hydrogen ion upon ionization

Eg. HNO3(aq) →H+(aq) + NO3-(aq)

Diprotic acid: give up two H+ ions, in two separate steps:

Eg.

Triprotic acids: yield three H+ ions, are relatively few in number.

Eg.

Conjugate Acid-Base Pairs

The concept of the conjugate acid-base pair, arises from the Brønsted definition of
acids and bases.

The conjugate base of a Brønsted acid is the species that remains after the acid
donates a proton.

Conversely, a conjugate acid results from the addition of a proton to a Brønsted

base.

Every Brønsted acid has a conjugate base, and vice versa.

Example:

 weak acids have strong conjugate bases

 weak bases have strong conjugate acids

Lewis Acid-Base Definition

A base is any species that donates an electron pair. Whereas an acid is any species that
accepts an electron pair. Protons act as Lewis acids in that they accept an electron

B: + H+ ⇌ B H+

A Lewis base has a lone pair of electrons to donate. And a Lewis acid is a vacant orbital

6.4.1.2 Dissociation of strong and weak

monoprotic acids and bases

Strong and Weak Acids: The reaction of an acid with its solvent (typically water) is
called an acid dissociation reaction. Acids are divided into two categories based on the
ease with which they can donate protons to the solvent.

 A strong acid is completely dissociated in aqueous solution.

 They almost completely transfer their protons to the solvent molecules.

 That is, the equilibrium constants for the following reactions are very large:

HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq)

 Weak acids, of which aqueous acetic acid is one example, cannot completely
donate their acidic protons to the solvent.

 Instead, most of the acid remains undissociated, with only a small fraction present
as the conjugate base.

CH3COOH(aq) + H2O(l) ⇌ H3O+(aq) + CH3COO–(aq)

The Acid-Dissociation Constant (Ka)

Strong acids dissociate completely into ions in water:

In a dilute solution of a strong acid, almost no HA molecules exist:

[H3O+] = [HA]init or [HA]eq = 0

Weak acids dissociate very slightly into ions in water:
+ -
HA +H O HO +A
(aq) 2 (aq) 3 (aq) (aq)

In a dilute solution of a weak acid, the great majority of HA

molecules are undissociated: [H3O+] << [HA]init or [HA]eq = [HA]init

at equilibrium, Q = K << 1
c c

Strong and Weak Bases: the basicity of an aqueous solution

 is a measure of the concentration of the hydroxide ion, OH–. The most common
example of a strong base is an alkali metal hydroxide, which completely
dissociates to produce the hydroxide ion. NaOH(aq) → Na+(aq) + OH–(aq)

 Weak bases only partially accept protons from the solvent and are characterized
by a base dissociation constant, Kb.

 For example, the base dissociation reaction and base dissociation constant for the
acetate ion are:

CH3COO-(aq) + H2O(l) ⇌ H3O+(aq) + CH3COOH (aq)

[𝐶𝐻3 𝐶𝑂𝑂𝐻][𝑂𝐻 − ]
Kb= [𝐶𝐻3 𝐶𝑂𝑂− ]
Amphiprotic Species: Some species can behave as either an acid or a base.
 A species that can serve as both a proton donor and a proton acceptor is called
amphiprotic.

HCO3–(aq) + H2O(l) ⇌ H3O+(aq) + CO32–(aq) ...... Ka2= 4.69×10–11

HCO3–(aq) + H2O(l) ⇌ OH–(aq) + H2CO3(aq) ......Kb2= 2.25×10–8

 Whether an amphiprotic species behaves as an acid or as a base depends on the

equilibrium constants for the two competing reactions.

Since, kb2> ka2, we expect that aqueous solutions of HCO3–will be basic.

6.4.1.3 Dissociation of water and pH of aqueous solutions

o Water is an amphiprotic solvent in that it can serve as an acid or a

base.

o An interesting feature of an amphiprotic solvent is that it is capable of

reacting with itself as an acid and a base.

o For pure water the concentration of hydroxyl and hydronium ions must be equal:

The ion-product for water, Kw:

2 + - -14
K [H O] = K = [H O ][OH ] = 1.0 x 10 (at 25°C)
c 2 w 3

[H3O+] = [OH-] = √1.0 x 10-14 = 1.0 x 10 -7M (at 25°C)

pH: a way to express acidity -- the concentration of H+ in solution

 pH = log (1/ [H+]) = - log [H+]=- log [H3O+]

 Eg. What is the pH of a solution that is 10-12 M in hydronium ion?

Solution: pH = -log[H3O+] = (-1)log 10-12 = (-1)(-12) = 12

6.4.1.4 Common ion effect

 The shift in the position of anequilibrium on addition of a substance that provides

an ion in common with one of the ions already involved in the equilibrium.

 Equilibrium is shifted in the direction that reduces the concentration of the

common ion

 Sources of common ions are

 Strong acid
 Strong bases
 Soluble salts

 The Common Ion Effect is important to

 Buffers

 Solubility Equilibrium

6.4.1.5. Buffer solutions

 A buffer solution resists changes in pH when it is diluted or when acids or bases

are added to it.

Generally, buffer solutions are prepared from a conjugate acid/base pair, such as
acetic acid/sodium acetate or ammonium chloride/ammonia.

 Has two components, one component is able to neutralize the acid and other
component is able to neutralize base
  Can be prepared from
 weak acid and its conjugate base Eg CH3COOH/CH3COO-
 weak base and its conjugate acid eg NH3/NH4+
 weak acid and strog base, if the weak acid is excess eg. CH3COOH/NaOH
in 2:1mole ratio
 weak base and strog acid, if the weak base is excess eg. NH3/HCl in
2:1mole ratio
 it is used to:
 human blood
 protein studies often must be performed in buffer media
 control of PH is often important in industrial process
 enzyme activity always working with PH 6&8

for a weak acid HA the dissociation equation and Ka expression are given by

HA +H2O H3O+ + A-

Ka =[ H3O+][ A-]/[HA], [ H3O+] = Ka.[HA]/[A-]

Put logarithm on both sides of the equation and rearranging

[𝐻𝐴] [𝑐𝑜𝑛𝑗.𝑏𝑎𝑠𝑒]
PH = Ka .log [𝐴−] for any acid- base pair PH = PKa + log [𝑎𝑐𝑖𝑑]

[𝐵] [𝑐𝑜𝑛𝑗. 𝑎𝑐𝑖𝑑]

For bases POH =Kb.log[𝐵𝐻+] POH = PKb + log
[𝑏𝑎𝑠𝑒]

Example: calculate the PH of a buffer solution containing 0.1M acetic acid and 0.2M
potassium acetate.

Solution:

[𝑐𝑜𝑛𝑗.𝑏𝑎𝑠𝑒]
PH = PKa + log [𝑎𝑐𝑖𝑑]

[0.2𝑀]
PH = -logKa + log [0.1𝑀] = -Log(1.8x10-5) + log2 = 5.05

6.4.1.6. Hydrolysis of salts

• Hydrolysis is a term applied to reactions of aquated ions that change the pH from
7

• The interactions between salts and water are called hydrolysis

 There are 4 salt forming processes of acid base reactions:

1. Salt of a strong acid and a strong base:

Eg: NaCl, KBr, and Ba(NO3)2.

 Neither the cation nor anion hydrolyzes, and the solution has a pH of 7.

2. Salt of a strong acid and a weak base:

Eg: NH4Br, ZnC12, and Al(NO3)3.

 The cation hydrolyzes, forming H+ ions, and the solution has a pH less than 7.

3. Salt of a weak acid and a strong base:

Eg: NaNO2, KC2H3O2, and Ca(OCl)2.

 The anion hydrolyzes, forming OH-ions, and the solution

has a pH greater than 7.

4. Salt of a weak acid and a weak base:

Eg: NH4F, NH4C2H3O2, and Zn(NO2)2.

 Both ions hydrolyze.

The pH of the solution is determined by the relative extent to which each ion
hydrolyzes.

6.4.2 Solubility Equilibria

 Many natural processes depend on the precipitation or dissolving of a slightly
soluble salt.

 Insoluble does not mean 100% insoluble.

– In the next section, we look at the equilibria of slightly soluble, or nearly

insoluble, ionic compounds.

– Their equilibrium constants can be used to answer questions regarding solubility

and precipitation.

The Solubility Product Constant

 When an excess of a slightly soluble ionic compound is mixed with water, an

equilibrium is established b/n the solid and the ions

 The equilibrium constant for this process is called the solubility product constant,
Ksp.

2 2
CaC2O4 (s) Ca 2 (aq)  C2O4 (aq) , K sp  [Ca 2 ][ C 2 O 4 ]

 Ksp basically relationship of amount of ions in solution at saturation point

 once product of concentration of ions = Ksp, ppt occurs

 In general, Ksp is the equilibrium constant for the solubility equilibrium of a

slightly soluble ionic compound.

 It is a constant that relates to amount of substance that is dissolve in solution at

saturation point not amount in flask

 It equals the product of the equilibrium concentrations of the ions in the

compound.

 Each concentration is raised to a power equal to the number of such ions in the
formula of the compound.

Calculating Ksp from the Solubility

A 1L sample of a saturated CaC2O4 contains 0.0061g of the salt at 25°C. Calculate the
Ksp for this salt at 25°C.

o We must first convert the solubility of calcium oxalate from 0.0061 g/liter to
moles per liter.

o Note to calculate the Ksp must have data on a saturated solution, if below
saturation point can't determine Ksp.

0.0061 g CaC2O4 1 mol CaC2O4

M CaC2O4  ( )  4.8  105 mol CaC2 O4 / L
L 128g CaC2O4
When 4.8 x 10-5 mol of solid dissolve it forms 4.8 x 10-5 mol of each ion.

You can now substitute into the equilibrium-constant expression

2
K sp  [Ca 2 ][ C 2 O 4 ] K sp  (4.8 105 )(4.8 105 ) K sp  2.3 109

The common-ion effect

 Caused by the addition of a compound having an ion in common with the

dissolved substance that shift the equilibrium.
 The presence of the common ion suppresses the ionization of a weak acid or a
weak base. For example, if sodium acetate and acetic acid are dissolved in the
same solution, they both dissociate and ionize to produce CH3COO– ions

 Sodium acetate, CH3COONa, is a strong electrolyte, so it dissociates completely

in solution, but acetic acid, CH3COOH, is a weak acid and ionizes partially.
 According to Le Chatelier’s principle, the addition of CH3COO– ions from
CH3COONa to a solution of CH3COOH will suppress the ionization of
 CH3COOH and decrease the concentration of hydrogen ions. Therefore, a
solution containing both CH3COOH and CH3COONa will be less acidic than a
solution containing only CH3COOH of the same concentration.
 The shift in equilibrium of the acetic acid ionization is caused by the acetate ions
from the sodium acetate. CH3 COO– is the common ion because it is supplied by
both CH3COOH and CH3COONa.

Exercise: Determine the [H3O+] and [CH3COO–] in a solution that is 0.10 M in both
CH3COOH and HCl?

6.4.3 Factors affecting stability of complexes

The main factors affecting stability of complexes are:

 Charge on the metal ion.
 Principal quantum number.
 Nature of ligands
 Macro cyclic ligands
 Hardness and softness
 Surrounding conditions
Unit 8
8. Introduction to organic chemistry (16 hours)
8.1 Introduction
Organic Chemistry was originally known by studying the compounds of carbon, which occurred in living
organisms. However, in 1828 a German chemist, Friedrich Wohler, synthesized the organic compound
Urea, CO (NH3)2, an animal waste product, from inorganic compound ammonium cyanate, NH 4OCN.
Therefore, the meaning of Organic chemistry has to be extended to include the compounds of carbon
whether or not they occur in living organisms.

Activity 8.1
1. Do you agree with the notion that says: “carbon compounds can be
synthesized only by animals and plants”?
2. Draw diagrams to show how carbon atoms can link to one another in
different ways to form a variety of compounds by considering only
four carbon atoms.
Discuss with your group and present it to the class.

Urea was the first organic compound synthesized in the laboratory. The synthesis of urea
by Friedrich Wöhler and subsequent synthesis of other organic compounds marked the
downfall of the ‘life force’ theory.
How do you explain organic compounds at present and define organic chemistry?
The common feature of organic compounds is that they all contain the element carbon.
Organic compounds are the compounds of carbon found in and derived from plants and
animals.
They also include those substances synthesized in laboratories except the oxides of
carbon, carbonates, hydrogen carbonates, cyanides and cyanates.
Besides carbon, these compounds contain a few other elements such as hydrogen,
oxygen, nitrogen, sulphur, halogens and phosphorus. The branch of chemistry that studies
carbon compounds is called organic chemistry. This branch of chemistry was developed,
starting from the theory of ‘life force’ to the era in which synthetic organic compounds
are used in our daily lives. The number of inorganic compounds discovered and prepared
may be in the region of some hundred thousand. At present, millions of organic
compounds have been discovered, synthesized, and used.

1.2 Classification of organic compounds

Activity 8.2

CH3CH2CHO and CH3COCH3 have the same chemical formula C3H6O. Write their
detailed structures and observe their difference? Why do these compounds have
different properties? Discuss with your group and present it to the class.

Organic compounds are generally classified based on their functional group.

What are functional groups and what groups of organic compounds are known on the
basis of this classification?
The functional group is the part of a molecule or a compound that determines the
chemical properties of that molecule or compound. This group also determines some of
the physical properties of a compound. Based on the functional groups they possess, the
most common classification of organic compounds including alkanes, alkenes, alkynes,
aromatics, alcohols, aldehydes, ketones, carboxylic acids and esters. The common
functional groups of some organic compounds are given in Table 8.1.
8.3 Hydrocarbons (classification, Naming, uses, and general Properties)
Hydrocarbons are the simplest organic compounds, containing only carbon and hydrogen, they can be
straight-chain, branched chain, or cyclic molecules. Carbon tends to form four bonds in a tetrahedral
geometry. Hydrocarbon derivatives are formed when there is a substitution of a functional
group at one or more of these positions.

8.3.1 ALKANES, CYCLOALKANES AND ALKYL GROUPS

Alkanes are saturated hydrocarbons containing only carbon-carbon single bond; they are
also known as paraffin. (i.e. small tendency for a chemical reaction). Alkanes have C – C
single bond functional group. Functional group is a part of a compound that determines
its physical and chemical property. For example: C - C single bond is a functional group
of Alkanes.

Write the general formula for homologous series of alkanes.

Alkanes have the general formula CnH2n +2, where n is the number of carbon atoms.
Alkanes have endings “ane”.

The first member of alkane is methane with carbon number one. Table 8.2 shows some of
the physical states of Alkanes.

Formula Name Boiling point (0C) Density (g/cm3) State at room

 temperature

CH4 Methane -161.7 0.42 Gas

C2H6 Ethane -88.5 0.57 Gas

C3H8 Propane -4.2 0.59 Gas

C4H10 Butane -0.4 0.60 Gas

C5H12 Pentane 36.2 0.63 Liquid

C6H14 Hexane 68.8 0.66 Liquid

C7H16 Heptane 98.5 0.68 Liquid

C8H18 Octane 125.8 0.70 Liquid

C9H20 Nonane 150.9 0.72 Liquid

C10H22 Decane 174.4 0.73 Liquid

SATURATED HYDROCARBONS OR ALKANES

Those hydrocarbons which contains as many hydrogen atoms as possible are said to be
saturated. On the other hand saturated hydrocarbons contains only a carbon-carbon single
bond. The saturated hydrocarbons are also known as alkanes. The simplest allkane is
methane (CH4).

The structural formula for methane is:

The second member of alkane is ethane with formula C2H6

The third member of alkane is propane which has a formula C3H8

The fourth member of alkane is Butane which has a formula of C4H10

If we see the above consecutive members of alkane each consecutive members differ
only by a constant – CH2 group.

A series of compounds in which each member differs from the next members by a
constant amount is called a Homologous series & the members of the series are called
homologous.
The general formula for alkanes is CnH2n+2 where n=1, 2, 3…
Alkyl radicals are obtained by removing one hydrogen atom from the corresponding
alkanes. Their general formula is CnH2n+1, where n = 1, 2, 3 . . . . The names of alkyl
radicals are derived from the names of the corresponding parent alkanes by changing the
suffix -ane to -yl.
Table 8.3 indicates formulas, names and structures of some alkyl groups.
NOMENCLATURE OF ALKANES:-
Common names such as pentane, isopentane and neopentane are sufficient to differentiate
between the three isomers with the formula C5H12. However as the size of the
hydrocarbon chain increases an international naming is required. The international union
of pure and applied chemistry (IUPAC) has developed a systematic approach to naming
of alkanes based on the following rules.
Rule – 1: Find the longest continuous chain atoms in the skeletal structure & name the
compound as the derivative of the alkane with this no of carbon atoms.
The names of the straight chain hydrocarbons is derived from Latin or Greek origin some
of the 1st 10 straight chain alkane are given in the above table.
Rule- 2: Nam the substituents on the chain, substituents derived from alkanes are named
by replacing the-ane ending with –yl.
Example: The substituent derived from methane is methyl (CH3-)

Rule- 3: Number the longest chain starting at the end nearest to the first substitute and
specify the no of carbon atoms on which the substituents are located use the lowest
possible numbers.

This compound is named as 2-methylpentane, not 4-methyl pentane.

Rule- 4: Use the prefixes di-,tri-, and tetra- to describe substituents that are found two,
three of four times respectively on the same chain.

This compound is named as 2,2,4-trimethyl pentane.

Rule- 5: Arrange the names of the substituents in alphabetical order.
 Is named as 4-ethyl- 2,2,7-trimethyl octane
Activity: Name the following compound

PHYSICAL AND CHEMICAL PROPERTIES OF ALKANES

What are four physical properties of alkanes?
 Insoluble in water.
 Less dense than water.
 Boiling point is directly related to length of chain.
 At standard conditions:
 Methane to butane are gases
 Pentane to heptadecane C17H36 are liquids
 Octadecane and up are solids (as indicated in table 2.1 above)

What are the chemical properties of alkanes?

Alkanes are also called paraffins which comes from Latin meaning lacking affinity.
Therefore, alkanes have low level of reactivity secondary to stable C-H and H-H bonds.
ISOMERS
Isomers are compounds which have the same molecular formula but different
arrangement of atoms and because of this difference in their structures isomers have
different physical properties such as (boiling point, refractive indexes, vapor pressures,
heat of combustions, solubility etc.)
Saturated hydro carbons shows structural isomerism. The maximum number of structural
isomers of alkane is determined by the general formula of 2n-4 +1 where n is the number
of carbon atoms.
Example: How many isomers does pentane have?
Solution: n=5
no of isomers = 2n-4+1
 = 25-4+1
=3

Preparations of alkanes:
Alkanes are prepared in the laboratory by decarboxylation of carboxylate salts.
i.e. RCOONa + NaOH heat R-H + Na2CO3
Carboxylate salt alkane
ALKENES S TRUCTURE
The chemical bonding in alkenes can be illustrated by reference to the simplest alkene,
ethene. This compound has the following structural form.

The double bond represents a four electron bond (i.e. two shared pairs of electrons).
However, the two bond between the carbon atoms have significantly different chemical
properties, and are formed in different ways. The first bond between the carbon atoms in
ethene is a s bonds (sigma bonds) and is similar to the carbon to carbon bond found in the
alkane series. However, the second bond between the carbon atoms in ethene is a p-bond
(pi-bond), which is much more reactive than the sigma bond and behaves differently in a
variety of experimental conditions.

The ethene molecule is planar (i.e. all atoms lie in the same plane) and the bond angle
between all the bonds (i.e. carbon to carbon and carbon to hydrogen) is 120 degrees. This
observed structure for ethene can be explained in terms of sp2 hybridisation of the
orbitals on the carbon atom.

Thus, ethene is a flat molecule, the distance between the carbon atoms being less than
that in ethene.

ALKENE NOMENCLATURE
Alkenes represent one of the most common functional groups in organic chemistry. An
alkene contains only carbon and hydrogen (a hydrocarbon) and contains at least one
double bond (termed an unsaturated hydrocarbon). Alkenes have the general formula
CnH2n,(n>2) thus, an alkene with 10 carbons (n = 10) will have 2(10) = 20 hydrogen, or
themolecular formula C10H20; each double bond therefore contributes one degree of
unsaturation.

The root, or parent name for an unbranched alkene is taken directly from the number of
carbons in the chain according to a scheme of nomenclature established by the
International Union of Pure and Applied Chemistry (IUPAC), as described previously for
alkanes.

To name alkenes:
1. Find the longest chain containing the alkene(double bond)
The IUPAC name for an alkene is constructed of two parts: 1) a prefix (meth... eth...
prop..., etc.) which indicates the number of carbons in the main, or parent, chain of the
molecule, and 2) the suffix ...ene to indicate that the molecule is an alkane.
For branched-chain alkanes, the name of the parent hydrocarbon is taken from the
longest continuous chain of carbon atoms containing the double bond.

2. Number the chain, giving the double bond the lowest possible number.
Numbering of the carbons in the parent chain is always done in the direction that gives
the lowest number to the double bond, or, the lowest number at the first point of
difference. If there are different substituents at equivalent positions on the chain, the
constituent of lower alphabetical order is given the lowest number.
If the same constituent occurs more than once in a molecule, the number of each carbon
of the parent chain where the constituent occurs is given and a multiplier is used to
indicate the total number of identical substituents; i.e., dimethyl... trimethyl...
tetraethyl..., etc.
In constructing the name, substituents are arranged in alphabetical order, without regard
for multipliers.

Example: Give IUPAC name for the following alkene

Parent name – Hexane

Position of double bond = 1- ene
Substitutes = 4, 4 – dimethyl
Therefore, the IUPAC name of the compound is: 4, 4- dimethyl-1-hexene.
Example: Give IUPAC name for the following Alkene

-1, 3- diene
- 2- methyl
Hence, IUPAC name is 2-methyl-1, 3-butadiene
ALKENES PHYSICAL PROPERTIES
The first three alkenes are gases, the intermediate alkenes are liquids and higher members
of the olefin series are wax like solids at room temperature. The alkenes are insoluble in
water, but are soluble in organic solvents. The liquids and solids have a density less than
water.

Compound Formula MP degC BP degC Density (g/ml)

Ethylene C2H4 -170 -102 0.6128

 Propene C3H6 -185 -47 0.6142

1-Butene C4H8 -130 -6.5 0.6356

ISOMERISM IN ALKENES
Unlike alkanes alkenes have geometric isomerism in addition to its position isomerism&
constitutional isomerism.

 Constitutional (structural) isomerism is due to substituents.

 Position isomerism is due to the position of the double bond.

Now let us look geometric isomerism.

GEOMETRIC ISOMERISM
Geometric isomerism (also known as cis-trans isomerism or E-Z isomerism) is a form of
stereoisomerism.

The E-Z system is better for naming more complicated structures but is more difficult to
understand than cis-trans. The cis-trans system of naming is still widely used - especially
for the sort of simple molecules you will meet at this level.

Where the atoms making up the various isomers are joined up in a different order, this is
known as structural isomerism. Structural isomerism is not a form of stereoisomerism.

How geometric isomers arise?

These isomers occur where you have restricted rotation somewhere in a molecule. At an
introductory level in organic chemistry, examples usually just involve the carbon-carbon
double bond - and that's what this page will concentrate on.

Think about what happens in molecules where there is unrestricted rotation about carbon
bonds - in other words where the carbon-carbon bonds are all single. The next diagram
shows two possible configurations of 1,2-dichloroethane.
ALKYNES
 Have a general form of R–C≡C–R. Where R- is alkyl group
 Have the general formula of CnH2n-2 where n= 2,3,4,…..
 They are isomeric with cyclo alkenes
Each triple bond reduces the number of hydrogen atoms by 4. For example, consider
compounds having the formula C5H8. The formula of the five-carbon alkane pentane is
C5H12 so the difference in hydrogen content is 4. The formula of 5 carbon alkyne is C5H8.
Their structures are shown below.
H2 H2
C C
H3C C CH3
H2
pentane
H2C CH3

HC C CH2
pentyne
NOMENCLATURE OF A LKYNES
THE IUPAC NAMING OF ALKYNE IS SIMILAR TO THAT OF ALKENE EXCEPT THAT THE

ENDING SUFFIX –‘YNE’ REPLACES THE SUFFIX-‘ENE’ OF THE ALKENE.

RULES:
1. Determine the longest continuous chain of carbons that have the triple bond.
2. Decide whether to number beginning on the right end or left end of the chain. If it
makes no difference to the triple bond then shift attention to the branched groups.
3. Name the branched groups in alphabetical order attaching (hyphenating) the
carbon number it is attached to along the continuous chain. of carbons to the front
of the branch name. If more than one of the same kind of branched group is
attached to the chain, identify the number carbon each group is attached to as a
series of numbers separated by commas between each number then a hyphen and
finally use a greek prefix attached to the branch name.
4. Attach a numerical prefix indicating the lowest carbon number the triple bond is
between onto the normal alkane name
5. Drop the "ane" ending and add the "yne" ending.

Let's try an example. Determine the IUPAC name of the following structure:

1. Determine the longest continuous chain of carbons that have the triple bond. The
chain does not necessarily have to be straight. That would be seven carbons long.
Actually there are two alternate continuous chains of carbpn that have seven
carbons, but teh rules say that we choose the one that results in the greatest
number of substituents.
2. Number the carbons in the chain so that the triple bond would be between the
carbons with the lowest designated number. This means that you have to decide
whether to number beginning on the right end or left end of the chain. If it makes
no difference to the triple bond then shift attention to the branched groups. We
would number from the left end and go toward the right end of the continuous
chain.
3. Identify the various branching groups attached to this continuous chain of carbons
by name There is a methyl group on carbon #4, an ethyl group on carbon #5, and
a Chloro group on carbon #6
 4. Name the branched groups in alphabetical order attaching (hyphenating) the
carbon number it is attached to along the continuous chain of carbons to the front
of the branch name. If more than one of the same kind of branched group is
attached to the chain, identify the number carbon each group is attached to as a
series of numbers separated by commas between each number then a hyphen and
finally use a greek prefix attached to the branch name. 6-Chloro-5-ethyl-4-methyl
5. Attach a numerical prefix indicating the lowest carbon number the triple bond is
between onto the normal alkane name 6-Chloro-5-ethyl-4-methyl-2-heptane
6. Drop the "ane" ending and add the "yne" ending associated with the Alkene
family 6-Chloro-5-ethyl-4-methyl-2-heptyne
Example:

a)

b
ALKYNES PHYSICAL PROPERTIES

Alkynes are compounds which have low polarity, and have physical properties that are
essentially the same as those of the alkanes and alkenes.
 They are insoluble in water.
 They are quite soluble in the usual organic solvents of low polarity (e.g. ligroin, ether,
benzene, carbon tetrachloride, etc.).
 They are less dense than water.
 Their boiling points show the usual increase with increasing carbon number.
 They are very nearly the same as the boiling points of alkanes or alkenes with the
same carbon skeletons.

Table 8.4 below show the physical properties of Alkynes

Name Formula MP degC BPdegC
Density(20C)
Acetylene HCCH -82 -75
Propyne HCCCH3 -101.5 -23
1-Butyne HCCCH2CH3 -122 91
1-Pentyne HCC(CH2)2CH3 -98 40 0.695
1-Hexyne HCC(CH2)3CH3 -124 72 0.719
1-Heptyne HCC(CH2)4CH3 -80 100 0.733
1-Octyne HCC(CH2)5CH3 -70 126 0.747
1-Nonyne HCC(CH2)6CH3 -65 151 0.763
1-Decyne HCC(CH2)7CH3 -35 182 0.770
2-Butyne CH3CCCH3 -24 27 0.694
2-Pentyne CH3CCCH2CH3 -101 55 0.714

ALCOHOLS

Alcohols are one of the oxygen derivative hydrocarbons. They have a general form of R-OH
; where R= alkyl or aryl group & -OH= the functional group.

NOMENCLATURE OF ALCOHOLS
The systematic (IUPAC) nomenclature for alcohols adds the ending -ol to the name of the
parent alkane and uses a number to identify the carbon that carries the OH group. The
systematic name for isopropyl alcohol, for example, is 2-propanol.
 Formula IUPAC name Common name

CH3-OH Methanol methyl alcohol

CH3CH2-OH Ethanol ethyl alcohol
CH3CH2CH2OH Propanol propyl- alcohol
CH3CHOHCH3 2-Propanol isopropyl alcohol

1) Prefix: Take the alkane (or alkene) name that corresponds to the correct number of
carbon atoms and remove the final e from the name.
2) Position number: Count the number of carbon atoms from the nearest chain end to
where the hydroxyl group, -OH, is attached. Put hyphens around this number and
place it after the prefix.
3) Suffix: Finally the suffix ol is added to denote an alcohol.

Examples:

- or
Ethanol

or

2-propanol

or

2-Methylpropan-2-ol
 Physical property of Alcohols
Lower alcohols are colorless and toxic liquids. They have a characteristic odor. The boiling
point of alcohol increases with increasing number of carbon atoms.

Between isomers as branching increases, the boiling point decreases. Boiling point of alcohol
is much higher than the corresponding alkenes because of the strong hydrogen bond in OH.
Lower alcohols (from C1 to C3) are completely soluble in water.

Classification of Alcohols
Alcohols are classified as either primary (10), secondary (20), or tertiary (30) on the basis of
their structures.

a. Primary alcohols: there is only one hydrocarbon group (R-group) attached to the carbon
containing “OH” functional group.

b. Secondary alcohols: alcohols containing two “R”–groups bonded to the carbon

containing OH- group. For example:

c. Tertiary alcohols: Contain three ‘R’ groups bonded to the carbon containing OH- group.
For example:

All the above alcohols have only one OH, and hence they are called monohydric alcohols.
And their general formula is CnH2n+1 OH. Alcohols containing more than one hydroxyl group
are also known as polyhydric alcohol. These are:

I. Dihydric alcohols: alcohols with two hydroxyl group, and their name ends in prefix
‘diols’. They are commonly known as glycol. For example:
II. Trihydric alcohols: alcohols that contain three OH groups and their name ends in prefix
‘triol’. For example:

PREPARATION OF ETHANOL
a. Fermentation: This basic and ages old procedure involves the action of zymase
(yeast, an enzyme) on an aqueous solution of sucrose (a simple sugar). The by-
products of the consumption of the sugar by the yeast are carbon dioxide and ethanol.
The carbon dioxide simply bubbles out of the solution, leaving behind a weak aqueous
solution of ethanol.

C6H12O6(aq) 2CH3CH2OH(aq) + 2CO2

This weak solution of ethanol in water can be made more concentrated by the process of
(fractional) distillation,

b. Hydration of ethene:

The fermentation of sucrose doesn't produce enough ethanol to be commercially viable, so a

different process is needed to produce the volumes of ethanol needed for industry.
The process of hydration is where the elements of water (2 H atoms and 1 O atom) are added
to a molecule. In this particular case water can be added to ethene to produce ethanol, with no
other by-products. It is not a straight forward process and requires a high temperature and
high pressure to generate the steam. The reaction also uses a catalyst of sulphuric acid to
speed up the whole process.

CH2=CH2 + H2O CH3CH2OH

REACTIONS OF ETHANOL
a. Combustion: Ethanol (along with all organic compounds) burns in excess oxygen to
give carbon dioxide and water as the only products.
 CH3CH2OH + 3O2(g) 2CO2(g) + 3H2O(g)
N.B.: Less oxygen is needed to ensure complete combustion of an alcohol, than a
corresponding alkane, because it contains an oxygen atom already.

b. With sodium metal: Just as with water, sodium metal will react with ethanol, though
not nearly as violently. The products are hydrogen gas (exactly the same as with water)
and sodium ethoxide (different to the reaction with water).
2CH3CH2OH + 2Na(s) 2CH3CH2ONa + H2(g)

The hydrogen gas produced does not ignite, as the enthalpy change of reaction is not high
enough. This is because the oxygen atom in water has two hydrogen atoms attached to it and
only one in ethanol.

c. Oxidation: Refluxing ethanol with oxidizing agents such as acidified manganate

(VII)(aq) and acidified dichromate(VI)(aq) ions changes it into ethanoic acid.

CH3CH2OH + [O] CH3COOH

Purple manganite (VII) ions (MnO4-) turn colorless (Mn2+) and orange dichromate (VI) ions
(Cr2O72-) turn green (Cr3+).

d. Halogenation: Whilst it is easy to hydrolyse haloalkanes into alcohols, the reverse

reaction requires quite extreme conditions.

A concentrated mixture of halide ions and acid are used (for example, sodium chloride in
concentrated sulphuric acid) along with heat, to ensure that the -OH group leaves the alcohol
(as water) and the carbon atom will accept a Cl- ion.

CH3CH2OH + Cl- + H+ CH3CH2Cl + H2O

There is an old test involving the halogenation of an alcohol to determine whether an alcohol
is primary, secondary or tertiary. A solution of zinc chloride in hydrochloric acid is added to
the alcohol. A primary alcohol will be the slowest to show a reaction and the tertiary will be
the fastest.

e. Dehydration: This process is the removal of the elements of water (2 H atoms and 1 O
atom) from ethanol leaving ethene. It is accomplished by refluxing ethanol with a
 catalyst of concentrated sulphuric or phosphoric acid, or by passing ethanol vapour over
heated aluminium oxide.

CH3CH2OH CH2=CH2 + H2O

f. Esterification: An ester is formed by reacting an alcohol with a carboxylic acid. It is a

similar type of reaction to the neutralization of an acid with a base. A catalyst of
concentrated sulphuric acid removes the water produced and helps to push the
equilibrium towards the products side of the equation.
CH3CH2OH + CH3COOH CH3COOCH2CH3 + H2O

Carboxylic acids
Carboxylic acid, any of a class of organic compounds in which a carbon (C) atom is bonded to an oxygen (O)
atom by a double bond and to a hydroxyl group (―OH) by a single bond. The carboxyl (COOH) group is so-
named because of the carbonyl group (C=O) and hydroxyl group. The general formula of a carboxylic acid is
R−COOH or R−CO2H, with R referring to the alkyl, alkenyl, aryl, or other group. Carboxylic acids occur
widely.

Structure of Carboxylic Acids

Carboxylic acids are organic compounds that contain at least one carboxyl group in their
structure. A carboxyl group is a functional group consisting of a carboxyl and a hydroxyl
which is usually written as —COOH or —CO2H
Nomenclature of Carboxylic Acids
Common names carboxlic acids
a. Straight chain monocarboxylic acids
A large number of carboxylic acids have widely used common names. Those with an even
number of carbon atoms ranging from 4 to 22 may be obtained by hydrolysis of animal and
vegetable fats and oils. They are referred to as fatty acids, and they have common names
derived from various sources. Formic acid derives its name from the Latin word for ants,
because it is one of the toxic ingredients of the secretion injected by the stinging ant.
Butanoic acid (butyric acid) derives its name from butter, in which it is found when the butter
becomes rancid. Caproic, caprylic, and capric acids are involved in the odor of a goat, and
their names derive from the Latin word, caper, for goat.
Table 8.5 lists common names of some of the most important monocarboxylic acids.
 b. Branched chain and substituted carboxylic acids
In common naming system, the branched chain and substituted acids are named as derivatives
of straight chain carboxylic acids. In this case, the position of the side chain or substituents is
indicated by Greek letters, α, β, γ, δ... for designating the 1st, 2nd, 3rd,… position of carbon
atoms as shown below:

Example: Write the common names for:

O OH

H3C CH2

CH3 CH3
a.
Solution:
It is monocarboxylic acid that consists of four carbon atoms. So, it is common name is
butyric acid. In addition to this, two methyl groups are attached to α -and β-positon carbon
atoms in the structure. Now, the complete common name is α, β -dimethylbutyric acid.
OH

O C Br

b. H2C CH2

Solution:
Similarly, the given acid contains three carbon atoms. Hence, the common name is propionic
acid. Next, the position of bromo group, which is attached to α-carbon atom in the structure.
This gives the complete common name as β -bromopropionic acid.
The IUPAC name of a carboxylic acid is derived from that of the longest carbon chain that contains the carboxyl
group by dropping the final -e from the name of the parent alkane and adding the suffix -oic followed by the
word “acid.” The chain is numbered beginning with the carbon of the carboxyl group.

In systematic nomenclature, the position of a substituent is designated by a number. The

carbonyl carbon of a carboxylic acid is always the C-1 carbon. In common nomenclature, the
position of a substituent is designated by a lowercase Greek letter, and the carbonyl carbon is
not given a designation. The carbon adjacent to the carbonyl carbon is the α-carbon, the
carbon adjacent to the α-carbon is the β-carbon and so on.

In systematic nomenclature, a carboxylic acid is named by replacing the terminal “e” of the
alkane name with “oic acid.” For example, the one-carbon alkane is methane, so the one-
carbon carboxylic acid is methanoic acid.