Atomic Theory of Matter
We now take for granted the idea that all matter is
comprised of atoms.
But how did the Atomic Theory of Matter develop, and how
was it proved?
The notion that matter was made up of atoms had to be
postulated and proved.
Experimental Evidence
Experimental evidence for the atomic nature of matter was
realized in the 18th century.
Since then scientists have proved conclusively the atomic
nature of matter
In the mid ‘80’s a tool was developed which for the first
time allowed scientists to actually “see” individual atoms
and molecules.
Surface of graphite as
imaged by a scanning
tunneling microscope
http://www.columbia.edu/cu/chemistry/groups/flynn/
Law of Conservation of Mass
1775 - Lavoisier
“Father of Modern Chemistry”
In every chemical operation an
equal amount of matter exists
before and after the operation.
Mass is conserved, the total mass
after the chemical operation must
be the same as that before.
Problem
Potassium chlorate (KClO3) decomposes to potassium
chloride (KCl) and oxygen (O2) when heated. In one
experiment 100.0 g of KClO3 generated 36.9g of O2 and 57.3 g
of KCl. What mass of KClO3 remained unreacted?
Mass of KClO3 before reaction = mass of KCl + mass of O2
+ mass of unreacted KClO3
100.0 g of KClO3 = 57.3 g KCl + 36.9g O2 + g unreacted KClO3
g unreacted KClO3 = 100.0 g - 57.3 g - 36.9 g
= 5.8 g
Law of Definite Proportions
Joseph Proust
In a given chemical compound, the proportions by mass
of the elements that compose it are fixed, regardless of
the source of the compound.
The ratio of elements in a compound is fixed regardless
of the source of the compound.
Water is made up of 11.1% by mass of hydrogen and
88.9% oxygen.
Problem
In a set of experiments very pure tin (Sn) was combined
with bromine (Br) forming tin tetrabromide (SnBr4). Using
the data below, confirm the law of definite proportions by
calculating the % of tin in each sample of SnBr4.
Grams of Sn reacted
Grams of SnBr4 formed
2.8445
10.4914
3.0125
11.1086
4.5236
16.6752
Need to determine (mass of Sn reacted)
(mass of Br reacted)
Mass of Br reacted =
Mass of SnBr4 formed - mass of Sn reacted
Grams of Sn reacted
2.8445
3.0125
4.5236
Grams of Br reacted
10.4914 - 2.8445 = 7.6469
11.1086 - 3.0125 = 8.0961
16.6752 - 4.5236 = 12.1516
(mass of Sn reacted)
(mass of Br reacted)
0.3721
0.3721
0.3723
Dalton’s Atomic Theory
•All matter consists of indivisible atoms
• Atoms of one kind of element are identical in mass and
properties; atoms of different kinds of elements are
different
•Compounds are made up of definite numbers of atoms of
the component elements
•The weight of a compound equals the sum of the weights
of the component elements
How do atoms combine to form compounds?
Gay Lussacs - Law of Combining Volumes
For chemical reactions involving gases, combinations
occur in simple proportions by volume. Moreover, the ratio
of the volume of each product gas to the volume of either
reacting gas is a ratio of small integers.
For example, if the pressure and temperature are kept
constant, two volumes of H2 gas reacts with one volume of
O2 gas, producing two volumes of water vapor. The two
reactants and the product form a 2:1:2 ratio.
Based on experimental observation:
2 volumes of hydrogen + 1 volume of oxygen  2 volumes
of water vapor
Avogadro’s Hypothesis
Equal volumes of different gases (at the same temperature
and pressure) contain equal numbers of particles
2 volumes of hydrogen + 1 volume of oxygen  2 volumes
of water vapor can be expressed as
2H2 + O2  2H2O
While at this time there was no direct evidence to show that
hydrogen and oxygen gas were H2 and O2, 50 years later
this was proven to be the case.
Are Atoms really indivisible?
Dalton reached his conclusions about atoms on the basis
of evidence gained on a macroscopic level.
As scientists developed more instrumentation capable of
probing phenomena at a microscopic level, more about
atoms was understood.
Example: The color of the emitted light characterizes the
element.
Electrons
J.J. Thompson showed that the cathode rays were in fact
particles of NEGATIVE charge, the rays could be deflected
by a magnetic field.
The term ELECTRON was coined for the negative particles.
Thompson also calculated the charge of each particle
Millikan’s Oil Drop Experiment
Robert Millikan’s oil drop experiment calculated the
charge/mass ratio of the electron, and combining
Thompson’s results the mass of the electron was
calculated to be 9.10 x 10-28 g.
(actual mass of the electron 9.10939 x 10 -28 g)
There must be a positive species which counters the
electron charge.
Radioactivity
Henri Becquerel in 1896 discovered high-energy radiation
was spontaneously emitted from uranium.
Later Marie Curie and her husband Pierre further
investigated this spontaneous emission of radiation which
was termed radioactivity.
Further studies of radioactivity by Rutherford showed that
the radiation consisted of three types of radiation , b, g
radiation.
 and b radiation are bent by an electric field, but in opposite
directions, and g radiation is unaffected.
Rutherford was able to show that  particles have charge of
+2 and that b –1 and that  particles combine with electrons
to form Helium atoms.
g rays are high energy electromagnetic radiation
The Structure of the Atom
J.J. Thompson, realized that electrons were sub-atomic
particles, and presented his theory of the model of the
atom.
The “PLUM-PUDDING” model
Rutherford’s conducted further experiments which
contradicted Thompson’s model.
To explain his results Rutherford postulated
Most of the mass of the atom and all its positive charge
was located in a concentrated core, called the nucleus.
Most of the total volume of the atom is empty in which
electrons move around the positive core.
Model of the Atom
Since the times of Rutherford, many more subatomic
particles have been discovered.
However, for chemists three sub-atomic particles are all that
we need to focus on – ELECTRON, PROTON, NEUTRON.
Electrons are –1, protons +1 and neutrons are neutral.
Atoms have an equal number of electrons and protons they
are electrically neutral.
Protons and neutrons make up the heavy, positive
core, the NUCLEUS which occupies a small volume
of the atom.
Isotopes
Atoms of different elements are distinguished by the
difference in the number of protons and the number of
electrons.
Since atoms are electrically neutral, the number of protons =
number of electrons
Since the number of protons (and electrons) differ, their
MASS differ. Hence atoms of different elements have
different masses.
So for example, hydrogen has ONE proton and ONE electron
Carbon has SIX protons and SIX electrons
Atoms of the SAME element can have different number of
NEUTRONS.
These atoms of the SAME elements but with different
number of neutrons are called ISOTOPES.
Hence isotopes of the same elements have the same
number of protons and electrons, but different number of
neutrons and hence different masses.
EXAMPLE – Carbon has three isotopes C12, C13, C14. Each
of these isotopes differ by the number of neutrons – ALL
have SIX protons. C12 has SIX neutron, C13 has SEVEN and
C14 has EIGHT.
To denote the number of protons and neutrons in an atom
the following symbol notation is used
12
6C
where 12 denotes SUM OF PROTONS + NEUTRONS
6 denotes the number of PROTONS
So for the isotopes of carbon the complete chemical
symbols are: 126C, 136C, 146C
The superscript, which is the sum of the number of
protons and neutrons, is called the MASS NUMBER (A).
The subscript indicates the number of protons and is called
the ATOMIC NUMBER (Z) .
How many protons, neutrons, and electrons are there in
197
79 Au
197 is the mass number and refers to the total number of
protons and neutrons. 79 is the atomic number and
refers to the number of protons. Hence this atom has 79
electrons and 197-79 = 118 neutrons
Atomic Units
Atoms are very, very light and very, very small
Since atomic dimensions are so small, it would be
cumbersome to use units we typically use for length (cm,
m) or mass (g).
Hence, on the atomic scale we define units appropriate for
this scale
MASS – unit typically used is an ATOMIC MASS UNIT (amu)
1 amu = 1.66054 x 10–24 g
Particle
Charge
Mass (g)
Mass (amu)
Proton
+1
1.6727x10-24
1.0073
Neutron
Neutral
1.6750x10-24
1.0087
Electron
-1
9.109x10-28
5.486 x 10-4
LENGTH – ANGSTROM(Å) = 10-10m
Typical atomic dimensions are 1 to 5 x10-10 m which
corresponds to 1 to 5 Å.
Relative Atomic Mass
A relative scale has been developed to compare the relative
masses of atoms. The ATOMIC MASS of an atom is its
relative mass on this scale.
Carbon- 12 (12C) has been set as the standard and assigned
a RELATIVE MASS of exactly 12.
Relative atomic masses have no units since they are the
ratio of two masses.
Average Relative Atomic Mass
Because the abundance of the isotopes of different
elements are essentially constant, we can define an
AVERAGE RELATIVE ATOMIC mass
Average Relative Atomic Mass = average mass of atoms of
an element =
(Abundance)A(Mass)A + (Abundance)B(Mass)A + …
(Table at the back of the text lists relative atomic masses of elements)
Problem
Naturally occurring chlorine has two isotopes, 3517Cl, 3717Cl.
The 35-Cl isotope has a relative atomic mass of 34.9688
and an abundance of 75.77% and the 37-Cl isotope has a
relative atomic mass of 36.9659 and an abundance of
24.23%. Calculate the average atomic mass of Cl.
Average Atomic Mass of Cl = (0.7577x34.9659) +
(.2423x36.9659)
= 35.4527
Average relative atomic mass of C is 12.0107 accounting
for 12C (98.892%, relative atomic mass 12.000000) and 13C
(1.108%, relative atomic mass 13.003354)
Relative Molecular Mass
The relative molecular mass is the sum of the relative
atomic masses of the atoms that make up the molecule.
Example, the chemical formula for water is H2O
Its relative molecular mass
= 2 (1.00794) + 15.9994 = 18.0153
Avogadro’s Number
Avogadro’s number: the number of atoms in exactly 12 g of
12C.
No = 6.022137 x 1023
Sodium (Na) has a relative atomic mass of 22.98977
Hence a sodium atom is (22.98977) times as heavy as
12
12C
If No atoms of 12C have a mass of 12g then, the mass of No
atoms of sodium must be
(22.98977) 12g = 22.98977 g
12
The mass, in grams, of No atoms of ANY element is
numerically equal to the relative atomic mass of that
element.
Same applies to molecules.
Since the relative molecular mass of water is 18.0152, the
mass of No water molecules is 18.0152g
MOLE: A mole has been defined as a unit containing 6.022137
x 1023 , Avogadro’s number, atoms or molecules,
One mole of any atom or molecule contains the same number
of atoms or molecules
The mass, in grams, of ONE MOLE of atoms or molecules is
numerically equal to relative atomic or molecular mass.
Hence 1 mole of Na weighs 22.9898 g, 1 mole of H2O weighs
18.0153 g
The MASS of one mole of atoms or molecules is called its
MOLAR MASS and has UNITS of g/mol
PROBLEM
How many moles of Fe are there in 8.232 g of Fe?
How many atoms are there in 8.232 g of Fe?
Moles of Fe = 8.232 g Fe x 1 mole
= 0.1474 mol Fe
55.85 g Fe
How many grams of water are there in 0.2000 moles of
water?
0.2000 mol H2O x 18.015 g H2O = 3.603 g H2O
1 mol H2O