# III

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```					III. ATOMIC THEORY
A. Basic Electrostatics

1. Interaction of charged particles

2. Force of interaction

kQ1Q2                  k - a constant
________
F=                           Q1 and Q2 - the charges on the bodies
2
r                   r - the distance between bodies

Example(1): different charges, same distance

Case A: +1, +1, 7 cm          Case B: +2, +2, 7 cm

Example(2): same charges, difference distance

Case C: +4, -6, 5 cm          Case D: +4, -6 15 cm
21 ATOMIC THEORY / Properties of Light

B. Properties of Light

1. Wave properties

a) Light is a propagating electromagnetic wave.

In general called electromagnetic radiation, EMR: includes visible, ultraviolet, X -rays etc.

b) Wavelength ((lambda)): the length of a full wave.
Units: meters, centimeters, etc.

c) Frequency ((nu)): number of oscillations/sec, waves/sec.
1
Units: 1/s, s-

d) Speed = wavelength  frequency

The speed of light, c, is constant: 3.00  10 m/s
8

c=  
22 ATOMIC THEORY / Properties of Light

2. Spectrum
10            8                   7     7             4                  2   0
m 10-               10-            4x10- 8x10-      10-          10-         10
8               -6              -5    -5        -2           0           2
cm 10-              10             4x10 8x10        10           10          10
| X-ray          | Ultraviolet | VIS | Infrared | Microwaves | Radio & TV |

inc 

inc 

3. Particle property

Einstein is credited with firmly establishing the particle property of light.

4. Energy of light
34
E = h              h - Planck’s constant, 6.63 x 10 -        Js

Planck is considered to be the father of quantum theory.

-5
Example(1): Find energy of a photon of light with a wavelength of 5.0 x 10        cm.
23 ATOMIC THEORY / Bohr Model

C. Bohr Model and the Periodic Table

1. Postulates    (These are not the original Bohr postulates. Bohr assumed that the angular momentum was quantized.)

a) The electron(s) of an atom circulate the nucleus in orbits with specific fixed

b) The electron(s) has energy.

Etotal =

c) The further the electron(s) is from the nucleus the higher is its energy.

d) Since the electron(s) can be at only certain distance from the nucleus, and since
the distance from the nucleus determines the energy, the electron(s) can have only
certain allowed values of energy.
i.e. the energy of an electron(s) is quantized.

e) Each orbit is associated with a particular amount of energy, so we will refer to them
as energy levels.

f) Each energy level is designated with a number (n), called the principal
quantum number.

2. Origin of light
24 ATOMIC THEORY / Bohr Model

3. The hydrogen spectrum

E
n
e
r
g
y
25 ATOMIC THEORY / Bohr Model

4. Electron configuration by the Bohr Model
2
a) max # e- = 2(n )

Example(1): Calculate the maximum number of electrons that energy levels 1 through 4 can hold.

Example(2): Give the ground state Bohr electron configuration for 1H, 2He, 5B, and 16S

5. Periodic table by the Bohr Model

a) Construction of periodic table

On the blank periodic table, on the next page, fill in the first 18 elements along
with their electron configuration according to Bohr.

b) Valence electron: The electrons in the highest occupied principal energy level
(i.e. largest n value) of an atom.

c) Group number: gives the number of valence electrons in a main group
element.

d) Period number: tells which principal energy level (n value) contains the
valence electrons.

Example(3): How many valence electrons are in 33As?

Example(4): In which principal energy level are the valence electrons of 33As located?
26 ATOMIC THEORY / Bohr Model
27 ATOMIC THEORY / Atomic Orbitals

D. Atomic Orbitals

1. What is wrong with the Bohr Model?

a) It does not predict the correct electron configurations for atoms past   18Ar.

b) Electrons have a dual nature (like light): particle/wave
Dual nature first proposed by DeBroglie

c) Heisenburg uncertainty principle: it is impossible to know both the position
and energy (actually momentum) of an
electron at the same time.

2. Orbitals and probability

a) Orbital: the region in space where there is a high probability of finding the
electron.

b) Types, or shapes, of orbitals
28 ATOMIC THEORY / Atomic Orbitals

c) Orbital orientations

Type     # Orientations    Designations

d) Any single orbital has a maximum capacity of two e- ‘s

3. Orbitals and energy

a) Orbitals: are sub-energy levels of the principal levels predicted by Bohr.
2
b) # orbitals per principal level = n

e.g. the 3 energy level hold 18 e- ‘s (2  3 ). To hold 18 e- ‘s you need 9
rd                                2

orbitals since each orbital can only hold two e- ‘s.
29 ATOMIC THEORY / Atomic Orbitals

c) Energy level diagram

E
n
e
r
g
y
30 ATOMIC THEORY / Atomic Orbitals

E
n
e
r
g
y
31 ATOMIC THEORY / Atomic Orbitals

4. Electron configuration using the energy diagram

Example(1): Give the electron configuration of 1H.

Example(2): Give the electron configuration of 2He.

a) Pauli exclusion principle: no two electrons in an atom can be exactly alike.

There are 4 distinguishing factors for an electron:
st
1 principal energy level

nd
2 orbital type

rd
3 orbital orientation

th
4 spin

b) Hund's rule: Given a chance electrons will remain unpaired.

Example(3): Give the electron configuration of 6C.

Example(4): Give the electron configuration of 8O.

Example(5): Give the electron configuration of 19K.
32 ATOMIC THEORY / Electron Configuration

E. Electron Configuration and Periodic Table

1. Relationship of orbital theory to the structure of the periodic table

On the blank periodic table, write in the last occupied orbital for each element and the number of
electrons in the orbital level. (e.g. the complete configuration of 6C is 1s22s22p2. Under 6C write ... 2p2 .
(The three dots means all lower level orbitals are completely filled.) You will soon see the relation
between the orbitals and the structure of the table.
33 ATOMIC THEORY / Electron Configuration

2. Electron configuration using the periodic chart

s

1                                                                             p

2

3                                d

4

5

6

7

Example(6): Write the complete electron configuration of   11Na,   using the periodic table to determine
the order of orbital filling.

Example(7): Write the complete electron configuration of   33As,   using the periodic table to determine
the order of orbital filling.

Be able to do complete electron configurations using the periodic table as your guide up to
56Ba.

3) Relative energy of orbitals using the periodic table

Example(8): Which orbital has the higher energy, 3s or 3p?

Example(9): Which orbital has the highest energy, 3s, 3p, 3d, or 4s?

Be able to do determine the relative order of energy for the orbitals up to 6s, using the
34 ATOMIC THEORY / Electron Configuration

3. Valence configuration

Example(10): Determine the valance configuration of 12Mg.

Example(11): Determine the valance configuration of   32Ge.

Example(12): Determine the valance configuration of   83Bi.

Be able to do valence configurations for all atoms in the main groups, using the periodic

4. Number of unpaired

Example(13): Determine the number of unpaired electrons in    27Co.

Example(14): Determine the number of unpaired electrons in    44Ru.

Example(15): Determine the number of unpaired electrons in    50Sn.

Be able to find the number of unpaired electrons in all main group atoms and
st       nd
the 1 and 2 row of transition element.
.
35 ATOMIC THEORY / Periodic Properties

F. Periodic Properties

1. Atomic size

The size of an atom (and the other periodic properties we will discuss here)
depends on the amount of force the nucleus exerts on the valence level of the
atom.

a) Across a period: Force increases due to the increase in the nuclear charge 
size decreases.

3Li          4Be   5B        6C       7N        8O        9F       10Ne

NOTE: The Bohr Model is NOT correct, but it is easy to draw and visualize. It does give
the same qualitative results for the atomic size as the atomic orbital theory.

b) Down a group: Force decrease due to increased shielding  size increases.

1H

3Li

11Na

19K
36 ATOMIC THEORY / Periodic Properties

2. Ionization Energy (IE): The amount of energy required to remove an electron
form an atom. (Neutral gaseous atom in its ground state.)

A                              H = IE
+1
A                    + e-

a) Across a period: Force increases due to the increase in the nuclear charge 
the IE increases.

3Li              4Be        5B        6C   7N   8O       9F        10Ne

b) Down a group: Force decrease due to increased shielding the IE decreases.

1H

3Li

11Na

19K

3. Electron affinity (EA): The energy associated with the addition of an electron to
an atom. (Neutral gaseous atom in its ground state.)

A + e-  A-                  H = EA
1

When an electron is added to a nonmetal, except group VIII, energy is
released.
37 ATOMIC THEORY / Periodic Properties

a) Across a period: Force increases due to the increase in the nuclear charge 
the EA increases (becomes more exothermic).

b) Down a group: Force decrease due to increased shielding the EA decreases
(becomes less exothermic).

4. Electronegativity (EN): A measure of an atoms ability to pull electrons to itself.
(Actually a pair of electrons in a chemical bond.)

5. SUMMARY of Periodic Properties

FORCE

SIZE

IE

EA

EN

FORCE

SIZE

IE, EA, EN

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