1. Unit 2 – Atomic Structure & Nuclear Chemistry
Part I – Atomic Theory, Subatomic Particles, and Average Atomic Mass

2. 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

3. 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

4. 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

5. Atomic Theory – 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.

6. 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)

7. 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

8. Thomson’s Experiment
Image used courtesy of

9. Thomson’s Model of the Atom
Plum Pudding Model
Image used courtesy of

10. 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

11. Rutherford’s Experiment

12. Rutherford’s Model of the Atom
Nuclear Atomic Model
Image used courtesy of

13. 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

14. Bohr’s Model of the Atom
Electron Shell Model
Image used courtesy of

15. 2 Regions of the Atom Nucleus Electron Cloud
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.

16. Subatomic Particles Electrons Protons Neutrons 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

17. Properties of the Atom Mass Charge Atomic Number
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

18. Example assuming neutral atom of Fluorine
Symbology
Example assuming neutral atom of Fluorine
Atomic number: 9
Mass Number: 19
Protons: 9
Neutrons: 10
(mass number – atomic number)
Electrons: 9 F 19 9

19. Isotopes
Atoms of the same element with different mass due to different number of neutrons

20. Isotope Atomic Mass (amu)
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
% mass oxygen-16: ( ) (.99759) =
% mass oxygen-17: ( ) (.00037) =
% mass oxygen-18: ( ) (.00204) =
Average Atomic Mass of Oxygen =
Isotope
Isotope Atomic Mass (amu)
Population (%)
Oxygen-16
Oxygen-17
0. 037
Oxygen-18

21. 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)

22. Number of electrons (2n2)
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
Number of electrons (2n2) 1
2(12) = 2 2
2(22)= 8 3
2(32)= 18 4
2(42)= 32 n
2n2

23. 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

24. 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

25. Total # of electrons per orbital type
Orbital Types
Orbital Type
General Shape
Orbital
Sublevels
# of electrons
per sublevel
Total # of electrons per orbital type s
Spherical 1 2 p
Dumbbell 3 6 d
Clover leaf 5 10 f
unknown 7 14

26. Electron Configuration
Energy Level
Orbital Type
Orbital
Sublevel
# of orbitals per energy level (n2)
# of electrons per orbital type
# of electrons per energy level (2n2) 1 s 2 p 3 4 6 8 d 5 9 10 18 f 7 16 14 32

27. 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 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

28. Samples of e- Configuration
Element Electron Configuration
H 1s1
He 1s2
Li 1s2 2s1
C 1s2 2s2 2p2
K s2 2s2 2p6 3s2 3p6 4s1
V 1s2 2s2 2p6 3s2 3p6 4s2 3d3
Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 (Note the overlap)
Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2

29. 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

30. 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

31. Valence Electrons
Shortcut to finding valence electrons for main group elements
Family 1A (1) valence electron
Family 2A (2) 2 valence electrons
Family 3A (13) 3 valence electrons
Family 4A (14) 4 valence electrons
Family 5A (15) 5 valence electrons
Family 6A (16) 6 valence electrons
Family 7A (17) 7 valence electrons
Family 8A (18) 8 valence electrons
Family 3-12 have multiple possibilities and shortcuts do not work

32. 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 Magnesium Chlorine Neon
Electron Dot Notation:
  • • •• ••
Na Mg : Cl : : Ne :
• • ••
Oxidation Numbers:

33. 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.

34. 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

35. 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

36. EM Spectrum
Image used courtesy of

37. 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 = f λ
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

38. 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

39. As frequency increases, the energy of the wave increases
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 = h f
Energy = Planck’s constant x frequency
E = (6.63 x Joule seconds) f

40. 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 Joule seconds) (5.76 x1014 s-1)
E = 3.82 x10-19 Joules

41. 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

42. 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

43. 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

44. 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

45. 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

46. Fission Process

47. Fission and Nuclear Reactors

48. Fusion
Lighter nuclei (such as hydrogen) combined to form heavier nuclei
Tremendous release of energy
2H H  He n + energy
Deuterium Tritium Helium
(occurs naturally in water)
Powers the sun and stars
No practical application to produce usable energy at this time

49. Fusion Process

50. 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)

51. 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

52. 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

53. 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

54. Alpha Decay Process

55. 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

56. 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

57. Beta Decay Process

58. 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

59. Radioactivity Decay Comparison
Radioactive Decay Type
Mass
Charge
Penetrating Power
Transmutation
Alpha
4 amu
Positive
Low
New Element Formed
Beta
0 amu
Negative
High
Gamma
None
(no particle)
Extremely High No

60. 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