General Chemistry

Principles and Modern Applications

Petrucci • Harwood • Herring
8th Edition

Chapter 2: Atoms and the Atomic Theory

 Contents

2-1 Early Chemical Discoveries and Atomic Theory

2-2 Electrons and Other Discoveries in Atomic Physics
2-3 The Nuclear Atom
2-4 Chemical Elements
2-5 Atomic Mass
2-6 Introduction to the Periodic Table
2-7 The Mole Concept and Avagadro Constant
2-8 Using the Mole Concept in Calculations
 2-1 Early Discoveries and the Atomic
Theory

Lavoisier 1774 - Law of conservation of mass

Proust 1799- Law of constant composition

and Law of definite proportions

Dalton (1803-1808)- Atomic Theory

Law of conservation of mass
Lavoisier 1774

Heat

Tin Tin oxide

Air Air
Flask
Flask

total mass = total mass

 Conservation of Mass

Reaction
AgCl + KCrO4 KCl + AgCrO4 (red)
104.50 g 104.50 g
 The Mass Laws
Law of Conservation of Mass:
The total mass of substances does not change during a
chemical reaction.
“Mass is neither created nor destroyed”

reactant 1 + reactant 2 product

total mass = total mass

calcium oxide + carbon dioxide calcium carbonate

CaO + CO2 CaCO3

56.08g + 44.00g 100.08g

Law of Constant Composition
French Chemist Joseph Proust (1799)

Cu metal dissolve in HNO3

Cu2+ + K2CO3 → CuCO3 (green color)

Malachite mineral Roof of the houses

 Law of Constant Composition
Joseph Proust (1754–1826)

• Also known as the law of definite proportions.

• The elemental composition of a pure substance
never varies.
• The relative amounts of each element in a
compound doesn’t vary.

H N
NH3
ammonia
ammonia always has 3 H and 1 N.
 Law of Conservation of Mass

The total mass of substances present at the

end of a chemical process is the same as the
mass of substances present before the
process took place.

3H2 + N2 2NH3
ammonia
The atoms on the right all appear on the left
The Law of Definite Proportions
When elements combine to form compounds,
they do so in definite proportions by mass.

Sample A Composition Sample B Composition

10,000 g 27,000 g
1,119 g H % H = 11,19 3,021 g H % H = 11,19
8,881 g O % O = 88.81 23,979 g O % O = 88.81
Two samples have the same proportions of the two elements.

Na (s) + Cl2 (g) NaCl (s)

39.3 g 60.7 g 100.0 g
23.0 g 35.5 g 58.5 g
39.3 g 23.0g
60.7 g = 35.5 g
= 0.647
Law of Definite (or Constant) Composition:
No matter what its source, a particular chemical
compound is composed of the same elements in
the same parts (fractions) by mass.

CaCO3 1 atom of Ca 40.08 amu

1 atom of C 12.00 amu

3 atoms of O 3 x 16.00 amu
100.08 amu
40.08 amu
= 0.401 parts Ca
100.08 amu

12.00 amu
= 0.120 parts C
100.08 amu

48.00 amu
= 0.480 parts O
100.08 amu
 Law of multiple proportions
Dalton’s theory leads to a prediction- the law of multiple proportions.
~ If two elements form more than a single compound, the masses of one element
combined with a fixed mass of the second are in the ratio of small whole
numbers.
Same elements to combine in different ratios to give different substances.
E.g. Oxygen and carbon can combine either in a 1: 1.333 mass ratio to make a
substance or in a 1: 2.667 mass ratio to make a substance.

first 1 g carbon per 1.333 g oxygen C:O mass ratio = 1: 1.333

second 1 g carbon per 2.667 g oxygen C:O mass ratio = 1: 2.667

comparison C:O mass ratio in first sample = (1 g C)/(1.333g O) = 2

of C:O ratios C:O mass ratio in second sample (1 g C)/(2.667g O)
Compare two substances clearly the second substance contains exactly twice as
much oxygen as the first for a given number of carbon. If the first oxide has the
molecular formula CO then the second oxide will be CO2.
 Dalton’s Atomic Theory

 Each element is composed of small particles called atoms.

‚ Atoms are neither created nor destroyed in chemical reactions.

ƒ All atoms of a given element are identical (in mass and other

properties)

„ Compounds are formed when atoms of more than one element

combine in a simple numerical ratio.

Atomic Theory of Matter

The theory of atoms:

Original to the Greeks
Leucippus, Democritus and Lucretius
(Aristotle thought they were nuts)
John Dalton (1805-1808)
Rewieved the idea and made it
science by measuring the atomic
weights of 21 elements.

That’s the key thing because then

you can see how elements
combine.
 Dalton’s Postulates
1- Each element is composed of extremely
small particles called atoms.
 Dalton’s Postulates

3- All atoms of a given element are identical (in mass

and other properties). But the atoms of one element are
different from the atoms of all other elements.

O N
 Dalton’s Postulates
2- Atoms are neither created nor destroyed in chemical reactions.

Atoms of an element are not changed into atoms of a different element

by chemical reactions; atoms are neither created nor destroyed in
chemical reactions. (As far as Dalton knew, they couldn’t be changed at
all).

O N O N

Red O’s stay Os and blue N’s stay N’s.

 Dalton’s Postulates
4-Compounds are formed when atoms of more than one element
combine in a simple numerical ratio.

Compounds are formed when atoms of more than one element

combine; a given compound always has the same relative number
and kind of atoms.

H N
NH3
ammonia

Chemistry happens when the balls rearrange

Law of Multiple Proportions

Ratio of oxygen-to-carbon in CO2 is

exactly twice the ratio in CO.
The Law of Multiple Proportions - In different com-
pounds containing the same elements, the masses of
one element combined with a fixed mass of the other
element are in the ratio of small whole numbers.
H O O/H
H2O2 11.2 g 178 g 15.9
=2
H2O 11.2 g 88.8 g 7.93

C O O/C
CO2 42.9 g 114 g 2.66
=2
CO 42.9 g 57.1 g 1.33
Law of Multiple Proportions
 Consequences of Dalton’s theory
u Law of Definite Proportions: combinations of elements are
in ratios of small whole numbers.

¶ In forming carbon monoxide, 1.33 g

of oxygen combines with 1.0 g of
carbon.

¶ In the formation of hydrogen

peroxide 2.66 g of oxygen combines
with 1.0 g of hydrogen.

Slide 23 of 25 General Chemistry: Chapter 2 Prentice-Hall © 2002

Summary:

These laws that govern chemical reactions and com-

pound composition allow us to do stoichiometric
calculations. They were all well known by 1803 when
John Dalton put forth his atomic theory of matter.
This theory was able to “explain” why these laws that
govern chemical reactions were true.
 2-2 Electrons and Discoveries in Atomic Physics
Behavior of charges
a) Two ++ or two - - charges repel each other.
b) Objects lack any electric charge exert no
forces on each other.
c) Positive and negative charges attract
each other.

When charged particles travel in a magnetic field,

they are deflected.
Negatively charged particles are deflected in one
direction.
Positively charged particles are opposite direction.
 The Discovery of Electrons
Cathode Ray Tube (CRT) also known as TV tube.
Discovered cathode rays, a type of radiation
emitted by the negative terminal or the cathode.
Cathode rays travel in straight lines and are
Michael Faraday
independent of cathode materials (whether it is
1791-1867
iron, platinium and so on).

Cathode rays are

invisible and can be
detected only by
flourescent materials
(ZnS).
 Properties of cathode rays

Electron m/e = -5.6857 ´ 10-9 g coulomb-1

 Thomson’s Cathode Ray Experiment
1897 J.J. Thomson
Cathode rays are deflected by electric and magnetic
field, as negative charged particles.

Thomson found the ratio of mass (m) to electric charge (e), that is
m/e (m/e = -5.6857 ´ 10-9 g/coulomb).

Thomson concluded that cathode rays are negatively charged

fundamental particles of matter found in all atoms.

Cathode rays later became known as electrons, this term first

proposed by George Stoney in 1874.
Thomson’s experiment involved the use of cathode-ray tube. When a
sufficiently high voltage is applied across the electrode, an electric current
 Thomson’s Cathode Ray Experiment

• Used many different metals and gases

• Beam was always the same
• By the amount it bent he could find the ratio of
charge to mass
• Was the same with every material
• Same type of piece in every kind of atom
 Millikan’s Oil Drop Experiment
Millikan found that ionized oil drops can be balanced
against the pull of gravity by an electric field.
Determined a value for the electron’s charge:
e = –1.6022 × 10–19 C
Robert Millikan
1868-1953

Charged droplet can

move either up or down,
depending on the charge
on the plates.

Radiation ionizes
a droplet of oil.

Magnitude of charge on
the plates lets us calculate
the charge on the droplet.
 Charge on The Electron

Value for the electron’s charge:

e = –1.6022 × 10–19 C

Mass of an electron
From m/e and the charge, the mass of an electron was
determined to be
m = 9.1094 × 10–31 kg/electron
 Properties of The Electron
• Thomson determined the mass-to-charge ratio; Millikan
found the charge; we can now find the mass of an electron:
me = 9.109 × 10–31 kg/electron
• This is almost 2000 times less than the mass of a hydrogen
atom (1.79 × 10–27 kg)

• Some investigators thought that cathode rays (electrons) were

negatively charged ions.
• But the mass of an electron is shown to be much smaller than
even a hydrogen atom, so an electron cannot be an ion.
• Since electrons are the same regardless of the cathode
material, these tiny particles must be a negative part of all
matter.
 J. J. Thomson – Atomic Model
Thomson proposed an atom with a positively charged
sphere containing equally spaced electrons inside

Proposed for a hydrogen atom that there was one

electron at the exact center of the sphere
Proposed for a helium atom that two electrons existed
along a straight line through the center, with each
electron being halfway between the center and the outer
surface of the sphere
 Plum Pudding Model

• Thomson proposed an atom with a

positively charged sphere containing
equally spaced electrons inside.
• He applied this model to atoms with
up to 100 electrons.
 X-rays and Radioactivity

In 1895, discovered x-rays when working with cathode

rays tube. Found that x-rays pass through different
materials at different temperatures.
Wilhelm C.
Rontgen
1845-1923

Discovered that uranium was able to expose a

photographic plate on black paper.
In1896, discovered radioactivity.

Antoine H.
Becquerel
1852-1908
http://hi.fi.tripod.com/timeline/images/wilhelm_rontgen.jpg
 Radioactivity
Radioactivity is the spontaneous emission of radiation by an
atom.
• First observed by Henri Becquerel.
• Also studied by Marie and Pierre Curie.

Three types of radiation

• α Particles were discovered by Ernest Rutherford
• β Particles were discovered by Ernest Rutherford
• γ Rays were discovered by Paul Villard

Properties of the three radioactive emissions discovered

Original name Modern name Mass (amu) Charge
a-ray a-particle 4.00 +2
b-ray b-particle (electron) 5.49x10-4 -1
g-ray g-ray 0 0_______
 Radioactivity

• a particles, attracted to negative electrode, so they have a

positive charge, much more mass than negative stuff (turn out
to be He2+ nuclei).
• b particles, attracted to positive electrode, so they have a
negative charge, 1000s of times less massive (turn out to be
electrons coming from nucleus).
• g rays, no charge, no mass.
 The nuclear atom

Hans Geiger and

Ernest Rutherford Alpha Scattering
1090 Experiment:
Ernest
Rutherford
1871-1937
 Most of the alpha Alpha Scattering
particles passed
through the foil.
Experiment:
Rutherford’s observations

Alpha particles
were “shot” into
thin metal foil.

A few particles
were deflected
A very few slightly by the
“bounced foil.
back” to the
source!
Alpha Scattering If Thomson’s model was correct, most of
Experiment: the α particles should have been deflected
Rutherford’s a little, like bullets passing through a
conclusions cardboard target.

Most of the alpha

particles passed
through the foil =>
An atom must be
mostly empty space.

The nucleus is far

A very few alpha particles smaller than is
bounced back => suggested here.
The nucleus must be very
small and massive.
 Rutherford’s Scattering Experiment
1909 Ernest Rutherford
~ use a particle to study the inner structure of atoms. When he
directed a beam of a-particles at a thin gold foil, he found that
• The majority of a-particles penetrated the foil undeflected.
• Some a particles experienced slightly deflections.
• A few (about one in every 20,000) suffered rather serious
deflections as they penetrated the foil.
• A similar number did not pass through the foil at all, but
bounced back in the direction from which they had come.
 Rutherford’s Revised Atomic Theory (1911)

Result: Most of the positively charged particles went straight through

the gold foil.
Atomic Theory: Most of the matter of the atom is found in a very small
part of the atom. This is called the nucleus of the atom. It is very tiny
and extremely dense.

Result: Some of the positively charged particles were deflected or even

bounced back.
Atomic Theory: Like charges repel so the nucleus must have a positive
charge. If electrons have a negative charge they could not be in a
positively charged nucleus. Electrons must surround the nucleus at
a distance.

Result: The diameter of the nucleus is 100,000 times smaller than the
diameter of the entire gold atom.
Atomic Theory: Atoms are mostly empty space with a tiny, massive
nucleus at the center .
 The Nuclear Atom: Protons and Neutrons
1911 Rutherford explained his results by proposing a model of the atom known as the
nuclear atom and having these features.

1 Most of the mass and all of the positive charge of an atom are
centered in a very small region called the nucleus. The atom is mostly
empty space.

2 The magnitude of the positive charge is different for different atoms

and is approximately one-half the atomic weight of the element.

3 There are as many electrons outside the nucleus as there are units of
positive charge on the nucleus. The atom as a whole is electrically
neutral.

Rutherford’s nuclear atom suggested the existence of positively charged fundamental

particles of matter in the nuclei of atoms- called protons. He predicted the existence in
the nucleus of electrically neutral particles.

1932 Jame Chadwick

~ verified that there is another type of particles in atom called neutron.
 •Mo
Structure of the Atom
con
The atom is composed of two kinds of particles: elec
Nucleus
Central core
Positively charged
Contains most of the atom’s mass
Electrons
Very light
Negatively charged
Exist in the region around the nucleus
 Discovery of Protons and Neutrons

• Rutherford’s experiments also told him the amount of positive

nuclear charge.
• The positive charge was carried by particles that were named
protons.
• The proton charge was the fundamental unit of positive charge.
• The nucleus of a hydrogen atom consists of a single proton.
 Discovery of Protons and Neutrons
• Scientists introduced the concept of atomic number,
which represents the number of protons in the
nucleus of an atom.
• James Chadwick discovered neutrons in the
nucleus, which have nearly the same mass as
protons but are uncharged.
Proton
A nuclear particle having a positive charge equal to that of the
electron and a mass more than 1800 times that of the electron

The number of protons in an atom is called the atomic number,

or the proton number, Z.
The number of electrons in the atom is also equal to Z because
the atom is neutral.
The mass number, A, is the total number of protons and neutrons
in the nucleus.
The number of neutrons, neutron number, is A-Z.
An element is a substance whose atoms have the same number of
protons and thus the same atomic number, Z.
Neutron
A nuclear particle having a mass almost equal
to that of the proton but no electrical charge

The mass number, A, is the total number of

protons and neutrons in the nucleus.

Isotopes are atoms whose nuclei have the

same atomic number (number of protons) but
different numbers of neutrons (mass number).
 2-4 Chemical Elements
All atoms of a particular element have the same atomic number, Z.
Known elements atomic numbers from Z=1 to Z=112.
Each element has a name and distinctive symbol.
Chemical Symbols are one or two-letter abbreviations of the name (usually
english name).
The first letter of the symbol is capitalized; for example: carbon,C;
oxygen,O neon, Ne; and silicon, Si.
Some elements known as Latin names, Fe for iron (Ferrum) and Pb for
lead (plumbum).
Sodium has Na symbol based on Latin Natrium.
Potassium has K symbol based on Latin Kalium.
Tungsten has W symbol based on german Wolfram.
Elements beyond uranium (Z=92) do not occur naturally.
 Isotopes
An isotope is specified by its atomic # and its mass #. The
notation used to designate isotopes is the chemical symbol of the
element written with its atomic # as a left subscript and its mass
# as a left superscript.

number p + number n
A
Z
E Symbol of element
number p
E.g.
Mass #
12
6
C Symbol of element
Atomic #
𝟐𝟕
𝟏𝟑𝐀𝐥 13 proton, 14 neutron and 13 electron
 Isotopes
Contrary to what Dalton thought, we know that atoms of an element do
not necessarily all have the same mass. In 1912, J. J. Thomson
measured the mass-to charge ratios of positive ions formed in neon gas.
He found that about 91% of the atoms had one mass and that the
remaining atoms were about 10% heavier. All neon atoms have 10
protons in their nuclei, and most have 10 neutrons as well. A few neon
atoms, however have 11 neutrons and some have 12.
Percent natural abundance
𝟐𝟎
𝟏𝟎𝐍𝐞 90.05%
𝟐𝟏
𝟏𝟎𝐍𝐞 0.27%
𝟐𝟐
𝟏𝟎𝐍𝐞 9.22%

Atoms that have the same atomic # (Z) but different mass
numbers (A) are called isotopes.
 Ions
When atoms lose or gain electrons, they are called ions and carry
net charges.
An atom that gains extra electrons becomes a negatively charged
ion, called an anion.
An atom that loses electrons becomes a positively charged ion,
called a cation.
The number of protons does not change when an atom becomes
an ion.

20Ne+ 10 protons, 10 neutrons and 9 electrons

22Ne2+ 10 protons, 12 neutrons and 8 electrons
 The charge on an ion is equal to the # of protons minus the # of
electrons.
number p + number n
A +? number p - number e
Z
E
number p

16O2- 8 protons (atomic number)

8 neutrons (mass number - atomic number)
10 electrons
Charge: 10- 8= -2

E.g. Determine numbers of electrons in Mg2+ cation and the S2- anion?
Mg2+ 12-number e = +2 number e =10
S2- 16-number e = -2 number e =18
 Isotopic Masses
Determining the masses of individual atoms must be done by
experiment. By international agreement, one type of atom has been
chosen and assigned a specific mass.

This standard is an atom of the isotope carbon-12, which is assigned

a mass of exactly 12 atomic mass unit, 12amu.

The masses of other atoms relative to carbon-12 are determined by

mass spectrometer. All atomic masses are given relative to the
mass of carbon-12 isotope.
Mass of one 12C atom = 12 amu (exactly)
1 amu = Mass of one 12C atom = 1.660539 x 10-24 g
12
 A Mass Spectrometer
A beam of gaseous ions passing through electric and magnetic
fields separates into components of differing masses. The
separated masses are focused on a measuring instrument that
records their amounts.

Heavy ions are

Light ions are deflected a little bit.
deflected greatly.
Ions are separated
according to mass.

Mass spectrum of
mercury,Hg.
Stream of positive
ions