ATOMIC STRUCTURE
SC Science Standards
• Interpret Dalton’s atomic theory in terms of the Laws of
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Conservation of Mass, Constant Composition, and
Multiple Proportions.
Compare and contrast the contributions of Dalton,
Thomson, Rutherford, Bohr, Planck and Schodinger to the
development of the current atomic model.
Based on the quantum theory, write electron
configurations and orbital notation for the representative
elements.
Use Bohr’s model of the atom to explain the bright line
spectrum in terms of electrons moving between energy
levels.
Describe and identify the regions of the electromagnetic
spectrum in terms of frequency, wavelength and energy.
Atomic Theory
• Democritus…(>2000 years ago)
• Greek philosopher
• First suggested the idea of atoms
• Matter is composed of tiny indivisible particles
• Named these particles “atomos”
• Now called atoms
• Ideas lacked experimental support
Atomic Theory
• John Dalton (1776 – 1844)
• English school teacher
• Studied chemistry
• Particularly interested in meteorology
• Performed experiments
• Studied the ratios that chemicals combine to form compounds
• Formulated the Atomic Theory
Dalton’s Atomic Theory
• All elements are
composed of
submicroscopic indivisible
particles called atoms
• Atoms of the same
element are identical. The
atoms of any one element
are different from those of
any other element.
Dalton’s Theory Cont…
• Atoms of different elements can physically mix together or
can chemically combine with one another in simple wholenumber ratios to form compounds.
• Chemical reactions occur when atoms are separated,
joined, or rearranged. However, atoms of one element are
never changed into atoms of another element as a result
of a chemical reaction.
What is an atom?
• The smallest particle of an element that retains the
properties of that element
• Individual atoms are visible with the proper instrument
Subatomic Particles
• Particles that are smaller than atoms
• Three main subatomic particles
• Protons
• Neutrons
• Electrons
Electrons
• Negatively charged
• Discovered by JJ Thomson in 1897
• Experimented with the flow of electric current through
gases in cathode ray tubes
• Found that the cathode rays were attracted to the metal
plates with a positive charge and repelled by metal plates
with a negative charge
Electrons cont…
• Thomson
• Concluded that cathode rays are composed of negatively charged
particles
• Called these negatively charged particles electrons
• Concluded that electrons are a part of the atoms of every element
• Electron has 1 unit of negative charge
• Electron has mass of about 1/2000 of a hydrogen atom
Thomson’s Plum Pudding Model
Protons
• Positively charged subatomic particle
• Discovered by E. Goldstein in 1886
• Has one unit of positive charge
Neutrons
• Discovered by Sir James Chadwick in 1932
• Subatomic particle with no electric charge
• Mass is equal to the mass of a proton
Structure of the atom
• Ernest Rutherford (1871 – 1937)
• Performed famous gold foil experiment
• Tested popular theory that atoms were composed of evenly
distributed protons and electrons
• Experimented with alpha particles (+ charges) aimed at a thin sheet
of gold foil
• Most particles went straight through
• Some (a very few) were bounced back
Rutherford’s experiment
• Rutherford proposed
• Most of the mass and all of the positive
charge of the atom is concentrated in a small
region at the center of the atom
• Called the center region the nucleus
• The nucleus is the center core of the atom
and is composed of protons and neutrons
The Nucleus
• Very small and dense
• If the nucleus were the size of a pea, its mass would be
250 tons!
• Has a positive charge
• Occupies a very small volume of the atom
• Electrons occupy the largest volume of the atom outside
of the nucleus
Atomic Number
• Different numbers of protons make atoms different
• Protons determine the identity of an element
• Atomic number is the number of protons in the nucleus of the
atom
• Each element has a unique atomic number
• Reported on the periodic table
Atoms
• Atoms are electrically neutral
• Number of protons must be equal to the number of
electrons
Mass Number
• Most of an atom’s mass is concentrated within the
nucleus
• Protons and neutrons contribute to the mass
• Mass number = # protons + # neutrons
Mass number
197
79
Au
Atomic number
Isotopes
• Atoms of the same element may have different nuclear
structures
• Number of neutrons may vary within atoms of the same element
• Isotopes are atoms that have the same number of protons, but
different numbers of neutrons
Average Atomic Mass
• Masses of atoms are measured in units called atomic mass
units (amu)
• An atomic mass unit is defined as 1/12 the mass of a carbon12 atom
• The mass of Carbon-12 is 12.000000 amu
• Mass of a single proton or neutron is approximately 1 amu
Atomic Mass
• In nature, most elements exist as a mixture of 2 or more isotopes
• Each isotope has a fixed mass and a natural percent abundance
• Atomic mass is the weighted average mass of the atoms in a
naturally occurring sample of the element
Calculating Atomic Masses
• You need to know:
• The number of stable isotopes of the element
• The mass of each isotope
• The natural percent abundance of each isotope
• Masses and relative abundances are values that can be looked up
in chemical reference books
Atomic Mass of Element X
• Element X has two natural isotopes. The isotope with
mass of 10.013 amu has a relative abundance of 19.90%.
The isotope with mass 11.0093 has a relative abundance
of 80.10%. Calculate the atomic mass of this element and
name it.
Nitrogen
mass number
14
15
exact weight percent abundance
14.003074
15.000108
99.63
0.37
Chlorine
mass
number
35
37
exact weight
34.968852
36.965903
percent
abundance
75.77
24.23
Silicon
mass number exact weight percent abundance
28
27.976927
92.23
29
28.976495
4.67
30
29.973770
3.10
mass number exact weight percent abundance
24
23.985042
78.99
25
24.985837
10.00
26
25.982593
11.01
mass number exact weight
92
94
95
96
97
98
100
91.906808
93.905085
94.905840
95.904678
96.906020
97.905406
99.907477
percent abundance
14.84
9.25
15.92
16.68
9.55
24.13
9.63
mass number
112
114
115
116
117
118
119
120
122
124
exact weight percent abundance
111.904826
0.97
113.902784
0.65
114.903348
0.36
115.901747
14.53
116.902956
7.68
117.901609
24.22
118.903310
8.58
119.902200
32.59
121.903440
4.63
123.905274
5.79