Chapter 4: Arrangement of
Electrons in Atoms
I. Rutherford model
A. Worked with it in chapter 3.
B. Protons and neutrons were in the nucleus
C. Electrons were around the nucleus.
D. Rutherford did not say anything about how the
electrons were arranged.
II. Light – It’s physics time!
A. Has properties of both waves and particles.
B. P. 92 – Electromagnetic Spectrum
1. Shortest wavelengths have the highest
a. Blue end of the spectrum
b. Has the most energy.
2. Longest wavelengths have the lowest frequency
a. Red end of the spectrum
b. Lowest amount of energy
3. Wavelength - λ – distance between crest to crest
(or trough to trough)
4. Frequency – v -number of waves/time period
5. Energy – E
6. E = hv (h = Plank’s constant)
C. Each light particle carries a quantity of energy
1. Called a photon
2. Has zero mass, but has energy
3. The color of the light tells the quantity of energy
III. Bohr Model
A. Niels Bohr 1913 – Danish
B. Passed current through hydrogen gas and used its line-
emission spectrum results. (p. 95 – Figure 4-5)
1. The electric current “excited” the electrons
2. When the electrons returned to the “ground” state
(the lowest energy levels), energy is given off in the
form of light.
3. The color of the band of light tells us the amount of
energy given off.
4. Bohr found that the spectrum for hydrogen is always
5. Bohr believed that the e- were only in specific orbits.
a. Like a bookshelf – the energy levels can only be
the shelves not the spaces in between the shelves.
b. Ground state – the lowest energy level allowed
for that electron.
c. Excited state
1. Energy has been added (usually by heat or
electricity) the e- moves to a state above the
2. The more energy that is absorbed by the e-,
the higher the energy level the e- will move to.
3. Excited e- are unstable and tend to drop down to
lower energy states. Excess energy is given off in
the form of light.
6. Bohr’s Model
N ) ) ) ) ) ) )
Nucleus 1 2 3 4 5 6 7
lowest energy highest energy
1= ground state
Quantitized energy is designated as n.
a. e- goes from level 2 to 1 some energy is given off.
b. e- goes from level 3 to 1 some other amount of
energy is given off. A different color of light is seen.
c. This results in the line spectrum seen.
d. This model works well for Hydrdogen, but not so
well for other elements.
e. Common model used to depict the atom or atomic
IV. Quantum Mechanical Model (Quantum Model)
A. Modern View – developed in the 1920’s.
B. Significant changes from the Bohr model.
1. e- move at very high speeds in complicated 3-D
patterns – based on calculus
2. Orbital is a 3-D region around the nucleus that indicate
the PROBABLE location of an e-.
a. p. 101 of text. Figure 4-11
b. Since e- are negatively charged and in rapid motion –
they are viewed as clouds of negative charge called
V. Quantum Numbers
A. Describes the e- probable location.
B. First or Principal Quantum Number
1. n – from the Bohr model
2. Indicates the energy level of the e- . – how far the e- is
from the nucleus.
3. e- wants to be at the LOWEST energy level possible.
a. Protons and electrons attract
b. Electrons repel each other.
C. Second Quantum Number or Angular momentum
2. Indicates the type or shape of the orbital.
3. s, p, d, f
p: pair of lobes
d: 4 lobes (usually)
f: 8 lobes (usually)
D. Third or magnetic quantum number
2. Describes the orientation or
direction in space of the orbital.
3. 3-D – x, y and z axis.
4. px, py, pz,
E. Fourth or spin Quantum number
1. Spin of the e- .
2. Clockwise or counterclockwise
3. Arrows used
4. Maximum of 2 electrons can fit into each orbital
Energy level Sublevels
2 2s, 2p
3 3s, 3p, 3d
4 4s, 4p, 4d, 4f
5-7 s, p, d, f
F. Several sublevels are available due to their orientation.
Sublevel # of possible # of spin Total # of
Orbitals directions electrons
s 1 2 1x2=2
p 3 2 3x2=6
d 5 2 5 x 2 = 10
f 7 2 7 x 2 = 14
G. Electron Configuration Notation
a. Find the total number of electrons that the atom
b. The electron always goes to the LOWEST energy
c. Superscripts indicate how many e- are in each
a. Has an atomic number of 6 – therefore 6 e-.
b. 1s2 2s22p2
H. Orbital Filling Order – p. 105 & p. 110
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
Draw parallel lines starting with 2p 3s.
Follow the arrows -
13 e-: 1s2 2s2 2p6 3s2 3p1
50 e- : 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p2
or 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p2
82 e-: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
6s2 4f14 5d10 6p2
or 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14
5s2 5p6 5d10 6s2 6p2
I. Orbital Notation
1. Rules –
a. e- are distributed in the lowest energy sublevels first
then they move to the next higher energy sublevel.
b. If 2 or more e- have energies that place them in the
same sublevel, the lower energy state will be the state in
which the e- are not in the same orbital.
2. Example: C
6 e- 1s2 2s22p2
__ _ ___ Note: This has a
1s 2s 2px 2py 2pz lower energy than ↓.
(Actually C uses something called sp3 – but that is another day...)
3. ___ ___ (incorrect)
Always spread out the e- within the same orbital type.
J. Noble – gas notation
1. Write the e- configuration for the noble gas
preceding the element, then list the remaining
2. Put it into [Noble gas] remaining electron
3. S [ Ne ]3s2 3p4
K. Visualizing the atom
1. The higher the quantum number, the further away
from the nucleus the orbit is, the larger the atom.
2. It works for p, d and f shells also.
L. The P.T. and orbital filling - we will cover more in
• Orbital filling in the P.T.
s d p
n n-1 n
+ 2+ 3+ +/- 3- 2- - 0
VI. Using this P.T. to fill your orbitals
A. Use your finger to fill the orbitals as you did before
n = row number and is the same for s & p orbitals
d orbitals are n-1
f orbitals are n-2
C. Fe - 26 electrons
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Noble gas configuration:
[Ar] 4s2 3d6
D. Pb – Noble gas configuration
[Xe] 6s2 4f14 5d10 6p2