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Chapter-3
Atoms: The
Building Blocks of
Matter
I. Early Theories of Matter
A. Greek Philosophers
1. Democritus and Leucippus
I. Early Theories of Matter
“Atomos” indivisible
Democritus, 400 BC
“Everything is made up of a few simple parts
called atomos.”
Atomos means “uncuttable or indivisible” in
Greek
He envisioned atomos as small, solid particles of
many different sizes and shapes
*His ideas were rejected later by Aristotle
I. Early Theories of Matter
A. Greek Philosophers
1. Democritus and Leucippus
2. Aristotle
I. Early Theories of Matter
 Matter IS infinitely divisible. He was so influential and
this is what people believed for the next 2000 years!
 He believed that all matter was “continuous”
Why? Aristotle invented
logic:
Terms: "A", "B", "rabbit", "Peter"
Propositions: "A" is "B" or "A" is not "B“
"Peter is a Rabbit"
Syllogisms: "A is B. B is C.
Therefore, A is C."
Two kinds of Syllogisms:
"All rabbits are brown. Peter is a rabbit.
Therefore, Peter is brown. (Induction)
"Flopsy, Mopsy, Cottontail and Peter are
brown. Flopsy, Mopsy, Cottontail and Peter
are rabbits. Therefore, all rabbits are brown."
(Deduction – what scientists do)
B. Alchemists
Contributions of Alchemists
• Invention of Lab
Equipment
• Important Chemical
Procedures
• Discovery of Elements
• Descriptive Chemistry
C. Lavoisier & Proust
Lavoisier – Law of
Conservation of Mass
C. Lavoisier & Proust
2. Proust – Law of
Definite Proportions
Law of Definite Proportions
 Compounds always use specific proportions
of their elements, regardless of what is
available.
 Water will always contain:
 11.2% H
&
88.8% O by mass
 Hydrogen peroxide contains:
 5.9% H
&
94.1% O by mass.
 Different ratios mean different properties!
Percent Mass
part
× 100
% Mass =
whole
mass of element
× 100
% Mass =
mass of compound
Practice Problem
 Two samples of compounds made of
only H and O have the following
compositions:
Compound A: 15.0 g Hydrogen
120.0 g Oxygen
Compound B:
 Are they same ?
2.0 g Hydrogen
34.0 g Oxygen
1 × 16.00
=
1 × 12.00
2 × 16.00
=
1 × 12.00
Atomic Theory – John Dalton 1808
3. Dalton – Law of
Multiple Proportions
Law of Multiple Proportions
 When two different compounds are made of the
same set of elements, different masses of one
element combine with the same mass of the
other in small, whole number ratios.
HUH?
 Example: Carbon dioxide CO2 has exactly
twice as much oxygen as carbon monoxide CO.
Law of Multiple Proportions Example
Law of Multiple Proportions
Dalton’s Atomic Theory
1) Matter is made of tiny particles called atoms
2) All atoms of an element are:
identical in size, mass and other properties different
from those of the other elements
3) Atoms cannot be subdivided, created or destroyed
4) Atoms of one element can combine with atoms of other
elements to form compounds. A given compound
always has the same relative number and types of atoms
5) In chemical reactions, atoms are combined, separated or
rearranged
 Modern Theory:
• Atoms of an element have a characteristic average
mass which is unique to that element, but elements
CAN have different isotopes!
• Atoms cannot be subdivided, created, or destroyed
in ordinary chemical reactions. But, changes CAN
occur in nuclear reactions
Dalton actually did have evidence
from experiments:
Mass is conserved in chemical
reactions
Elements combine to make compounds
in predictable mass ratios:
Law of Definite Proportions
Law of Multiple Proportions
Summary of Law’s of 1790’s
The transformation of a substance(s) into one or more new
substances is called “chemical reaction”
Law of conservation of mass: mass is neither created nor
destroyed during ordinary chemical reactions or physical
changes.
Law of definite proportions: A pure compound always
contains the same elements in the same proportions by mass.
Law of multiple proportions: When elements combine to
form more than a single compound, the ratios of the masses
of the combining elements can be expressed by a ratio of
small whole numbers.
Dalton’s Model/Chart
Today we can even see atoms!
A penny contains
29,000,000,000,000,000,000,000 atoms
III. Dalton’s model vs. Modern Model
How has it changed?
IV. Finding Subatomic Particles
A. Radioactivity - the phenomenon of rays
being produced spontaneously by unstable
atomic nuclei.
B. X-rays & radioactivity
1. Wilhelm Roentgen - 1895
a. shot cathode rays at metal and
radiation was given off that would
darken photographic film.
b. the rays could pass
through flesh, but not
bone to expose photo
plate.
c. rays were
unknown, so they
were given the name
"x-rays“.
2. Becquerel - obtained a sample of uranium
ore (pitchblende) to study phosphorescence
a. He noticed uranium also darken
photographic paper
b. Radioactive Decay
c. Transmutation - change of one
element into another
3. Marie Curie- Noticed polonium and
radium also gave off radiation:
a. radiation of radium is two
million times greater than uranium.
b. it conducted electricity in air
"ionized air“.
c. ionizing radiation has sufficient
energy to change atom and
molecules into ions.
d. non-ionizing radiation does not,
such as radio waves.
Non-ionizing vs. Ionizing Radiation
C. Discovery of Electrons
1. William Crookes (1879) discovers
cathode rays (streams of electrons)
2. JJ Thomson (1897) – Cathode Ray Tube
Experiment

Cathode ray tube (CRT) to deduce the presence of a
negatively charged particles.
 Discovery of electrons using CRT
 Electron’s charge-to-mass ratio
The beam of negative
particles bends downward
The magnet is turned around and the
beam bends in opposite direction
atoms were not indivisible particles
2. JJ Thomson (1897) – Plum Pudding Model
An electron is much
smaller than a
hydrogen atom?
Dalton was wrong:
atoms can be divided!
2. JJ Thomson (1897)
 All elements should contain electrons
 All atoms are electrically neutral
(so there must be positive particles present)
 Mass of electrons are so small  other particles?
CRT experiment
Electrons discovered
Plum-pudding model
Mass/charge ratio
Mass of electron
(1/2000th of a H atom)
3. Robert A. Millikan (1909) – Oil Drop
Experiment
 Oil drop experiment findings:
 Measured the charge of an electron
(1.602× 10-19 C)
 Calculated the mass of an electron
(9.109 × 10-31 kg)
4. Rutherford, Marsden & Geiger (1911)
4. Ernest Rutherford – Gold foil Experiment
Conclusions–
Nucleus is:
 small
 dense
 positively
charged
atoms are mostly
empty space
4. Ernest Rutherford – Gold foil Experiment
--A question arose:
Why don’t the electrons fall into the nucleus ?
4. Ernest Rutherford – Other Findings
 Characterized radioactivity as 3 types of radiation:
 Alpha rays (): positively charged particles
 Beta rays ():
negatively charged particles
 Gamma rays (): no charge
Nuclear Symbols:
Neutron
Proton
Electron
Alpha
Beta
Gamma
Positron
High energy
Lower energy
High energy
Beta particles and gamma rays are usually emitted together!
5. Eugen Goldstein’s Canal Ray
Experiment (1886)
Goldstein used a perforated anode
in the cathode ray tube and found
that there were these positively
charged rays (canal rays or anode
rays), which consisted on positively
charged particles called as protons.
6. James Chadwick – Discovery of Neutron
(1932)
He discovered that most nuclei also
contain another neutral particle:
neutron.
It is slightly more massive than
proton but has no charge.
D. Bohr’s Model (to be continued…)
-Niels Bohr proposed a
model in which orbiting
electrons don't lose
energy. A change to the
Laws of Physics!
Five Major Atomic Models
1. Dalton’s Atomic Model 1808
2. Plum Pudding Model
1909
3. Rutherford’s Nuclear Model 1910
4. Bohr’s Solar System Model 1913
5. Shrodinger’s Wave Mechanical Model 1927
E. Subatomic Particles of Atom
Nuclear particles are held together by
Nuclear Forces!
Location
e- cloud
Nucleus
Nucleus
Nucleons: particles in the nucleus - Proton (p+) neutron (n0)
E. Subatomic Particles of Atom
Thomson
electrons
Chadwick
neutrons
Rutherford
protons
E. Subatomic Particles of Atom
Of course – scientists have
found even more particles –
here’s a brief introduction!
Did you know that if a
nucleus were the size of the
dot in the exclamation point
at the end of this sentence,
its mass would be
approximately as much as
that of 70 automobiles!
The Modern View of the
Atomic Structure






Atoms: Consist of tiny, dense, positively charged nucleus
surrounded by cloud of (-) electrons
Nucleus: Contains positively charged protons and neutral
neutrons
Atomic number(Z): number of protons (determines the element)
Mass number: sum of protons and neutrons (determines the
isotope)
Isotopes: atoms of an element that differ in the number of
neutrons
Atomic Mass: Average mass of all isotopes based on the
abundance (expressed in amu)
23
sodiu

23
Na
11
Atomic Vocabulary
Mass number = p+ + n0 (in the nucleus)
Element
symbol
A
E
Z
Atomic number = # p +
(= # e– )
Elements are put in order of increasing atomic
number on the periodic table, identifies an element.
Atomic Vocabulary
Carbon
6
C
12.011
name
Atomic
number
Element
symbol
Average
atomic mass
(in amu)
F. Isotopes
 Atoms of one element that have different mass numbers
(different # of neutrons)
 There are 3 naturally occurring isotopes of hydrogen:
Isotope
Name
#p
Protium
#n
% of all H
1
0
99.985
Deuterium
1
1
0.015
Tritium
1
2
< 0.001
Periodic Table by Number of
Stable Isotopes
Neutral Isotopes
Arsenic
18
75
Phosphorous
8
8
33
16 15
18
75
31
Example:
An atom of carbon with 7 neutrons:
13C
6
An atom of lead with 125 neutrons:
207Pb
82
Relative Atomic Masses
It is convenient to use relative atomic mass because
the atoms are too small.
(Oxygen atom weighs 2.657 ×10-23g).
The atomic mass of any nuclide is determined by
comparing it with carbon-12 atom.
One atomic mass unit (amu) is exactly 1/12 of the
mass of one C-12 atom. (1.660540 ×10-27 kg)
Average Atomic Masses of Elements
 Different isotopes will have different masses.
 Most elements consist of mixtures of isotopes.
 Average Atomic Mass is the weighted average of
the element’s average atomic masses of the
naturally occurring isotopes of an element
 There are three naturally occurring isotopes of
uranium:
234U
235U
238U
92
92
92
How to Calculate Average Atomic Masses?
Mass #
39
19
K
40
19
K
41
19
K
Atomic #
Abundance of isotopes?
93.2581 %
0.0117 %
6.7302 %
Atomic Masses of isotopes?
38.963707 amu
39.963998 amu
40.961826 amu
How to Calculate Average Atomic Masses?
%
Symbol
abundance
39
K
amu
decimal
product
93.2581
38.963707 × 0.932581 = 36.3368
0.0117
39.963998 × 0.000117 =
0.00468
6.7302
40.961826 × 0.067302 =
2.7568
19
40
K
19
41
K
19
+
39.0983
Avogadro’s Number
It is defined as the number of particles in one
mole of any pure substance.
This number is found to be 6.022×1023.
The MOLE
One mole of something contains 6.022×1023
particles of that substance.
Example: One mole of Hydrogen atoms
contains exactly 6.022×1023 atoms of
Hydrogen.
Example: 1 mole of H2O has 6.022×1023 water
molecules.
Example: 1 mole of NaCl has 6.022×1023
formula units of NaCl.
Molar Mass
Mass of one mole of a substance. By
definition, a molar mass is the mass of 1 mol
of a substance (i.e., g/mol).
Molar mass can be calculated by expressing
the mass in amu or in grams.
Molar mass of C = 12 g/mol
Mole Relationships
One mole of atoms, ions, or molecules contains
Avogadro’s number of those particles.
One mole of molecules or formula units contains
Avogadro’s number times the number of atoms or ions
of each element in the compound.
Mole is a Bridge to Mass, Particles, Volume
Moles provide a bridge from the molecular
scale to the real-world scale.
Moleville
Molar Mass
Railroad
Mass Junction
Avagodro’s
Number
Particles
Conversions