Brad Collins
Electron Arrangement - Part 2
Chapter 9
Some images Copyright © The McGraw-Hill Companies, Inc.
Review Energy Levels
Multi-electron
n=4
4p
4p
3p
3s
n=2
n=1
3p
4d
4d
4d
4d
3d
3d
3d
3d
3d
4p
4s
n=3
4d
3p
n=3, l = 2
n=3, l = 1
n=3, l = 0
2p
2p
2p
2s
n=2, l = 0
1s
n=1, l = 0
n=2, l = 1
9.1
“Fill up” electrons in lowest energy orbitals (Aufbau principle)
B
electrons
O
C
N
F
Ne
95
6
7
810
electrons
electrons
222s
22s
222p
22p
51
24 6
3
O
Ne
C
N
B
F 1s
1s
1s
2s
2p
22s
22s
12
Be 1s
1s
Li
Be34electrons
electrons
Li
H
He12electron
electrons
He
H 1s
1s12
9.1
Electron ‘Filling’ Order
•
Electrons fill from lowest to highest energy
•
Filling order by shell and subshell:
•
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f,
5d, 6p, 7s, 5f, 6d, 7p
•
No more than 2 electrons can occupy an orbital (Pauli
exclusion principal)
•
Within a subshell, one electron fills each orbital before
the electrons pair up (Hund’s rule).
9.1
Order of orbitals (filling) in multi-electron atom
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
9.1
Electron Configuration
Electron configuration is how the electrons are
distributed among the various atomic orbitals in an
atom.
number of electrons
in the orbital or subshell
1
1s
principal quantum
number n
angular momentum
quantum number l
Electron configuration of carbon: 1s2, 2s2 2p2
9.1
9.1
General Rules for Assigning
Electrons to Atomic Orbitals
1. Each shell or principal level of quantum number n contains
n subshells. If n = 2, there are two subshells (two values
of l ) with angular momentum quantum numbers 0 and 1.
2. Each subshell of quantum number l contains 2l + 1
orbitals. For example, if l = 1, there are three p-orbitals.
3. No more than two electrons can be placed in each orbital.
4. The maximum number of electrons that an atom can have
in a principal level n is equal to 2n2.
9.1
What is the electron configuration of Mg?
Mg 12 electrons
1s < 2s < 2p < 3s < 3p < 4s
1s2 2s22p6 3s2
2 + 2 + 6 + 2 = 12 electrons
Abbreviated as [Ne]3s2
[Ne] =1s2 2s22p6
What are the possible quantum numbers for the
last (outermost) electron in Cl?
Cl 17 electrons
1s < 2s < 2p < 3s < 3p < 4s
1s2 2s22p6 3s23p5
2 + 2 + 6 + 2 + 5 = 17 electrons
Last electron added to 3p orbital
n=3
l=1
ml = -1, 0, or +1
ms = ½ or -½
9.1
Valence Electrons
•
Valence electrons are electrons used by atoms to make chemical bonds.
•
For a particular atom, the valence electrons are electrons with the highest
n-value
•
H has 1 electron in shell n = 1, so one valence electron
•
•
Li has 1 electron in shell n = 2, so one valence electron
•
•
1s1!
1s2 2s1
Cl has 7 electrons in shell n = 3, so seven valence electrons
•
1s2 2s22p6 3s2 3p5
9.1
Periodic Table and Electron
Configuration
•
Groups (columns) of representative elements in the periodic table have the
same number of valence electrons
•
Transition metals are predicted to have 2 valence electrons, but often do not.
•
Filling d-orbitals
•
Varying valence related to d-electron configurations
•
2
2
6
2
6
2
Iron (Fe) configuration: 1s , 2s 2p , 3s 3p , 4s 3d
•
•
•
2+
Fe
3+
Fe
2
2
6
2
6
0
6
2
2
6
2
6
0
5
6
configuration: 1s , 2s 2p , 3s 3p , 4s 3d
configuration: 1s , 2s 2p , 3s 3p , 4s 3d
Note: The 5 d-electrons in Fe
2+
all have parallel spins (Hund’s rule)
9.2
H
Electron Configurations and the
Periodic Table
Li Be
B C N
He
O F Ne
9.2
Electron Configurations - Alternate
Approaches
Noble (rare) Gas abbreviation
Ne = 1s2 2s2 2p6
Mg = 1s2 2s2 2p6 3s1 or [Ne] 3s1
Orbital Diagram
•
Uses boxes and arrows to represent orbitals and electrons
H
He
1s1
1s2
9.3
Anomalous Electron Configurations
•
Some elements have anomalous electron configurations (do not conform to Aufbau
principle)
•
Transition metals e.g., Cr, Cu
•
•
2
Predicted for Cr = [Ar] 4s 3d
1
•
Reason: Half-full d-orbital more stable, parallel spins (Hund’s rule)
2
9
Predicted for Cu = [Ar] 4s 3d
1
Actual for Cu = [Ar] 4s 3d
•
•
5
Actual for Cr = [Ar] 4s 3d
•
•
4
10
Reason: Full d-orbital more stable
More common for Inner-Transition metals e.g., La
•
Predicted for [Xe] 6s2 4f1 Actual for [Xe] 6s2 5d1
9.3
Transition Metal Cations: Electron Configurations
When a cation is formed from an atom of a transition metal,
electrons are always removed first from the ns orbital and
then from the (n – 1)d orbitals.
Fe:
[Ar]4s23d6
Mn:
[Ar]4s23d5
Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
9.3
Electron Configurations of Ions
•
Ions are created when neutral elements gain or lose electrons
•
•
•
Electrons are gained or lost from the valence (outermost) shell
•
Sodium: [Ne] 3s1, Sodium ion: [Ne], Na+
•
Chlorine: [Ne] 3s23p5, Chlorine ion: [Ar], Cl–
The tendency of ions to attain noble gas electron configurations is called
the Octet Rule!
Some elements cannot achieve a Noble Gas configuration
•
Gallium: [Ar] 4s2 3d10 4s1, Gallium ion: [Ar] 3d10, Ga3+
•
Called a psuedo-noble gas configuration!
•
Consists of a noble gas abbreviation plus d and f subshells
9.4
Practice Problem
How many valence electrons does sulfur have?!
Sulfur is in Group 6, [Ne] 3s2 3p4
Sulfur has 6 valence electrons
How many valence electrons does nitride ion have?!
Nitrogen is in Group 6, [Ne] 3s2 3p4
Nitride ion electron configuration is [Ne] 3s2 3p6
Nitride ion has 8 valence electrons
Electron Configurations of Ions
Na+: [Ne]
Al3+: [Ne]
F–: 1s2 2s22p6 or [Ne]
O2–: 1s2 2s22p6 or [Ne] N3–: 1s2 2s22p6 or [Ne]
Na+, Al3+, F–, O2–, and N3– are all isoelectronic with Ne
What neutral atom is isoelectronic with H–?
H–: 1s2 same electron configuration as He
9.4
Periodic Properties
•
Properties of elements depend on their position on the periodic
table
•
Group 1 metals form 2:1 di-metal oxides, M2O
•
Group 2 metals form 1:1 metal oxides, MO
•
Atomic radius (CH 9.5)
•
Ionization energy (CH 9.6)
•
Electron affinity (CH 9.7)
•
Electronegativity (CH 10.3)
Atomic Radius
•
Radius increases
down a group
•
n increases with
each period
•
Trends: radius
increases down
and to the left
Measured from nucleus to nucleus
9.5
Ionic Radii
•
Cation radius is always smaller than neutral atom
•
Anion radius is always bigger than neutral atom
9.5
Ionic Radii
2+extra
This
is due
Na+ and
Mgto
both proton
have
2+ that pulls
in
the2+ is
10Mg
electrons,
but Mg
electrons
closer
smaller than
Na+to the
nucleus than Na+
9.5
Ionization Energy
•
Ionization energy (IE) is the energy required to remove an
electron from an atom in its gaseous state.
•
Ionization energy tends to increase as more electrons are
removed from an atom (Table).
•
Large jumps in IE mark transitions to stable electron
configurations (Octet rule, Hund’s rule).
9.6
First Ionization Energy
Trends
•
First ionization energy is the energy to remove the first electron from an atom.
•
Tends to increase across a period and up a group
•
Exceptions due to more stable electron configurations, e.g., carbon (half-full psubshell, Hund’s rule)
9.6
General Trends in First Ionization Energies
Increasing First Ionization Energy
(Mostly) Increasing First Ionization Energy
9.6
Electron Affinity
•
Electron affinity (EA) is the enthalpy change that occurs when an
electron is added to an atom in the gaseous state.
•
The greater the EA, the more energy an atom is willing to ‘spend’
to obtain an electron.
•
Fluorine will release 328 kJ/mol to obtain an electron.
•
Electron affinities are expressed as –∆H
•
F(g) + 1 e– —> F–(g)
∆H = –328 kJ/mol, EA = 328 kJ/mol
•
O(g) + 1 e– —> O–(g)
∆H = –141 kJ/mol, EA = 141 kj/mol
9.7
Electron Affinity Trends
•
First EA is the energy to add the first
electron to an atom.
•
First EA ends to increase across a
period and up a group
•
Exceptions due to more stable electron
configurations, e.g., carbon (half-full psubshell, Hund’s rule)
9.7
General Trends in Electron Affinity
Increasing Electron Affinity
(Mostly) Increasing Electron Affinity
9.7