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Unit 2 – Atomic
Structure & Nuclear
Chemistry
Part I – Atomic Theory, Subatomic Particles, and Average
Atomic Mass
Part I Key Terms
• Atomic mass - The mass of an atom of a chemical element expressed in atomic mass
units. It is approximately equivalent to the number of protons and neutrons in the atom
(the mass number)
 Average atomic mass – Weighted average of all atoms of a particular element and is
dependent on the mass of isotopes for an element and the relative population of each
isotope
• Bohr model - Devised by Niels Bohr, depicts the atom as a small, positively charged
nucleus surrounded by electrons that travel in circular orbits around the nucleus
• Dalton’s Postulates - States that matter is composed of extremely small particles called
atoms; atoms are invisible and indestructable; atoms of a given element are identical in
size, mass, and chemical properties; atoms of a specific element are different from those
of another element; different atoms combine in simple whole-number ratios to form
compounds; in a chemical reaction, atoms are separated, combined, or rearranged
• Isotope -Atoms of the same element with different numbers of neutrons
Part I Key Terms (cont.)
• Isotope notation - Subscripts and superscripts can be added to an element’s symbol to
specify a particular isotope of the element and provide other important information. The
atomic number is written as a subscript on the left of the element symbol, the mass
number is written as a superscript on the left of the element symbol
• Mass number - The total number of protons and neutrons in a nucleus.
• Subatomic particles - The three kinds of particles that make up atoms: protons,
neutrons, and electrons
• Theory - An explanation supported by many experiments; is still subject to new
experimental data, can be modified, and is considered valid if it can be used to make
predictions that are proven true
Early Development of Atomic Theory
• Major Contributors to Understanding Atomic Structure
• Democritus – ancient Greek philosopher that originally stated all
matter consists of atoms
• 1605: Francis Bacon – published the scientific method
• 1803: John Dalton – Postulates of Atomic Theory
• 1897: J.J. Thomson – Discovery of the negatively charged
electron and the mass to charge ratio of the electron
• 1908 Robert Millikan – Determines the charge of the electron
• 1911: Ernest Rutherford – Discovers positively charged
nucleus
• 1913: Niels Bohr – Theorizes structure of the electron
cloud with energy levels and planetary orbits of electrons
• 1932: James Chadwick – Discovers neutrons
Atomic Theory – John Dalton
• John Dalton
• English physicist
• Experimented extensively with multiple gases and gaseous
compounds
• Contributions – Five Postulates of Atomic Theory
•
•
•
•
1. All matter consists of tiny particles called atoms
2. Atoms are indestructible and unchangeable.
3. Elements are characterized by the mass of their atoms.
4. When elements react, their atoms combine in simple, whole
number ratios.
• 5. When elements react, their atoms sometimes combine in
more than one simple whole, number ratio.
Dalton’s Model of an Atom
• He made no prediction about the construction of atoms
believing them to be solid spheres.
• Conclusions made based on his experiments and postulates:
• Law of the Conservation of Mass – when chemical reactions
occur, the atoms are only rearranged and there is no difference in
mass following a chemical reaction
• Law of Definite Proportions – elements combine in simple, low
number ratios to form compounds (examples – H20, CO2)
• Law of Multiple Proportions –elements combine in different
simple, low number ratios to form different compounds
(examples – H20 and H202; CO and CO2)
Atomic Theory – J.J. Thomson
• Discovered the negatively charged electron and the mass to
charge ratio of the electron
• Used cathode ray tube
• Beam of electrons deflected toward positive plate indicated
the electron has negative charge
• Amount of deflection indicates the mass to charge ratio
Thomson’s Experiment
Image used courtesy of http://www.chemteam.info/AtomicStructure/Disc-of-Electron-Images.html
Thomson’s Model of the Atom
• Plum Pudding Model
Image used courtesy of http://www.kutl.kyushu-u.ac.jp/seminar/MicroWorld1_E/Part2_E/P24_E/Thomson_model_E.htm
Atomic Theory – Ernest Rutherford
• Discovered positively charged nucleus
• Used gold foil & detector ring
• Fired alpha particles at foil which are positively charged
• Most went through – atom mostly empty space
• Some deflected – nucleus positively charged
• Some bounced back – solid mass indicates nuclear core
Rutherford’s Experiment
Rutherford’s Model of the Atom
• Nuclear Atomic Model
Image used courtesy of http://www.bbc.co.uk/manchester/content/articles/2008/09/10/100908_rutherford_physics_feature.shtml
Atom Theory – Niels Bohr
• Discovered electrons reside in energy levels with discrete
amounts of energy
• Mathematic modeling
• Needed to explain why negatively charged electrons do not
get absorbed into positively charged nucleus
• Used information from Balmer, Lyman, & Paschen series
• Emission spectra for Hydrogen explained by Rydberg equation
Bohr’s Model of the Atom
• Electron Shell Model
Image used courtesy of http://www.blurtit.com/q982327.html
2 Regions of the Atom
 Nucleus
 Contains the protons and neutrons
 Accounts for virtually all of the mass, but only a very small
portion of the volume of the atom.
 Has a positive charge equal to the number of protons.
 Electron Cloud
 Contains the electrons in orbitals
 Has virtually no mass, but accounts for virtually all of the volume
 Has a negative charge equal to the number of electrons.
Subatomic Particles
 Electrons
 Charge = -1
 Mass ≈ 0 amu
 Location: in orbitals in the electron cloud (outside the nucleus)
 Protons
 Charge = +1
 Mass = 1 amu
 Location: Inside the nucleus
 Neutrons
 Charge = 0
 Mass = 1 amu
 Location: Inside the nucleus
Properties of the Atom
 Mass
 Measured in Atomic Mass Units (amu)
 Equal to the sum of the number of protons and neutrons
 Represented by the Mass Number
 Charge
 Neutral unless electrons gained or lost (ionized)
 Number of electrons and protons is equal and, therefore balance
out
 Atomic Number
 Equal to the number of protons
 Define the element and its chemical properties
Symbology
19
9
F
Example assuming neutral atom of
Fluorine
Atomic number:
9
Mass Number:19
Protons:
9
Neutrons:
10
(mass number – atomic number)
Electrons:
9
Isotopes
• Atoms of the
same element
with different
mass due to
different
number of
neutrons
Average Atomic Mass
 Weighted average of all atoms of a particular element
 Dependent on the mass of isotopes for an element and the relative
population of each isotope
Isotope
Oxygen-16
Oxygen-17
Oxygen-18




Isotope Atomic
Mass (amu)
15.99491
16.99913
17.99916
% mass oxygen-16: (15.99491) (.99759)
% mass oxygen-17: (16.99913) (.00037)
% mass oxygen-18: (17.99916) (.00204)
Average Atomic Mass of Oxygen
Population (%)
99.7590
0. 037
0.20400
= 15.9564
= 0.0063
= 0.0367
= 15.9994
Naming Isotopes
 Name of the element followed by the mass number of the
isotope
 Carbon – 12 = the name of the carbon atom with a mass number
of 12 (6 protons and 6 neutrons)
 Carbon – 14 = the name of the carbon atom with a mass number
of 14 (6 protons and 8 neutrons)
 Fluorine – 19 = the name of the Fluorine atom with a mass
number of 19 (9 protons and 10 neutrons)
Energy Levels
• Energy levels correspond to the energy of individual
electrons. Each energy level has a discrete numerical
value.
• Different energy levels correspond to different numbers
of electrons using the formula 2n2 where “n” is the
energy level
Energy Level
1
2
3
4
n
Number of electrons (2n2)
2(12) = 2
2(22)= 8
2(32)= 18
2(42)= 32
2n2
Quantum Mechanical Model of
Atomic Structure
• 1900: Max Planck – Develops law correlating energy to
frequency of light
• 1905: Albert Einstein – Postulates dual nature of light as both
energy and particles
• 1924: Louis de Broglie – Applies dual nature of light to all
matter
• 1927: Werner Heisenberg – Develops Uncertainty Principle
stating that it is impossible to observe both the location and
momentum of an electron simultaneously
• 1933: Erwin Schrodinger – Refines the use of the equation
named after him to develop the concept of electron orbitals to
replace the planetary motion of the electron
Orbitals
 Impossible to determine the location of any single electron
 Orbitals are the regions of space in which electrons can most
probably be found
 Four types of orbitals




s – spherically shaped
p – dumbbell shaped
d – cloverleaf shaped
f – shape has not been determined
 Each additional energy level incorporates one additional orbital type
 Each type of orbital can only hold a specific number of electrons
Orbital Types
Orbital
Type
General
Shape
Orbital
Sublevels
1
# of
electrons
per
sublevel
2
Total # of
electrons
per orbital
type
2
s
Spherical
p
Dumbbell
3
2
6
d
Clover leaf
5
2
10
f
unknown
7
2
14
Electron Configuration
Energy
Level
Orbital
Type
Orbital
Sublevel
1
s
1
s
p
s
p
d
s
p
d
f
1
3
1
3
5
1
3
5
7
2
3
4
# of
# of
# of
orbitals
electrons electrons
per energy per orbital per energy
level (n2)
type
level (2n2)
1
4
9
16
2
2
6
2
6
10
2
6
10
14
2
8
18
32
Electron Configuration
Notation
• Find the element on the periodic table
• Follow through each element block in order by stating the
energy level, the orbital type, and the number of electrons per
orbital type until you arrive at the element.
1s
2s
3s
4s
5s
6s
7s
4f
5f
3d
4d
5d
6d
2p
3p
4p
5p
6p
7p
Samples of e- Configuration
• Element Electron Configuration
•
•
•
•
•
•
•
•
H
He
Li
C
K
V
Br
Pb
1s1
1s2
1s2 2s1
1s2 2s2 2p2
1s2 2s2 2p6 3s2 3p6 4s1
1s2 2s2 2p6 3s2 3p6 4s2 3d3
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 (Note the overlap)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2
Noble Gas Electron Configuration Notation
 Find element on the Periodic Table of Elements
• Example: Pb for Lead
 Move backward to the Noble Gas immediately preceding the
element
 Example: Xenon
 Write symbol of the Nobel Gas in brackets
 Example: [Xe]
 Continue writing Electron Configuration Notation from the
Noble Gas
 Example: [Xe] 6s2 4f14 5d10 6p2
Valence Electrons
• The electrons in the highest (outermost) s and p orbitals of an
atom
• The electrons available to be transferred or shared to create
chemical bonds to form compounds
• Often found in incompletely filled energy levels
Valence Electrons
 Shortcut to finding valence electrons for main group
elements








Family 1A (1)
Family 2A (2)
Family 3A (13)
Family 4A (14)
Family 5A (15)
Family 6A (16)
Family 7A (17)
Family 8A (18)
1 valence electron
2 valence electrons
3 valence electrons
4 valence electrons
5 valence electrons
6 valence electrons
7 valence electrons
8 valence electrons
 Family 3-12 have multiple possibilities and shortcuts do not
work
Electron Dot Notation
 Electron configuration notation using only the valence electrons of an
atom.
 The valence electrons are indicated by dots placed around the element’s
symbol.
 Used to represent up to eight valence electrons for an atom. One dot is
placed on each side before a second dot is placed on any side.
Valance Electrons:
Sodium
1
Electron Dot Notation:
•
Na
Oxidation Numbers:
+1
Magnesium
2
Chlorine
7
Neon
8
•
••
••
Mg
•
: Cl :
: Ne :
•
••
+2
-1
0
Part II Key Terms
• Alpha particle: A helium nucleus emitted by some radioactive
substances
• Beta particle: An energetic electron or positron produced as the result
of a nuclear reaction or nuclear decay
• Beta radiation: Radioactive decay in which an electron is emitted
• Electron Configuration Notation -Consists of an element’s symbol,
representing the atomic nucleus and inner-level electrons, that is
surrounded by dots, representing the atom’s valence electrons.
• Emission spectrum: The range of all possible wave frequencies of
electromagnetic radiation, waves created by the systematic interactions
of oscillating electric and magnetic fields
• Energy Levels - A certain volume of space around the nucleus in which
an electron is likely to be found. Energy levels start at level 1 and go to
infinity.
• Excited state: The state of an atom when one of its electrons is in a
higher energy orbital than the ground state.
Part II Key Terms (cont.)
• Gamma radiation: Electromagnetic radiation emitted during
radioactive decay and having an extremely short wavelength
• Ground state: The lowest energy state of an atom or other particle
• Nuclear fission: Splitting of the nucleus into smaller nuclei
• Nuclear fusion: Combining nuclei of light elements into a larger
nucleus
• Nucleon: a constituent (proton or neutron) of an atomic nucleus
• Planck’s constant: As frequency increases, the energy of the wave
increases
• Radioactive decay: Spontaneous release of radiation to produce a
more stable nucleus
• Radioactive isotope: An isotope (an atomic form of a chemical
element) that is unstable; the nucleus decays spontaneously, giving
off detectable particles and energy
Electromagnetic (EM) Spectrum
• The EM Spectrum is the range of all possible wave frequencies
of electromagnetic radiation, waves created by the systematic
interactions of oscillating electric and magnetic fields
• The general term for all electromagnetic radiation is light
• The range of the EM Spectrum is from very low frequency
known as radio waves to very high frequency known as
gamma radiation
• The visible spectrum of light is in the center portion of this EM
Spectrum
• All EM Spectrum travels at the same speed in a vacuum – this
speed is known as the speed of light, 3.00 x 108 m/s
EM Spectrum
Image used courtesy of http://9-4fordham.wikispaces.com/Electro+Magnetic+Spectrum+and+light
Speed of Light and Frequency
• Since the speed of all EM radiation is the same, there is a clear
mathematical relationship between the frequency of the light
and its wavelength
• All waves travel at a speed that is equal to the product of its
frequency (the reciprocal of time) and its wavelength (distance)
c=fλ
• The speed of EM radiation is fixed at 3.00 x 108 m/s
• Therefore:
3.00 x 108 m/s = f λ
Speed of light = frequency x wavelength
• As frequency increases, wavelength decreases. As wavelength
increases, frequency decreases
• Example: If frequency doubles, wavelength is cut in half
As f ↑, λ↓: Calculations
• If the wavelength of a radio wave is 15 meter, what is its
frequency?
3.00 x 108 m/s = f (10 m)
(3.00 x 108 m/s) / 15 m = f
2.0 x107 s-1 = f
Frequency = 2.0 x107 Hertz
• If the frequency of gamma radiation is 6.25 x 1022 Hertz,
what is its wavelength?
3.00 x 108 m/s = (6.25 x 1022 s-1) λ
(3.00 x 108 m/s) / (6.25 x 1022 s-1) = λ
4.80 x10-15 m = f
Wavelength = 4.80 x10-15 m
Planck’s Law
• Max Planck determined in 1900 there was a mathematical
relationship between the energy of EM radiation and the
frequency of that radiation:
As frequency increases, the energy of the wave increases
E=hf
Energy = Planck’s constant x frequency
E = (6.63 x 10-34 Joule seconds) f
Planck’s Law Calculations
• Example: If the wavelength of green light is 5.21 x 10-7 meters,
what is the energy of this light?
3.00 x 108 m/s = f (5.21 x 10-7 m)
(3.00 x 108 m/s) / 5.21 x 10-7 m = f
5.76 x1014 s-1 = f
Frequency = 5.76 x1014 Hertz
E = (6.63 x 10-34 Joule seconds) (5.76 x1014 s-1)
E = 3.82 x10-19 Joules
Implication of Planck’s Law
• In order to move an electron to a higher energy level, excite
an electron, energy must be absorbed to move the electron
• Since electrons exist in fixed energy levels with a specific
amount of energy, the amount of energy needed is a finite
amount equal to the difference in the energy associated with
the ground state of the electron and the energy associated
with the level to which the electron is excited
• If the energy related to the excited electron is removed, the
electron will return to its ground state and the energy
released is equal to the energy absorbed to excite it
• The energy released is released as light
• The overall result is that every element has a unique spectra
of light associated with it and the spectra can be used to
identify the element
Nuclear Reactions
 All nuclear reactions are based on Einstein’s Theory of
Relativity
 At speeds approaching the speed of light, energy and mass
are interchangeable
 E = mc2
Energy = mass x (speed of light)2
 Mass can be converted to energy and vice versa
Mass Defect
 There is a difference between the mass of an atom and the
various particles that make up the atom
 This difference is called the mass defect of the atom
 This mass defect is the binding energy of the atom
 In nuclear reactions, the binding energy is released as energy
(heat, light, or gamma radiation) and/or particles with
measureable mass
Types of Nuclear Reactions
• Fission – Splitting of the nucleus into smaller nuclei
• Fusion – Combining nuclei of light elements into a larger
nucleus
• Radioactive Decay – Spontaneous release of radiation to
produce a more stable nucleus
Fission
 Nucleus splits into smaller nuclei when struck by a neutron of
sufficient energy
 Tremendous release of energy
 When controlled can produce huge amounts of power in nuclear
reactors
 Naturally occurs in uranium and other ores in spontaneous fission
 Clean source of energy with no carbon footprint
 Produces radioactive nuclear waste with long term environmental
and health considerations
Fission Process
Fission and Nuclear Reactors
Fusion
 Lighter nuclei (such as hydrogen) combined to form heavier
nuclei
 Tremendous release of energy

2H
+
Deuterium

Tritium
3H
4He
+
1n
+ energy
Helium
(occurs naturally in water)
 Powers the sun and stars
 No practical application to produce usable energy at this time
Fusion Process
Radioactive Decay
• Spontaneous release of radiation by unstable nuclei in order to
increase stability
• Radiation can be either energy alone (gamma) or energy
accompanied by release of a particle (all of the other forms of
decay)
Forms of Radioactive Decay
 Alpha decay – release of alpha particle and energy
 Beta decay – release of beta particle and energy
 Gamma Emission – release of electromagnetic
radiation (energy)
 Positron Emission – release of a positron and energy
 Electron Capture – absorption of and electron and release of energy
 Neutron Emission – release of a free neutron and energy
Alpha Decay
 Typically found in heavier nuclei and the means to achieve stability
is to reduce mass
 Nuclei shed mass in the form of a helium nucleus to become more
stable
 Helium nucleus that is released is ionized and called and Alpha
Particle
Alpha Decay (cont.)
 Alpha Particle is positively charged (no electrons present)
 Alpha Particles are very massive, but travel slower (low penetrating
power)
 Can cause significant tissue damage if not shielded
 Shielding can be accomplished with clothing or paper
Alpha Decay Process
Beta Decay
 Common in nuclei of any size where instability is caused by the number
of neutrons
 Neutron decays into a proton and an electron
 Proton remains in the nucleus
 The electron leaves the atom and is called a Beta Particle
Beta Decay (cont.)
 Beta Particle is negatively charged
 Mass of the nucleus is unchanged
 Beta particles have very low mass but are travelling at very high
speed
 Beta particles can penetrate through the skin and cause deep tissue
damage
Beta Decay Process
Gamma Radiation
 Nucleus becomes more stable through the release of
electromagnetic energy
 No change in mass
 No change in the element
 The Gamma radiation can be reduced by shielding, but Gamma
radiation cannot be stopped
 Usually found with another type of decay, but not always
Radioactivity Decay Comparison
Radioactive
Decay Type
Mass
Charge
Penetrating
Power
Transmutation
Alpha
4 amu
Positive
Low
New Element
Formed
Beta
0 amu
Negative
High
New Element
Formed
Gamma
None
(no particle)
None
Extremely High
No
Nuclear Reaction
Mass Conservation
• All nuclear reactions must conserve the overall mass of the
particles involved in the reaction
• Two properties must be the same on both sides of a nuclear
equation
• Total Mass Number – the sum of the mass numbers of all
particles must be the same on both sides of the reaction
• Total Atomic Number – the sum of the atomic numbers of all
particles must be the same on both sides of the reaction