Chapter 7; Electronic Structure
of Atoms
I.
II.
III.
IV.
V.
Electromagnetic Radiation
Flame Test/ Emission Spectra
Quantized Energy Levels
Bohr Model/ Rydberg Equation
Principal Energy Levels, n
a)
b)
First Ionization Energy
2nd , 3rd, 4th, etc Ionization Energy
Chapter 7; Electronic Structure
of Atoms
VI. Sublevels (s, p, d, f)
a)
Photoelectron Spectroscopy
VII. Electron Configuration
VIII. Valence Electrons/ Core
IX. Good/ Bad Point of Atom Model
X. Quantum Theory
a)
b)
Dual Nature of the Electron
Heisenberg Uncertainty Principle
Chapter 7; Electronic
Structure of Atoms
XI. Quantum Numbers (n, l, ml, ms)
XII. Oribtal Diagrams
a)
Paramagnetism and Diamagnetism
Electronic Structure Model
Experimental Evidence What it means
1. Line Spectra
2. Ionization Energies
3. Photoelectron
Spectrum
4. Intensity/detail of Line
Spectra
1. Electrons in quanitized ‘n’
2. # electrons in each ‘n’
3. # electrons in each ‘n’ and
each sublevel
4. Indicates ‘n’ have sublevels
associated with them
Electronic Structure
1
# of
# e- in n
Sublevel (2n2)
1
2
Sublevel # e- in each
sublevel
Names
s
s-2
2
2
8
s,p
s-2, p-6
3
3
18
s,p,d
4
4
32
s,p,d,f
s-2, p-6,
d-10
s-2, p-6,
d-10, f-14
n
Order of orbitals (filling) in multi-electron atom
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
7.7
What is the electron configuration of Mg?
Mg 12 electrons
1s < 2s < 2p < 3s < 3p < 4s
1s22s22p63s2
2 + 2 + 6 + 2 = 12 electrons
Abbreviated as [Ne]3s2
[Ne] 1s22s22p6
What is the electron configuration of Cl?
Cl 17 electrons
1s22s22p63s23p5
1s < 2s < 2p < 3s < 3p < 4s
2 + 2 + 6 + 2 + 5 = 17 electrons
7.7
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
8.2
Na+: [Ne]
Al3+: [Ne]
O2-: 1s22s22p6 or [Ne]
F-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 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
8.2
Electron Configurations of
Transition Metals
• Completely filled or half-completely filled
d-orbitals have a special stability
– Some “irregularities” are seen in the electron
configurations of transition and inner-transition
metals.
Electron Configurations of Cations of Transition Metals
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
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
Order of filling; 3s<3p<4s<3d
But when removing electrons to form + ions for transition metals
Order of removing electrons; 4s<3d<3p<3s
8.2
Electronic Structure
Good Points
Bad Points
• Electrons in Quantized
Energy Levels
• Maximum # electrons
in each n is 2n2
• Sublevels (s,p,d,f) and
# electrons they hold
• Electrons are placed in
orbits about nucleus
• Only explains
emission spectra of H2
• Does not address all
interactions
• Treats electron as
particle
There are less interactions to take into
account in H than other elements
Interactions
1. Attraction between
+ nucleus and negative
electrons
Interactions
1. Attraction between + nucleus
and negative electrons
2. Repulsion between electrons
in same energy level.
3. Shielding effect of filled
principal energy levels.
+1
H
+4
Be
Quantum Theory – Revised
Electronic Structure Model
1. Dual Nature of the Electron
2. Heisenberg Uncertainty Principle
Dual Nature of Electron
Previous Concept;
A Substance is Either Matter or Energy
• Matter; Definite Mass and Position
Made of Particles
• Energy; Massless and Delocalized
Position not Specificed
Wave-like
Dual Nature of Electron
• Electron is both “particle-like” and “wavelike” at the same time.
• Previous model only considered “particlelike” nature of the electron
Heisenberg Uncertainty Principle
• Act of measuring the position and energy of
electron changes the position of electron
– Better one variable is known (energy); the less
well the other variable is known (position)
Orbitals Replace Orbits
• Orbits- Both electron position and energy
known with certainty
• Orbitals – Regions of space where an
electrons of a given energy will most likely
be found
Quantum Theory
Orbitals Replace Orbits
Orbits
Orbitals
Schrodinger Wave Equation (Y)
Describes size/shape/orientation of orbitals
• Wave Equation is based on…
1. Dual Nature of Electron (Electron both
particle and wave-like at the same time.)
2. Heisenberg Uncertainty Principle
(Orbitals describe a region in space
an electron will most likely be.)
7.5
Wave Equation (Y)
• Wave Equation describe the size, shape, and
orientation of the orbital the electron (of a
given energy) is in. There are four variables
in the function
(n, l, ml, ms)
-n; Energy and size of orbital
– l; Shape of orbital
– ml; Orientation of orbital
– ms; Electron Spin
1. Each electron has a unique set of 4
Quantum Numbers
2. Each orbital described by the Quantum
Numbers can hold a maximum of 2
electrons.
Schrodinger Wave Equation;
1st Quantum Number
Y = fn(n, l, ml, ms)
principal quantum number n
n = 1, 2, 3, 4, ….
distance of e- from the nucleus
n=1
n=2
n=3
7.6
Schrodinger Wave Equation
2nd Quantum Number
Y = fn(n, l, ml, ms)
angular momentum quantum number l
for a given value of n, l = 0, 1, 2, 3, … n-1
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
l=0
l=1
l=2
l=3
s orbital
p orbital
d orbital
f orbital
Shape of the “volume” of space that the e- occupies
7.6
Principal Energy
Level, n
Sublevel,
l
Quantum #
Electron
Configuration
1
0
(1,0, , )
1s
2
0
(2,0, , )
2s
1
(2, 1, , )
2p
0
(3,0, , )
3s
1
(3, 1, , )
3p
2
(3,2, , )
3d
0
(4,0, , )
4s
1
(4, 1, , )
4p
2
(4, 2, , )
4d
3
(4, 3, , )
4f
3
4
l = 0 (s orbitals)
l = 1 (p orbitals)
7.6
l = 2 (d orbitals)
f-orbitals
Orbital Shapes
Orbital Type
Shape Name
s
Spherical
p
Dumbbell
d
Complex
f
More complex
Schrodinger Wave Equation
3rd Quantum Number
Y = fn(n, l, ml, ms)
magnetic quantum number ml
for a given value of l
ml = -l, …., 0, …. +l
if l = 1 (p orbital), ml = -1, 0, or 1
if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2
orientation of the orbital in space
7.6
Number of Degenerate Orbitals Needed
for Each Type of Orbital (Sublevel)
Type of Orbital Maximum # of # of Degenerate
electrons in
Orbitals
Orbital
s
2
1
p
6
3
d
10
5
f
14
7
ml = -1
ml = -2
ml = 0
ml = -1
ml = 0
ml = 1
ml = 1
ml = 2
Schrodinger Wave Equation
4th Quantum Number
Y = fn(n, l, ml, ms)
spin quantum number ms
ms = +½ or -½
ms = +½
ms = -½
7.6
Valid Possibilities for Quantum Numbers
Chemistry; The Science in Context; by Thomas R Gilbert, Rein V Kriss, and Geoffrey
Davies, Norton Publisher, 2004, p125
How many 2p orbitals are there in an atom?
n=2
If l = 1, then ml = -1, 0, or +1
2p
3 orbitals
l=1
How many electrons can be placed in the 3d
subshell?
n=3
3d
l=2
If l = 2, then ml = -2, -1, 0, +1, or +2
5 orbitals which can hold a total of 10 e7.6
Three Manners to Convey How
Electrons are Arranged
1. Electron Configuration ; List Orbitals and
Number of Electrons in Each
(1s22s22p63s2…)
2. Quantum Numbers (2,0,0,+1/2)
3. Orbital Diagrams; List Orbitals and show
location of electrons and their spin
1s
2s
2p
Orbital Diagrams
Pauli exclusion principle - no two electrons in an atom
can have the same four quantum numbers.
The most stable arrangement of electrons
in subshells is the one with the greatest
number of parallel spins (Hund’s rule)
or maximum # of unpaired electrons.
Orbital Diagrams
Carbon; 6 electrons
Electron Configuration; 1s22s22p2
Orbital Diagram
1s
2s
2p
7.7
Orbital Diagrams
Oxygen; 8 electrons
Electron Configuration; 1s22s22p4
Orbital Diagram
1s
2s
2p
Paramagnetic
unpaired electrons
2p
Diamagnetic
all electrons paired
2p
7.8