CHE 101 Chapter 11

The Atom
• Has a small dense nucleus which is positively
charged
• contains protons (+1 charge)
• contains neutrons (no charge)
• Remainder of the atom is mostly empty space
• contains electrons (–1 charge)
Chapter 11
Electrons

The Atom Ground State and Excited States

• The nuclear charge (n+) is balanced by the presence of • Ground State
n electrons moving in some way around the nucleus
• lowest energy state
• most stable arrangement of electrons
• Excited States
• higher energy states
• result when energy is added to an atom
• relax back to ground state by releasing a photon
• Where are the electrons?
• How are the electrons arranged?

Figure 11.1

Ground State and Excited States Ground State and Excited States

Figure 11.8 Figure 11.9

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CHE 101 Chapter 11

Orbitals Orbitals
• Electrons are in orbitals • Differ in:
• region of space with a high probability of finding • energy
an electron • size
• orbitals do not have sharp boundaries • shape
• define an orbital’s size as the sphere that contains • Electrons in different orbitals differ in energy
90% of the total electron probability

Figure 11.20

Orbital Energy Orbitals

• Principal Energy Level • Principal Energy Levels
• n n = 7 • have sublevels (subshells)
n = 6
• also called shells n = 5
• shape of orbital
• energy of orbital
Energy

n = 4
• specifies the main energy level
n = 3
for the orbital
n = 2
• 7 total levels
• correspond to the rows in the n = 1
periodic table
• higher number = higher energy

Figure 11.22

Principal Energy Level, n = 1 Principal Energy Level, n = 2

• n=1 • n=2
• one sublevel = s orbital • 2 sublevels
• label with number and letter 2s
2p

Figure 11.23

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CHE 101 Chapter 11

Principal Energy Level, n = 2 Principal Energy Level, n = 2

• 2s • 2p
• contains 3 orbitals
• bigger than 1s
• have node at nucleus
• higher energy than 1s
• node = zero probability of finding an electron
• higher in energy than 2s sublevel

Figure 11.24 Figure 11.25

Principal Energy Level, n = 3 Principal Energy Level, n = 3

• n=3 • 3rd sublevel
• 3 sublevels (s, p, d) • d orbitals
• 1st = 3s (bigger and higher in energy than 2s) • contains 5 orbitals
• 2nd = 3p (bigger and higher in energy than 2p) • have 2 nodes
• higher in energy than 3p sublevel

Figure 11.27 Figure 11.28

Principal Energy Level, n = 4 Principal Energy Level, n = 4

• n=4 • f orbitals
• 4 sublevels (s, p, d, f)
• 1st = 4s (bigger and higher in energy than 3s)
• 2nd = 4p (bigger and higher in energy than 3p)
• 3rd = 4d (bigger and higher in energy than 3d)
• 4th = 4f
• contain 7 orbitals
• higher in energy than 4d sublevel

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CHE 101 Chapter 11

Principal Energy Levels Orbital Energy Diagram

7p _ _ _
• n=5 6d _ _ _ _ _
5f _ _ _ _ _ _ _

• only use 5s, 5p, 5d, 5f sublevels 7s _

6p _ _ _

• n=6
5d _ _ _ _ _
Increasing 4f _ _ _ _ _ _ _
6s _
• only use 6s, 6p, 6d sublevels Energy 5p _ _ _
4d _ _ _ _ _
• n=7 5s _
4p _ _ _
• only use 7s, 7p sublevels 3d _ _ _ _ _
4s _
3p _ _ _
3s _
2p _ _ _
2s _
1s _

Orbital Energies from the Periodic

Orbital Energies
Table
• Use periodic table to help

Figure 11.31

Orbital Energies Trick Filling Orbitals with Electrons

• Starting with n = 1 at • Aufbau Principle
the top, write out the • “building up”
orbital labels for each • fill lower energy orbitals first
principal energy level • Pauli Exclusion Principle
• Draw diagonal arrows • each orbital can hold a
down and to the left, maximum of 2 electrons and
starting at 1s those 2 electrons must have
opposite spins

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CHE 101 Chapter 11

Filling Orbitals with Electrons Orbital Diagrams

• Hund’s Rule • Represent an orbital as a square
• degenerate orbitals (orbitals with the same energy) • called a box diagram
are occupied singularly with electrons of the same
• write orbital label under box
spin before pairing electrons up
• orbital energies increase left to right
• i.e. electrons occupy separate orbitals in the same
subshell with parallel (i.e. same) spin before pairing up • orbitals in same subshell have adjacent boxes
• want to maximize unpaired electrons • Represent electrons as arrows
Example: • the direction of the arrow represents the spin of
the electron

2p

Example Example
• Write the orbital diagram for hydrogen atom. • Write the orbital diagram for carbon atom.

H-atom C-atom
Z = 1, 1 e- Z = 6, 6 e-

1s 1s 2s 2p

Example Electron Configuration

• Write the orbital diagram for silver atom. • Shorthand to represent location of electrons
Ag-atom • n(subshell)(number of electrons)
Z = 47, 47 e- • listed in order of energy
Example:
1s 2s 2p 3s 3p 4s
H-atom

3d 4p 5s

4d

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CHE 101 Chapter 11

Example Examples
• Write the electron configuration for carbon.

C-atom 1s 2s
Z = 6, 6 e- Lithium has 1 unpaired electron

1s22s22p2

1s 2s
Beryllium has no unpaired electrons

Examples Example
• Write the electron configuration for silver atom.
Ag-atom
1s 2s 2p
Z = 47, 47 e-
Nitrogen has 3 unpaired electrons
1s 2s 2p 3s 3p 4s

3d 4p 5s 4d
1s 2s 2p
Oxygen has 2 unpaired electrons 1s22s22p63s23p64s23d104p65s24d9

Condensed Electron Configuration Example

• Also called abbreviated electron configuration • Write the abbreviated electron configuration for
• simplifies electron configuration by abbreviating Ag.
inner electrons (electrons in lower principle energy Z = 47, 47 e-
levels) Complete electron configuration:
• use the noble gas that comes before the element 1s22s22p63s23p64s23d104p65s24d9
to abbreviate the inner electrons
Example: Compare to Kr (36 e-):
He = 1s2 1s22s22p63s23p64s23d104p6
Li = 1s22s1 Use [Kr] to represent the first 36 e-
abbreviate: Li = [He]2s1 Ag = [Kr]5s24d9

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CHE 101 Chapter 11

Example Periodic Table

• Write the abbreviated electron configuration for • Use the periodic table as a guide
Se. Example:
Z = 34, 34 e- What is the electron configuration of phosphorous?
Complete electron configuration:
1s22s22p63s23p64s23d104p4 Look at periodic table:
P is in the 3rd row, p-block
Compare to Ar (18 e-):
the last electron will go into a 3p orbital
1s22s22p63s23p6
Use [Ar] to represent the first 18 e- 1s22s22p63s23p3
Se = [Ar]4s23d104p4 or [Ne]3s23p3

Phosphorous:
1s22s22p63s23p3 or [Ne]3s23p3 Using the Periodic Table
• What is the electron configuration of arsenic?
Look at periodic table:
As is in the 4th row, p-block
the last electron will go into a 4p orbital

1s22s22p63s23p64s23d104p3
or [Ar]4s23d104p3

Using the Periodic Table Summary

• What is the electron configuration of manganese? • In a principal energy level that has d orbitals,
Look at periodic table: the s orbital from the next level fills before the
Mn is in the 4th row, d-block d orbitals in the current level
the last electron will go into a 3d orbital • After lanthanum, which has the electron
configuration [Xe]6s25d1, a group of fourteen
elements (the lanthanides) occurs. This series
1s22s22p63s23p64s23d5 of elements corresponds to the filling of the
or [Ar]4s23d5 seven 4f orbitals

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CHE 101 Chapter 11

Summary Exceptions to the Aufbau Principle

• After actinum, which has the configuration • For some elements, half filled and filled subshells
[Rn]7s26d1,a group of fourteen elements (the are apparently more stable
actinides) occurs. This series corresponds to the • Examples:
• chromium (Cr)
filling of the seven 5f orbitals expected: [Ar]4s23d4
• Except for helium, the group numbers indicate actual: [Ar]4s13d5
the sum of electrons in the ns and np orbitals in
the highest principal energy level that contains • copper (Cu)
electrons (where n is the number that indicates a expected: [Ar]4s23d9
particular principal energy level) actual: [Ar]4s13d10

Exceptions to the Aufbau Principle Exceptions to the Aufbau Principle

Figure 11.30 Figure 11.34

Alkali Metals Classifying Electrons

• Li: [He]2s1 • Core electrons
• Na: [Ne]3s1 • also called inner electrons
• K: [Ar]4s1 • all the inner shell electrons, typically those of the
• Rb: [Kr]5s1 previous noble gas
• Cs: [Xe]6s1 • Valence electrons
• Fr: [Rn]7s1 • electrons in the outermost (highest) principal
energy level of an atom
• predict chemical reactivity
Notice: all have one electron in the outer most shell

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CHE 101 Chapter 11

Valence Electrons Counting Valence Electrons

• Elements in the same group have the same • Main Group Elements
valence electron configuration • number of valence electrons is equal to group number
Group 1 (Alkali Metals) = 1 valence electron
• Elements with the same valence electron
Group 2 (Alkaline Earth Metals) = 2 valence electrons
arrangement show very similar chemical behavior
Group 3 = 3 valence electrons
Example: Group 4 = 4 valence electrons
Li: [He]2s1 Na: [Ne]3s1 Group 5 = 5 valence electrons
K: [Ar]4s1 Rb: [Kr]5s1 Group 6 = 6 valence electrons
Cs: [Xe]6s1 Fr: [Rn]7s1 Group 7 (Halogens) = 7 valence electrons
Group 8 (Nobel Gases) = 8 valence electrons
Alkali Metals all have 1 valence electron (ns1)

Counting Valence Electrons Ions

• Transition Metals • Group 1 metals always form +1 cations
• number of valence electrons is 2 • Group 2 metals always form +2 cations
• found in the ns2 subshell • Group 3 metals always form +3 cations
(n = row number of transition metal)
Example: • Group 7 nonmetals form -1 anions
Iron • Group 6 nonmetals and Te always form -2 anions
[Ar]4s23d6 • Group 5 nonmetals and As always form -3 anions
outer most principal energy level is n = 4
have two n = 4 electrons
WHY?????
have 2 valence electrons

Cations Anions
• Metals lose valence electrons to reach a noble • Nonmetals gain valence electrons to reach a
gas-like configuration noble gas configuration
Examples: Examples:
Na  Na+ + 1e- F + 1e-  F-
11 e- 10 e- 9 e- 10 e-
[Ne]3s1 [Ne] 1s22s22p5 1s22s22p6 = [Ne]

Ga  Ga3+ + 3e- As + 3e-  As3-

31 e- 28 e-
33 e- 36 e-
[Ar]4s23d104p1 [Ar]3d10
[Ar]4s23d104p3 [Ar]4s23d104p6 = [Kr]

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CHE 101 Chapter 11

Isoelectronic Isoelectronic
• Species with the same number of electrons
Example:
Na+, Mg2+, Al3+, O2-, F-, and Ne = 10 electrons

Example Example
• What is the electron configuration for the • What is the electron configuration for the
most stable ion of calcium? most stable ion of iodine?

Ca = [Ar]4s2 I = [Kr]5s24d105p5

Ca2+ = [Ar] I- = [Kr]5s24d105p6 = [Xe]

Electrons
• Be sure to practice some problems!
• Extra Credit Puzzle
• due next Tuesday

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