1. Atomic Theory

2. The Evolution of the Atomic Model
Basic building block of matter
Cannot be broken down chemically
A single unit of an element

3. Dalton (1803) Known as the founder of the atomic theory
Dalton’s Postulates:
1. All matter is composed of indivisible particles called atoms
2. all atoms of a given element are identical in mass and properties. Atoms of different elements have different masses and different properties
3. Compounds are formed by a combination of 2 or more atoms
4. Atoms cannot be created, destroyed, or converted into other kinds of atoms during chemical reactions

4. Cannonball Theory/Model
1. Spherical
2. Uniform Density

5. J.J. Thomson (1897)
Used a cathode ray tube with charged particle field (+/-)
Cathode ray deflected by negative electrode toward positive electrode
Discovered subatomic particle called the ELECTRON
Small
Negatively charged

6. J.J Thomson Plum Pudding Theory/Model Positive “pudding”
Negative electrons embedded (just like raisin bread)

7. Rutherford (1909)
Conducted the GOLD FOIL EXPERIMENT where he bombarded a thin piece of gold foil with a positive stream of alpha particles
Expected virtually all alpha particles to pass straight through foil
Most passed through but some were severely deflected
Conclusion led to…

8. Rutherford’s Experiment

9. Rutherford Nuclear Theory/Model 1. The atom is mostly empty space
2. At the center of the atom there is a DENSE, POSITIVE CORE called the NUCLEUS
**Provided no information about electrons other than the fact that they were located outside the nucleus**

10. Neils Bohr (1913) Bohr Model or Planetary Model
1. Electrons travel AROUND the nucleus in well defined paths calls ORBITS (like planets in the solar system)
2. Electrons in different orbits possess different amounts of energy

11. Neils Bohr
3. Absorbing/Gaining a certain amount of energy causes electrons to jump to a higher energy level or an excited state
When excited electrons emit/lose a certain amount of energy which causes electrons to fall back to a lower energy level or the Ground State

12. Wave-Mechanical/Cloud Model
Modern, present day model
Electrons have distinct amounts of energy and move in areas called orbitals
Orbital= an area of high probability for finding an electron (not necessarily a circular path)
Developed after the famous discovery that energy can behave as both waves & particles
Many scientists have contribute to this theory using X-Ray diffraction

14. Vocabulary (of the Periodic Table)
Atomic#= the number of protons in every atom of the element (NEVER CHANGES!)
Element Symbol= the letter(s) used to identify an element
Atomic Mass= average mass of all the isotopes of an element

15. **Nucleons= Protons and Neutrons
(any subatomic particle found within the nucleus)
Subatomic Particle
Charge
Relative Mass
Location
Symbol
How to Calculate
Proton +1
1 amu
Nucleus or
Look at atomic #
Neutron
Mass # - Atomic #
Electron -1
1/1836 amu
Outisde nucleus
P= e
(in neutral atoms)

16. Determining Subatomic Particles (p, n, e)
# of Protons = sum of the protons and neutrons in an atom of an element (Atomic Number)
# of Neutrons = Mass # - Atomic # or Mass # - # Protons
# of Electrons = sum of the protons and neutrons in an atom of an element ( equal to # protons in a neutral atom)

17. Determining Subatomic Particles (p, n, e)
Atomic Number:
1. Look at the element symbol and locate on the periodic table
2. same as the # of protons or the nuclear charge
Mass Number = # Protons + # Neutrons
Nuclear Charge = # of protons or the atomic #
P = nuclear charge

18. Atoms vs. Ions Vocabulary Term Definition Example Neutral Atom
An atom with the same number of protons and electrons
P = e
(no charge indicated)
Ion
Two Types
Anion and Cation
Protons and electrons are not the same

19. Atoms vs. Ions

20. Ions Anion Cation aNion An atom that has GAINED one or more electrons
e > p
NEGATIVE ION
ca+ion
An atom that has LOST one or more electrons
p > e
POSITIVE ION

21. Isotope
Atoms of the same element with different mass numbers; same atomic number, same number of protons, different number of neutrons

22. Isotopes
Example 1: Carbon (12, 13, 14)
p =
e =
n =

23. Isotopes
Example 2: Uranium (238, 240)
p =
e =
n =
U-240

24. Practice
1. Two different isotopes of the same element must contain the same number of
2. Two different isotopes of the same element must contain a different number of *
a. protons
b. neutrons
c. electrons *
a. protons
b. neutrons
c. electrons

25. * Practice 3. Isotopes of a given element have
a. the same mass number and a different atomic number
b. the same atomic number and a different mass number
c. the same atomic number and the same mass number *

26. Calculating Atomic Mass (for any element)
Atomic mass = the weighted average of an element’s naturally occurring isotopes
% abundance of isotope 1 x (mass of isotope 1)
% abundance of isotope 2 x (mass of isotope 2)
+ % abundance of isotope 3 x (mass of isotope 3)
Average Atomic Mass of the Element

27. Calculating Atomic Mass
Example 1: The exact mass of each isotope is given
Chlorine has two naturally occurring isotopes, Cl-35 (isotopic mass amu) and Cl-37 (isotopic mass amu). In the atmosphere, 32.51% of the chlorine is Cl-37, and 67.49% is Cl-35. What is the atomic mass of atmospheric chlorine?
Step 1: Multiply the mass of each separate isotope by its percent abundance
CL-35 = amu x =
(.6749)
CL-37 = amu x =
(.3251)

28. Calculating Atomic Mass
Step 2: Add the products of all the calculated isotopes together from step 1
***This is your average atomic mass***

29. Calculating Atomic Mass
The element Carbon occurs in nature as two isotopes. Calculate the average atomic mass for Carbon based on the information below
C-12 = 98.89% C-13 = 1.11%
**Since the mass numbers were not given for either isotope, use the mass number instead**
C-12 = 12
C-13 = 13

30. Calculating Atomic Mass
Step 1: Multiply the mass by the percent abundance
C-12 = 12 x (.9889) =
C-13 = 13 x (.0111) =
Step 2: Add the products together
0.1443
*These are weighted masses*

31. Calculating Atomic Mass
Practice: The element Boron occurs in nature as two isotopes. Calculate the average atomic mass for Boron, using the information below.
Isotope Mass % abundance
Boron amu %
Boron amu %
Average Mass of Boron =

32. Average Mass= 1.014 Isotope % abundance
Practice: The element Hydrogen occurs in nature as three isotopes. Calculate the average atomic mass of Hydrogen.
Isotope
% abundance
Protium
99.0%
Deuterium
0.6%
Tritium
0.4%
Average Mass= 1.014

33. Mass Number Atomic Mass
The MASS of ONE isotope of a given element
The AVERAGE MASS of ALL isotopes of a given element

34. Electron Configurations
The dashed chain of numbers found in the lower left hand corner of an element box
Tells the number of energy levels as well as the number of electrons in each level
(how the electrons are arranged around the nucleus)

35. Electron Configuration
Is the representation of the arrangement of electrons distributed among the orbitals
Used to describe the orbitals of an atom in the ground state
Can be used to describe ionized atoms (cations and anions)
Many of the chemical and physical properties of elements can be correlated to their electron configuration

36. Electron Configurations
All electron configurations on the Periodic Table are NUETRAL ( p = e)
For IONS, add or subtract electrons from the LAST NUMBER in the electron configuration only

37. Electron Configuration
Orbitals
There are four types (s, p, d, and f)
They have different shapes
Each orbital can hold a maximum of two electrons
P, d and f orbitals have different sublevels

38. Electron Configuration
Is unique to an element’s position on the periodic table
Energy level is determined by the period
Number of electrons is given by the atomic number

40. Electron Configuration
Orbitals on different energy levels are similar to each other but they occupy different areas in space
ex. 1s and 2s orbitals both have s orbital characteristics ( radial nodes, spherical volume, can only hold two electrons) but since they are found in different energy levels they occupy different spaces around the nucleus

41. Electron Configuration
Electrons fill orbitals to minimize energy
Electrons fill the principal energy levels in order of increasing energy
Electrons are getting further from the nucleus
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

43. Electron Configuration
A single orbital can hold a maximum of two electrons
These electrons must have opposing spins
One electron spins up and the other spins down
Ensures they have different quantum numbers
Pauli Exclusion Principal

44. Electron Configuration
S subshell 1 orbital that can hold 2 electrons
P subshell 3 orbitals that can hold up to 6 electrons
D subshell 5 orbitals that can hold up to 10 electrons
F subshell 7 orbitals that can hold up to 14 electrons

46. Principal Energy Level (n)
Electron energy levels consisting of orbitals which designated s, p, d, or f.

47. Electron Configurations
Valence Electrons:
Electrons found in the OUTERMOST shell or orbital
Kernel Electrons:
INNER electrons (all non-valence electrons)

49. Orbital Notation

51. Orbital Notation When filling in the orbital
Electrons fill the lowest vacancy levels first
When there’s more than one subshell at a particular energy level (ex. 3p or 4d) only one electron fills each subshell until each subshell has one electron. Then electrons start pairing in each subshell

52. Hund’s Rule
Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

53. Energy Level Diagram: Oxygen. Oxygen is atomic number 8.
8 protons  How many electrons?
8 electrons
The first electron goes into the 1s orbital, filling the lowest energy level first, and the second one spin pairs with the first one.
Electrons 3 and 4 spin pair in the next lowest vacant orbital — the 2s.
Electron 5 goes into one of the 2p subshells (it doesn’t matter which one — they all have the same energy), and electrons 6 and 7 go into the other two totally vacant 2p orbitals.
The last electron spin pairs with one of the electrons in the 2p subshells (again, it doesn’t matter which one you pair it with). The completed energy level diagram for oxygen appears in the following illustration.

54. Bohr Diagrams A method for showing electron location in an atom/ion
All electrons must be drawn
Look up the electron configuration and the period (row) the element is in on the periodic table
Draw a circle (nucleus) write in the number of protons and neutrons
Draw the shells (energy levels) around the nucleus

55. Bohr Models Add the electrons The first shell only holds 2 electrons
Now add the rest of the electrons
Add one at a time starting on the right side and going counter clockwise

56. Maximum number of electrons
Shell (energy level)
Maximum number of electrons 1 2 8 3 18 4 32 5 50 6 72 7 98
The larger the number the further the electrons are from the nucleus, the larger the size of the orbital and the larger the atom is.
The first shell is the ground state or lowest energy state
N can never be 0 or a negative number because an atom cannot not have zero or a negative amount of energy
When an electron is in an excited state it gains energy
2n2 n=
Formula for maximum # of electrons per quantum number (or energy level)

57. Draw the Bohr Model for Carbon
Bohr Models
Example:
Draw the Bohr Model for Carbon

58. Draw the Bohr Model for Oxygen
Bohr Models
Example
Draw the Bohr Model for Oxygen

59. Draw the Bohr Model for Na+
Bohr Models
Example
Draw the Bohr Model for Na+
E- configuration 2-8

60. Lewis Dot Diagrams Only illustrates valence electrons
Write the element symbol
Look up the electron configuration (use the last number in the configuration- # of valence electrons)
Use either an X or dot to represent the electrons
Place that many electrons around the symbol at (12, 3, 6 and 9)

61. Lewis Dot Diagram
Example: Carbon e- configuration 2-4

62. Practice

63. The number of unpaired electrons is equal to the number of BONDS that an element can form with other elements
When determining the number of bonds an element can form, arrange the valence electrons so that you have the MAXIMUM number UNPAIRED
Example: Carbon
How many bonds can carbon form? 4

64. Ground State vs. Excited State
Ground State electrons are in the lowest energy configuration possible (the configuration found on the periodic table)
Excited State electrons are found in a higher energy configuration (any configuration not listed on the periodic table)

65. Ground State
Excited State

66. EXCITED
EXCITED
GROUND
GROUND
EXCITED
GROUND
EXCITED
EXCITED

67. Excited electrons fall rapidly to a lower energy level
The greater the distance from the nucleus, the greater the energy of the electron
When ground state electrons absorb energy they jump to a higher energy level or an excited state
This is very unstable/temporary condition
Excited electrons fall rapidly to a lower energy level
When excited electrons fall from an excited state to a lower energy level, they release energy in the form of light

68. Ground  Excited Excited  Ground Energy is absorbed
Dark line spectrum is produced
Excited  Ground
Energy is released
Bright line spectrum is produced
Dark Lines
Absorbed
Bright Lines
Emitted

69. Different elements produce different colors of light or spectra
Balmer Series: electrons falling from an excited state down to the second (2nd) energy level give off visible light (Bright Line Spectrum or Visible Light Spectrum)
Different elements produce different colors of light or spectra
These spectra are unique for each element
Spectral lines are used to identify different elements

70. *
Gas A and Gas D