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SOME BASIC CONCEPTS OF CHEMISTRY
SOME BASIC CONCEPTS OF CHEMISTRY
1. CHEMISTRY 2. MATTER
Chemistry is defined as the study of the composition, Matter is defined as any thing that occupies space
properties and interaction of matter. Chemistry is often possesses mass and the presence of which can be felt by
any one or more of our five senses.
called the central science because of its role in connecting
the physical sciences, which include chemistry, with the Matter can exist in 3 physical states viz. solid, liquid, gas.
life sciences and applied sciences such as medicine and Solid - a substance is said to be solid if it possesses a
engineering. definite volume and a definite shape, e.g., sugar, iron, gold,
wood etc.
Various branches of chemistry are
Liquid- A substance is said to be liquid, if it possesses a
1.1 Physical chemistry definite volume but no definite shape. They take up the
shape of the vessel in which they are put, e.g., water, milk,
The branch of chemistry concerned with the way in which oil, mercury, alcohol etc.
the physical properties of substances depend on and Gas- a substance is said to be gaseous if it neither possesses
influence their chemical structure, properties, and reactions. definite volume nor a definite shape. This is because they
fill up the whole vessel in which they are put, e.g., hydrogen,
1.2 Inorganic chemistry
oxygen etc.
The branch of chemistry which deals with the structure, The three states are interconvertible by changing the
composition and behavior of inorganic compounds. All the conditions of temperature and pressure as follows
substances other than the carbon-hydrogen compounds
are classified under the group of inorganic substances.
1.3 Organic chemistry
The discipline which deals with the study of the structure,
composition and the chemical properties of organic
compounds is known as organic chemistry.
1.4 Biochemistry
The discipline which deals with the structure and behavior
of the components of cells and the chemical processes in
living beings is known as biochemistry.
1.5 Analytical chemistry 3. CLASSIFICATION OF MATTER AT
MACROSCOPIC LEVELL
The branch of chemistry dealing with separation,
identification and quantitative determination of the At the macroscopic or bulk level, matter can be classified as
compositions of different substances. (a) mixtures (b) pure substances.
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SOME BASIC CONCEPTS OF CHEMISTRY
These can be further sub-divided as shown below
Pure substances can be further classified into elements
and compounds.
Element- An element is defined as a pure substance that
contains only one kind of particles. Depending upon the
physical and chemical properties, the elements are further
subdivided into three classes, namely (1) Metals (2) Non-
metals and (3) Metalloids.
Compound- A compound is a pure substance containing
two or more than two elements combined together in a fixed
proportion by mass. Further, the properties of a compound
are completely different from those of its constituent
elements. Moreover, the constituents of a compound
cannot be separated into simpler substances by physical
(a) Mixtures : A mixture contains two or more substances
methods. They can be separated by chemical methods.
present in it (in any ratio) which are called its
components. A mixture may be homogeneous or
4. PROPERTIES OF MATTER
heterogeneous.
Homogeneous mixture- in homogeneous mixture the Every substance has unique or characteristic properties.
components completely mix with each other and its These properties can be classified into two categories –
composition is uniform throughout i.e it consist of only physical properties and chemical properties.
one phase. Sugar solution and air are thus, the examples of
homogeneous mixtures. 4.1 Physical Properties
Heterogeneous mixtures- In heterogeneous mixture the
Physical properties are those properties which can be
composition is not uniform throughout and sometimes the
different phases can be observed. For example, grains and measured or observed without changing the identity or the
pulses along with some dirt (often stone) pieces, are composition of the substance. Some examples of physical
heterogeneous mixtures. properties are color, odor, melting point, boiling point,
density etc.
4.2 Chemical properties
Any distinct portion of matter that is uniform throughout Chemical properties are those in which a chemical change
in composition and properties is called a Phase. in the substance occurs. The examples of chemical properties
are characteristic reactions of different substances; these
(b) Pure substances :- A material containing only one
include acidity or basicity, combustibility etc.
substance is called a pure substance.
5. MEASUREMENT
5.1 Physical quantities
In chemistry, a substance is a form of matter that has
constant chemical composition and characteristic All such quantities which we come across during our
properties. It cannot be separated into components by scientific studies are called Physical quantities. Evidently,
physical separation methods, i.e. without breaking the measurement of any physical quantity consists of two
chemical bonds. They can be solids, liquids or gases. parts
SOME BASIC CONCEPTS OF CHEMISTRY
(1) The number, and (2) The unit International d’Unités – abbreviated as SI) was established
by the 11th General Conference on Weights and Measures
A unit is defined as the standard of reference chosen to (CGPM from Conference Generale des Poids at Measures).
measure any physical quantity. The CGPM is an inter governmental treaty organization
created by a diplomatic treaty known as Meter Convention
5.2 S.I. UNITS which was signed in Paris in 1875.
The International System of Units (in French Le Systeme The SI system has seven base units and they are listed in
table given below.
These units pertain to the seven fundamental scientific quantities. The other physical quantities such as speed, volume,
density etc. can be derived from these quantities. The definitions of the SI base units are given below :
Definitions of SI Base Units
Unit of length metre The metre is the length of the path travelled by light in
vacuum during a time interval of 1/299 792 458 of a second.
Unit of mass Kilogram The kilogram is the unit of mass; it is equal to the mass of
the internationl prototype of the kilogram.
Unit of time second The second is the duration of 9 192 631 770 periods of the
radiation corresponding to the transition between the two
hyperfine levels of the ground state of the caesium-133
atom.
Unit of electric current ampere The ampere is that constant current which, if maintained in
two straight parallel conductors of infinite length, of
negligible circular cross-section, and placed 1 metre apart
in vacuum, would produce between these conductors a
force equal to 2 × 10–7 newton per metre of length.
Unit of thermodynanic kelvin The kelvin, unit of thermodynamic temperature, is the
temperature fraction 1/273. 16 of the thermodynamic temperature of the
triple point of water.
Unit of amount of substance mole 1. The mole is the amount of substance of a system
which contains as many elementary entities as there
are atoms in 0.012 kilogram of carbon-12; its symbol
is “mol.”.
2. When the mole is used, the elementary entities must
be specified and may be atoms, molecules, ions,
electrons, other particles, or specified groups of such
particles.
Unit of luminous intensity candela The candela is the luminous intensity, in a given direction,
of a source that emits monochromatic radiation of
frequency 540 × 1012 hertz and that has a radiant intensity
in that direction of 1/683 watt per steradian.
SOME BASIC CONCEPTS OF CHEMISTRY
7. LAW OF CHEMICAL COMBINATION
The mass standard is the kilogram since 1889. It has been
defined as the mass of platinum-iridium (Pt-Ir) cylinder 7.1 Law of conservation of mass
that is stored in an airtight jar at International Bureau of
“In a chemical reaction the mass of reactants consumed
Weights and Measures in Sevres, France. Pt-Ir was chosen
and mass of the products formed is same, that is mass is
for this standard because it is highly resistant to chemical
conserved.” This is a direct consequence of law of
attack and its mass will not change for an extremely long
conservation of atoms. This law was put forth by Antoine
time.
Lavoisier in 1789.
6. SOME IMPORTANT DEFINITION 7.2 Law of Constant / Definite Proportions
6.1 Mass and Weight The ratio in which two or more elements combine to form a
compound remains fixed and is independent of the source
Mass of a substance is the amount of matter present in it of the compound. This law was given by, a French chemist,
while weight is the force exerted by gravity on an object. Joseph Proust.
The mass of a substance is constant whereas its weight
7.3 Law of Multiple Proportions
may vary from one place to another due to change in gravity.
The SI unit of mass is the kilogram (kg). The SI derived When two elements combine to form two or more compounds
unit (unit derived from SI base units) of weight is newton. then the ratio of masses of one element that combines with
a fixed mass of the other element in the two compounds is a
6.2 Volume
simple whole number ratio. This law was proposed by Dalton
Volume is the quantity of three-dimensional space enclosed in 1803.
by some closed boundary, for example, the space that a
7.4 Law of Reciprocal Proportions
substance (solid, liquid, gas, or plasma) or shape occupies
or contains. Volume is often quantified numerically using
the SI derived unit, the cubic meter. When three elements combine with each other in
combination of two and form three compounds then the
6.3 Density ratio of masses of two elements combining with fixed mass
of the third and the ratio in which they combine with each
The mass density or density of a material is defined as its other bear a simple whole number ratio to each other. This
mass per unit volume. The symbol most often used for Law was given by Richter in 1792.
density is U (the lower case Greek letter rho). SI unit of
density is kg m–3. 7.5 Gay Lussac’s Law of Gaseous Volumes
6.4 Temperature This law was given by Gay Lussac in 1808. He observed
that when gases combine or are produced in a chemical
Temperature is a physical property of matter that
reaction they do so in a simple ratio by volume provided all
quantitatively expresses the common notions of hot and
gases are at same temperature and pressure.
cold. There are three common scales to measure temperature
— °C (degree celsius), °F (degree fahrenheit) and K (kelvin).
7.6 Avogadro Law
The temperature on two scales is related to each other by
the following relationship: In 1811, Avogadro proposed that equal volumes of gases at
°F = 9/5 (°C) + 32 the same temperature and pressure should contain equal
number of molecules.
K = °C + 273.15
SOME BASIC CONCEPTS OF CHEMISTRY
8. DALTON’S ATOMIC THEORY 9.3 Method 2
In 1808, Dalton published ‘A New System of Chemical Mass of 6.022 × 10 23 atoms of that element taken in grams.
Philosophy’ in which he proposed the following: This is also known as molar atomic mass.
1. Matter consists of indivisible atoms.
2. All the atoms of a given element have identical properties
including identical mass. Atoms of different elements differ
in mass. Mass of 1 atom in amu and mass of
6.022 × 1023 atoms in grams are numerically equal.
3. Compounds are formed when atoms of different elements
combine in a fixed ratio. When atomic mass is taken in grams it is also called
the molar atomic mass.
4. Chemical reactions involve reorganization of atoms. These
are neither created nor destroyed in a chemical reaction. 6.022 × 1023 is also called 1 mole of atoms and this
number is also called the Avogadro’s Number.
9. ATOM
Mole is just a number. As 1 dozen = 12;
Atom is the smallest part of an element that can participate in a
chemical reaction. {Note : This definition holds true only for 1 million = 10 6; 1 mole = 6.022 × 1023.
non-radioactive reactions}
9.1 Mass of an Atom
10. MOLECULES
A group of similar or dissimilar atoms which exist together in
There are two ways to denote the mass of atoms. nature is known as a molecule. e.g. H2, NH3.
9.2 Method 1 The mass of molecules is measured by adding the masses of
the atoms which constitute the molecule. Thus, the mass of
Atomic mass can be defined as a mass of a single atom which a molecule can also be represented by the two methods used
is measured in atomic mass unit (amu) or unified mass (u) for measuring the mass of an atom viz. amu and g/mol.
where
1 a.m.u. = 1/12th of the mass of one C12 atom
SOME BASIC CONCEPTS OF CHEMISTRY
is called the Excess Reagent. e.g. if we burn carbon in air
11. CHEMICAL REACTIONS (which has an infinite supply of oxygen) then the amount
of CO2 being produced will be governed by the amount of
A chemical reaction is only rearrangement of atoms. Atoms
carbon taken. In this case, Carbon is the LR and O2 is the
from different molecules (may be even same molecule) rearrange ER.
themselves to form new molecules.
13. PERCENT YIELD
Points to remember :
As discussed earlier, due to practical reasons the amount of
Always balance chemical equations before doing any product formed by a chemical reaction is less than the amount
calculations
predicted by theoretical calculations. The ratio of the amount
The number of molecules in a reaction need not to be of product formed to the amount predicted when multiplied
conserved e.g. by 100 gives the percentage yield.
N2 + 3 H2 Æ 2 NH3. The number of molecules is not
conserved Actual Yield
Percentage Yield = Theoretical Yield × 100
If we talk about only rearrangement of atoms in a
balanced chemical reaction then it is evident that the
mass of the atoms in the reactants side is equal to the
14. REACTIONS IN AQUEOUS MEDIA
sum of the masses of the atoms on the products side. Two solids cannot react with each other in solid phase
This is the Law of Conservation of Atoms and Law of and hence need to be dissolved in a liquid. When a solute
Conservation of Mass. is dissolved in a solvent, they co-exist in a single phase
called the solution. Various parameters are used to measure
the strength of a solution.
12. STOICHIOMETRY
The strength of a solution denotes the amount of solute
The study of chemical reactions and calculations related which is contained in the solution. The parameters used to
to it is called Stoichiometry. The coefficients used to balance denote the strength of a solution are:
the reaction are called Stoichiometric Coefficients.
Mole fraction X : moles of a component / Total moles
of solution.
Points to remember :
Mass% : Mass of solute (in g) present in 100g of
The stoichiometric coefficients give the ratio of
solution.
molecules or moles that react and not the ratio of
masses. Mass/Vol : Mass of solute (in g) present in 100mL of
solution
Stoichiometric ratios can be used to predict the moles of
product formed only if all the reactants are present in the v/v : Volume of solute/volume of solution {only for
stoichiometric ratios. liq-liq solutions}
Practically the amount of products formed is always g/L : Wt. of solute (g) in 1L of solution
less than the amount predicted by theoretical
calculations mass of solute
ppm : mass of solution u 10
6
12.1 Limiting Reagent (LR) and Excess Reagent (ER)
moles of solute
If the reactants are not taken in the stoichiometric ratios Molarity (M) : volume of solution (L)
then the reactant which is less than the required amount
determines how much product will be formed and is known moles of solute
as the Limiting Reagent and the reactant present in excess Molality (m) : mass of solvent (kg)
SOME BASIC CONCEPTS OF CHEMISTRY
IMPORTANT RELATIONS e.g. molality remains unchanged with temperature. Formulae
involving volume are altered by temperature e.g. Molarity.
1. Relation between molality (m) Molarity (M), density (d) of
solution and molar mass of solute (MO) 17. INTRODUCTION TO EQUIVALENT CONCEPT
d : density in g/mL Equivalent concept is a way of understanding reactions
and processes in chemistry which are often made simple
MO : molar mass in g mol–1
by the use of Equivalent concept.
M u 1000 17.1 Equivalent Mass
Molality, m
1000d MM O
2. Relationship between molality (m) and mole fraction (XB) of “The mass of an acid which furnishes 1 mol H+ is called
the solute its Equivalent mass.”
“The mass of the base which furnishes 1 mol OH– is called
XB 1000 1 X A 1000
m u m u its Equivalent mass.”
1 XB MA XA MA
17.2 Valency Factor (Z)
Points to remember :
Valency factor is the number of H+ ions supplied by 1
Molarity is the most common unit of measuring
molecule or mole of an acid or the number of OH- ions
strength of solution.
supplied by 1 molecule or 1 mole of the base.
The product of Molarity and Volume of the solution
gives the number of moles of the solute, n = M × V Molecular Mass
Equivalent mass, E
All the formulae of strength have amount of solute. Z
(weight or moles) in the numerator.
17.3 Equivalents
All the formulae have amount of solution in the
denominator except for molality (m).
wt. of acid / base taken
No. of equivalents =
15. DILUTION LAW Eq. wt.
When a solution is diluted, more solvent is added, the moles
of solute remains unchanged. If the volume of a solution
having a Molarity of M1 is changed from V1 to V2 we can
write that: It should be always remembered that 1 equivalent of an
acid reacts with 1 equivalent of a base.
M1V1 = moles of solute in the solution = M2V2
16. EFFECT OF TEMPERATURE 18. MIXTURE OF ACIDS AND BASES
Volume of the solvent increases on increasing the Whenever we have a mixture of multiple acids and bases
temperature. But it shows no effect on the mass of solute in we can find whether the resultant solution would be acidic
the solution assuming the system to be closed i.e. no loss of or basic by using the equivalent concept. For a mixture of
mass. multiple acids and bases find out the equivalents of acids
and bases taken and find which one of them is in excess.
The formulae of strength of solutions which do not involve
volume of solution are unaffected by changes in temperature.
SOME BASIC CONCEPTS OF CHEMISTRY
22. EQUIVALENT VOLUME OF GASES
19. LAW OF CHEMICAL EQUIVALENCE
Equivalent volume of gas is the volume occupied by 1
According to this law, one equivalent of a reactant
equivalent of a gas at STP.
combines with one equivalent of the other reactant to
give one equivalent of each product . For, example in a Equivalent mass of gas = molecular mass /Z.
reaction aA + bB o cC + dD irrespective of the
Since 1 mole of gas occupies 22.4L at STP therefore 1
stoichiometric coefficients, 1 eq. of A reacts with 1 eq. of
equivalent of a gas will occupy 22.4/Z L at STP. e.g. Oxygen
B to give 1 eq. each of C and 1eq of D
occupies 5.6L, Chlorine and Hydrogen occupy 11.2L.
20. EQUIVALENT WEIGHTS OF SALTS 23. NORMALITY
To calculate the equivalent weights of compounds which
The normality of a solution is the number of equivalents of
are neither acids nor bases, we need to know the charge on
solute present in 1L of the solution.
the cation or the anion. The mass of the cation divided by
the charge on it is called the equivalent mass of the cation
equivalents of solute
and the mass of the anion divided by the charge on it is N
volume of solution (L)
called the equivalent mass of the anion. When we add the
equivalent masses of the anion and the cation, it gives us
the equivalent mass of the salt. For salts, Z in the total The number of equivalents of solute present in a solution is
amount of positive or negative charge furnished by 1 mol given by Normality × Volume (L).
of the salt. On dilution of the solution the number of equivalents of
the solute is conserved and thus, we can apply the
21. ORIGIN OF EQUIVALENT CONCEPT
formula : N1V1 = N2V2
Equivalent weight of an element was initially defined as
Caution :
weight of an element which combines with 1g of hydrogen.
Later the definition wad modified to : Equivalent weight of Please note that the above equation gives rise to a lot of
an element is that weight of the element which combines confusion and is a common mistake that students make.
with 8g of Oxygen. This is the equation of dilution where the number of
equivalents are conserved. Now, since one equivalent of
a reactant always reacts with 1 equivalent of another
reactant a similar equation is used in problems involving
titration of acids and bases. Please do not extend the same
logic to molarity.
Same element can have multiple equivalent weights
depending upon the charge on it. e.g. Fe2+ and Fe3+.
Relationship between Normality and Molarity
N = M × Z ; where ‘Z’ is the Valency factor
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