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2
The Atomic Nature
of Matter
Introduction: The Evolution of Atomic Theory
2.1 - Atomic Theory
2.2 - Atomic Architecture: Electrons and Nuclei
2.3 – Atomic Diversity: The Elements
2.4 – Charged atoms: Ions
2.5 – Energy of Atoms and Molecules
The Greek Concept of Atomos:
The Indivisible Atom
• Around 440 BC, Leucippus originated the atom concept.
• His pupil, Democritus (c460-371 BC) extended it
• There are five major points to their atomic idea.
• All matter is composed of atoms, which are too small to be seen.
These atoms CANNOT be further split into smaller portions.
• There is a void, which is empty space between atoms.
• Atoms are completely solid.
• Atoms are homogeneous, with no internal structure.
• Atoms can differ in size, shape, and weight
Aristotle (384-322 BC)

Spoke openly against the concept
—


Only a few scholars gave it much thought.
The Catholic Church accepted Aristotle's position
—


the atom concept diminished
equated atomistic ideas with Godlessness
It was not until 1660 that Pierre Gassendi succeeded
in separating the two
not until 1803 that John Dalton put the atom on
a solid scientific basis.
John Dalton 1803-1807
Modern Atomic Theory
1) All matter is composed of tiny particles called
atoms.
2) All atoms of a given element have identical
chemical properties that are characteristic of
that element.
3) Atoms form chemical compounds by
combining in whole-number ratios.
4) Atoms can change how they are combined,
but they are neither created nor destroyed in
chemical reactions.
2.1 Atomic Theory
Conservation of Atoms & Mass
Atoms are neither created nor destroyed during
physical or chemical processes.
Fig 2-4
2.1 Atomic Theory
Conservation of Atoms & Mass
Mass is neither created nor destroyed during physical
or chemical processes.
Fig 2-5
Courtesy Patrick Watson
2.1 Atomic Theory
Atoms & Molecules are Continually in Motion
Diffusion: the gradual mixing of atoms & molecules due to their
continual motion.
Fig 2-7
2.1 Atomic Theory
Atoms & Molecules are Continually in Motion
Fig 2-9
Courtesy Patrick Watson
2.1 Atomic Theory
Atoms Combine to Make Molecules
Fig 2-6
2.1 Atomic Theory
Dynamic Molecular Equilibrium
the condition in which a forward and reverse process occur at
equal rates, so the system undergoes no net change.
Fig 2-10
When this example is at dynamic equilibrium, the number of
molecules vaporizing equals the number of molecules
condensing at any given moment.
Laws of Mass Conservation &
Definite Composition
Law of Mass conservation: The total mass of substances
does not change during a chemical reaction.
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.
Mass Percent Composition of Na2SO4
Na2SO4 = 2 atoms of Sodium + 1 atom of Sulfur + 4 atoms of Oxygen
Elemental masses
Percent of each Element
2 x Na = 2 x 22.99 = 45.98
1 x S = 1 x 32.07 = 32.07
4 x O = 4 x 16.00 = 64.00
142.05
Check
% Na = Mass Na / Total mass x 100%
% Na = (45.98 / 142.05) x 100% =32.37%
% S = Mass S / Total mass x 100%
% S = (32.07 / 142.05) x 100% = 22.58%
% O = Mass O / Total mass x 100%
% O = (64.00 / 142.05) x 100% = 45.05%
% Na + % S + % O = 100%
32.37% + 22.58% + 45.05% = 100.00%
Calculating the Mass of an Element
in a Compound Ammonium Nitrate
How much Nitrogen is in 455 kg of Ammonium Nitrate?
Ammonium Nitrate = NH4NO3 The Formula Mass of Cpd is:
4 x H = 4 x 1.008 = 4.032 g
2 x N = 2 X 14.01 = 28.02 g
3 x O = 3 x 16.00 = 48.00 g
Therefore gm Nitrogen/ gm Cpd
80.052 g
28.02 g Nitrogen
= 0.35002249 g N / g Cpd x 100 = 35.00%
80.052 g Cpd
455 kg x 1000g / kg = 455,000 g NH4NO3
455,000 g Cpd x 35.00 g N / 100g Cpd = 1.59 x 105 g Nitrogen
or: 455 kg NH4NO3
X 28.02 kg Nitrogen = 159 kg Nitrogen
80.052 kg NH4NO4
Law of Multiple Proportions
If elements A and B react to form two compounds,
the different masses of B that combine with a fixed
mass of A can be expressed as a ratio of small whole
numbers:
Example: Nitrogen Oxides I & II
Nitrogen Oxide I : 46.68% Nitrogen and 53.32% Oxygen
Nitrogen Oxide II : 30.45% Nitrogen and 69.55% Oxygen
Cmpd I
Cmpd II
in 100 g of each Compound: g O = 53.32 g & 69.55 g
g N = 46.68 g & 30.45 g
g O /g N = 1.142 & 2.284
Cmpd II 2.284
2
=
Cmpd I 1.142
1
2.2 Atomic Architecture: Electrons & Nuclei
Forces
Gravitational force: the force which pulls object toward the
center of the Earth.
Electrical force: the
attraction or repulsion
between two charged
objects.
Fig 2-12
2.2 Atomic Architecture: Electrons & Nuclei
Forces
Gravitational force: the force which pulls object toward the
center of the Earth.
Electrical force: the attraction or repulsion between two
charged objects.
Magnetic force: the force generated by charged objects in
motion.
Fig 2-12
Courtesy Patrick Watson
Davy, Faraday
1807 Humphry Davy:
• Forces holding compounds together are
electrical
1833 Michael Faraday:
—
Atomic mass and the electricity
needed to free elements from
compounds are related
More Important developments
• 1864 Heinrich Geissler: developed a pump
• produced good vacuums in sealed glass tubes
• 1870’s William Crookes: Cathode rays
• are negatively charged
• same regardless of cathode metal
• are particles with mass
• 1891 George Stoney
• electricity exists in units called "electrons”
2.2 Atomic Architecture: Electrons & Nuclei
Electrons
Fig 2-13 Gas Discharge Tube
Conclusions:
•Atoms are made up of smaller positive and negative fragments.
•The negatively charged particles are electrons, which are
uniform in behavior, regardless of their source.
1897 J.J. Thomson
DISCOVERED THE ELECTRON
1897 J.J. Thomson
• Deflected an electron beam by both magnetic
and electric attraction/repulsion
• Measured the electron’s mass/charge ratio
2.2 Atomic Architecture: Electrons & Nuclei
Electrons
Fig 2-14 Cathode Ray Tube
Conclusion:
Charge
e

  1.76x1011 C/kg
Mass
m
Cathode Rays
•
•
•
•
Attracted to the positive electrode
Not visible but could make things “glow”
Traveled in a straight line
Could be bent by electric or magnetic
fields
• A plate in it’s path acquired a negative
charge
• Same regardless of material
The quest continues
• 1905 Albert Einstein: Photoelectric Effect
• light causes electrons to be emitted
from metals
• quantized energy transfer causes the
emission
—E=mc2
He was SOOOOO excited!!
2.2 Atomic Architecture: Electrons & Nuclei
Electrons
Charge = n (-1.6x10-19 C)
Charge
e

  1.76x1011 C/kg
Mass
m
Conclusions:
•Electrons are particles.
•Electrons have a mass
of 9.1 x 10-19 kg
Fig 2-15 Millikan’s Oil Drop Experiment
Nucleus Discovered

1911 Ernest Rutherford:
small
dense positive nucleus
nucleus is most of the mass of an atom
electrons are in the space around the nucleus
Ernest Rutherford (1871-1937)
• Won the Nobel Prize in Chemistry
in 1908
• “It was quite the most incredible
event..... It was almost as if a gunner
were to fire a shell at a piece of tissue
and the shell bounced right back!!!!! ”
2.2 Atomic Architecture: Electrons & Nuclei
The Nucleus
• Most of the mass & all of the
positive charge of the atom
are in the nucleus, which
occupies only 1 part in 1014 of
the atoms volume.
• Electrons occupy a huge
volume in comparison to the
nucleus, but have relatively
small masses.
Fig 2-17 Schematic drawing of an atom
2.2 Atomic Architecture: Electrons & Nuclei
The Nucleus
• Protons account for the
positive charge of the nucleus
and have a positive charge
with a magnitude equal to the
negative charge of an
electron. Protons have a
mass about 2000 times that
of an electron.
• The mass of a neutrons is
about the same as the mass
of a proton, but a neutron is
electrically neutral.
Fig 2-17 Schematic drawing of an atom
2.2 Atomic Architecture: Electrons & Nuclei
The Nucleus
Table 2-1 Atomic Building Blocks
Name
Symbol
Electron
Proton
Neutron
e
p
n
Charge
-1.6022 x10-19 C
+1.6022 x10-19 C
0
Mass
9.1091 x10-31 kg
1.6726 x10-27 kg
1.6749 x10-27 kg
2.3 Atomic Diversity
An element is identified by the charge of its nucleus
Atomic number, Z: nulcear charge, number of
protons
Atomic Definitions I: Symbols,
Isotopes,Numbers
A
X
Z
The Nuclear Symbol of the Atom, or Isotope
X = Atomic symbol of the element, or element symbol
A = The Mass number; A = Z + N
Z = The Atomic Number, the Number of Protons in the Nucleus
N = The Number of Neutrons in the Nucleus
Isotopes = atoms of an element with the same number of protons,
but different numbers of Neutrons in the Nucleus
2.3 Atomic Diversity
Isotopes
Atom with the same number of protons, but
different number of neutrons.
Mass number A
Atomic number  Z
X Element symbol
Example: How many protons, neutrons and electrons
do each of the following have?
16
8
O
12
6
C
14
6
C
2.3 Atomic Diversity
Isotopes
A
ZX
Examples:
16
O
8
12 C
6
14
6C
8 protons, 8 neutrons, 8 electrons
6 protons, 6 neutrons, 6 electrons
6 protons, 8 neutrons, 6 electrons
2.3 Atomic Diversity
Isotopes
Courtesy of Sachtleben Chemie GmbH (Paint)
Courtesy M. Freeman/PhotoLink/Phoyo Disc (Knee Joint)
46Ti
= 8.2%
47Ti
= 7.4%
48Ti
= 73.8%
49Ti
= 5.4%
50Ti
= 5.2%
2.3 Atomic Diversity
Isotopes
Cl
Cr
Ge
Sn
Fig 2-18 Natural abundance of the isotopes of Cl, Cr, Ge, Sn
Problem: Calculate the abundance of the two Bromine isotopes:
79Br = 78.918336 g/mol and 81Br = 80.91629 g/mol , given that
the average mass of Bromine is 79.904 g/mol.
Plan: Let the abundance of 79Br = X and of 81Br = Y and X + Y = 1.0
Solution:
X(78.918336) + Y(80.91629) = 79.904
X + Y = 1.00 therefore X = 1.00 - Y
(1.00 - Y)(78.918336) + Y(80.91629) = 79.904
78.918336 - 78.918336 Y + 80.91629 Y = 79.904
1.997954 Y = 0.985664
or Y = 0.4933
X = 1.00 - Y = 1.00 - 0.4933 = 0.5067
%X = % 79Br = 0.5067 x 100% = 50.67% = 79Br
%Y = % 81Br = 0.4933 x 100% = 49.33% = 81Br
Pg 55 2.3.2 – 2.3.3
Modern Reassessment of the Atomic Theory
1. All matter is composed of atoms. Although atoms are composed
of smaller particles (electrons, protons, and neutrons), the atom
is the smallest body that retains the unique identity of the element.
2. Atoms of one element cannot be converted into atoms of another
element in a chemical reaction. Elements can only be converted into
other elements in Nuclear reactions in which protons are changed.
3. All atoms of an element have the same number of protons and
electrons, which determines the chemical behavior of the element.
Isotopes of an element differ in the number of neutrons, and thus
in mass number, but a sample of the element is treated as though
its atoms have an average mass.
4. Compounds are formed by the chemical combination of two or more
elements in specific ratios, as originally stated by Dalton.
2.4 Charged Atoms: Ions
Formation of Cations
Ions: electrically charged atomic or molecular particles
Cations: ions with positive charges
2.4 Charged Atoms: Ions
Formation of Anions
Anions: ions with negative charges
Net electrical charge is always conserved
2.4 Charged Atoms: Ions
Ionic Compounds
A solid containing cations
& anions in a balanced
whole-number ratio.
2 Na(s) + Cl2(g) 2 NaCl(s)
Fig 2-21
Courtesy Michael Watson
Fig. 2.19
2.4 Charged Atoms: Ions
Salt + water
Ionic Solutions
Pure water
Fig 2-23
Sugar + water
Courtesy Ken Karp
2.4 Charged Atoms: Ions
Ionic Solutions
+
-
Fig 2-24
2.5 Energy of Atoms and Molecules
Fig 2-24
2.5 Energy of Atoms and Molecules
Kinetic Energy, Ekinetic: the energy of directed motion
of an object.
Ekinetic = 1/2 mu2
(2-1)
1 kg m2s-2 = 1 J
Thermal Energy: the energy of random motion,
translational, rotational and vibrational. The thermal
energy of an object is equal to the sum of the kinetic
energy of its atoms.
2.5 Energy of Atoms and Molecules
Potential Energy: energy that is stored
Chemical Energy: potential energy stored as
chemical bonds
Electrical Energy: potential energy that is the result of
electrical forces between charged objects
q1q2
Eelectric  k
r
(2-2)
Radiant Energy: the energy of electromagnetic radiant
(light, photons)
2.5 Energy of Atoms and Molecules
Conservation of Energy
Energy is neither created nor destroyed in any
process only transferred from one body to another,
or changed from one form to another.