AP Chem: Notes chapters 7, 8, 9
Read/Do Problems for Chapters 7, 8, 9
Chapter 7: Atomic Structure and Periodicity
Electromagnetic Radiation:
Radiant energy that exhibits wavelength-like behavior and travels
through space at the speed of light in a vacuum.
Label the parts of the graph below: wavelength, amplitude, node
Waves have 3 primary characteristics:
1.
Wavelength (λ): distance between two peaks in a wave.
2.
Frequency (ν): number of waves per second that pass a given point in space.
3.
Speed (c): speed of light is 2.9979 x 108 meters/second (m/s)
Wavelength and frequency can be inter-converted:
c = λν
ν = frequency (s-1, 1/s, Hz, cyc/s, or waves/s )
λ = wavelength (m)
c = speed of light (m/s)
1 hertz (Hz) = 1/s
Max Planck (1858-1947)
An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA
Transfer of energy is quantized, and can only occur in discrete units, called quanta.
∆ 𝐸 = ℎ𝜈 =
ℎ𝑐
𝜆
ΔE = change in energy, (J)
h = Planck’s constant, 6.626 x 10-34 J · s
ν = frequency, (s-1)
λ = wavelength, (m)
Photoelectric Effect:
Albert Einstein (1879-1955)
Experiment demonstrates the particle nature of light. Einstein called light particles “photons.”



Classical physics viewed light as a form of energy so it
predicted that a metal would eventually collect enough
energy to eject an electron regardless of the wavelength
of the light.
The emission of electrons from a metal when light shines
on the metal.
For a given metal no electrons were emitted if the
frequency of light was below a “threshold” value
regardless of how long the light was shining on the metal.
Classical theory said that E of ejected electron should increase with increase in light intensity… but
 No e- observed until light of a certain minimum E is used.
 Number of e- ejected depends on light intensity.
Energy has mass:
E = mc2
E = energy
m = mass
c = speed of light

Ephoton =
ℎ𝑐
mphoton =
ℎ
𝜆
𝜆𝑐
Hence the dual nature of light or wave-particle duality.
Line Spectra of excited atoms:
Excited atoms emit light of only certain wavelengths
The wavelengths of emitted light depend on the element.
Observe spectrum tubes with diffraction gratings (spectroscopes)
Visible lines in H atom spectrum are called the BALMER series:
Continuous spectrum: Contains all the wavelengths of light.
Line (discrete) spectrum: Contains only some of the wavelengths of light.
Niels Bohr (1885-1962): Bohr’s greatest contribution to science was in building a simple
model of the atom. It was based on an understanding of the SHARP LINE SPECTRA of
excited atoms.
One view of atomic structure in early 20th century was that an electron (e-) traveled about the
nucleus in an orbit. (Bohr’s Atomic Model)
+
Electron
orbit
1. Any orbit should be possible and so is any energy.
2. But a charged particle moving in an electric field should emit energy.
End result should be destruction!
Bohr said classical view is wrong.
 Need a new theory — now called QUANTUM or WAVE MECHANICS.
 e- can only exist in certain discrete orbits — called stationary states.
 e- is restricted to QUANTIZED energy states.
 Energy of state = - C/n2
 where C is a constant & where n = quantum no. = 1, 2, 3, 4, ....
Bohr Model:
 The e- in a hydrogen atom moves around the nucleus only in certain allowed circular orbits.

E = -2.178 × 10
-18
 z2 
J  2
n 
E = energy of the levels in the H-atom
z = nuclear charge (for H, z = 1)
n = an integer
E becomes more negative as the electron moves closer to the nucleus.
Ground State: The lowest possible energy state for an atom (n = 1).
Shielding Effect: Higher energy electrons are in outer shells, and are easier to remove
because they are more shielded from the positive nucleus by inner shell electrons.
Energy Changes in the Hydrogen Atom
ΔE = Efinal state – Einitial state
𝜆 =
ℎ𝑐
𝛥𝐸
Failure of Bohr Model:
but

The Bohr Model of the atom paved the way for the Quantum Mechanical Theory,

current theory is in no way derived from the Bohr Model of the atom.
The Bohr Model of the Atom was fundamentally incorrect--electrons do not move in circular
orbits about the nucleus.
Um…so what is quantum mechanics?
Louis deBroglie (1892-1987)
de Broglie (1924) proposed that all moving objects have wave properties.
λ=
ℎ
𝑚v
λ = wavelength (m)
h = Planck’s constant, 6.626 x 10-34 J · s = kg m2 s-1
m = mass, (kg)
v = velocity (m/s) – do not confuse with nu (frequency) or c (speed of light)
See Sample Exercise 7.3 on page 298.
Baseball (0.100 kg) at 35
𝑚
𝑠
λ = 1.9 x 10-34 m
e- with velocity = 1.0 x 107
λ = 7.27 x 10-11 m
𝑚
𝑠
Erwin Schrodinger 1887-1961
Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms.
Wave Equation: Based on the wave properties of the atom
Ĥψ = Eψ
ψ = wave function
Ĥ = mathematical operator
E = total energy of the atom
A specific wave function is often called an orbital. Electrons move in orbitals not Bohr
orbits.
Werner Heisenberg 1901-1976
Heisenberg Uncertainty Principle:
 Problem of defining nature of electrons in atoms solved by Werner Heisenberg.
 Cannot simultaneously define the position and momentum (p = m•v) of an electron.
 We define e- energy exactly but accept the limitation that we do not know exact position.
x   mv  
h
4
x = position
mv = momentum (p)
h = Planck’s constant
The more accurately we know a particle’s position, the less accurately we can know its momentum.
Probability Distribution: square of the wave function (ψ2)
 probability of finding an electron at a given position
 Radial probability distribution is the probability distribution in each spherical shell.
Quantum Numbers: describe the location of the esymbol
name
possible values
what it tells
about electron
cloud
n
principal quantum number
1-7
size
(energy level)
l
angular momentum
(sublevel) quantum
number
l =0 to n-1
shape
l = 0 (s sublevel), l = 1 (p sublevel), l = 2 (d sublevel), l = 3 (f sublevel)
ml
magnetic quantum number
ms
spin quantum number
m = l to -l
l
±
1
2
3D orientation in
space
spin states of
electron
Pauli Exclusion Principle:
 In a given atom, no two electrons can have the same set of four quantum numbers
(n, l, ml ,ms).
 Therefore, an orbital can hold only two electrons, and they must have opposite spins.
Shapes of Electron Orbitals
The following are graphical representations of the shapes of the orbitals within atoms. Remember
these shapes do not exist but are mainly three-dimensional probabilities of locating the electron.
Also remember, a maximum of two electrons can occupy any of the following shapes.
How many different kinds of orbitals does each energy level have?



n = the number of the energy level.
n2 = the number of orbitals in an energy level.
2n2 = the number of electrons in an energy level.
s-orbital (spherical)
The following picture shows the relative sizes of 1s, 2s, and 3s orbitals.
p-orbital (dumbbell)
3 orbitals: px, py. pz - The three p orbitals lie 90° apart in space
When l = 1, there is a PLANAR NODE thru the nucleus.
When n = 2, then l = 0 and 1
Therefore, in n = 2 shell there are 2 types of orbitals — 2 subshells
For l = 0 m = 0
this is a s subshell
For l = 1
l
m = -1, 0, +1
l
this is a p subshell with 3 orbitals
d-orbital (double dumbbell and dumbbell/donut)
5 orbitals
When n = 3, what are the values of l?
l = 0, 1, 2
and so there are 3 subshells in the shell.
For l = 0,
m = 0 ---> s subshell with single orbital
l
For l = 1, m = -1, 0, +1 ---> p subshell with 3 orbitals
For l = 2,
l
m = -2, -1, 0, +1, +2---> d subshell with 5 orbitals
l
s orbitals have no planar node (l = 0) and so are spherical.
p orbitals have l = 1, and have 1 planar node, and so are “dumbbell” shaped.
This means d orbitals (with l = 2) have 2 planar nodes
f-orbital (quadruple dumbbell and dumbbell/double donut)
7 orbitals
When n = 4, l = 0, 1, 2, 3 so there are 4 subshells in the shell.
For l = 0, m = 0 ---> s subshell with single orbital
l
For l = 1, m = -1, 0, +1 ---> p subshell with 3 orbitals
l
For l = 2, m = -2, -1, 0, +1, +2 ---> d subshell with 5 orbitals
l
For l = 3, m = -3, -2, -1, 0, +1, +2, +3 ---> f subshell with 7 orbitals
l
Aufbau principle:
As protons are added one by one to the nucleus to build up the elements, electrons are
similarly added to these hydrogen-like orbitals.
electron filling order (aufbau):
Hund’s Rule:
The lowest energy configuration for an atom is the one having the maximum number of
unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same
energy) orbitals.
Valence electrons: electrons in the outermost principle quantum level of an atom.
The inner electron are called core electrons.
atom
Ca
N
Br
valence electrons
2
5
7
Arrangement of Electrons in Atoms:
Shells (energy levels) = n
↓
Sublevels = l
↓
Orbitals = m
l
Assigning electrons to atoms:
Electrons generally assigned to orbitals of successively higher energy.
For H atoms, - C/n2 . E depends only on n.
For many-electron atoms, energy depends on both n and l.
Effective Nuclear Charge, Zeff
Zeff is the nuclear charge experienced by the outermost electrons. Explains why E(2s) < E(2p)
Zeff increases across a period owing to incomplete shielding by inner electrons.
Estimate Zeff by --> [ Z - (no. inner electrons) ]
Charge felt by 2s e- in Li
Be
Z
B
Z
and so on!
Z
eff
eff
eff
=3-2=1
=4-2=2
=5-2=3
Writing Atomic Electron Configurations:
Orbital Notation

Each electron is represented by a half-arrow

When electrons share an orbital, opposite spins are represented by half-arrows going in
opposite directions
Examples: Write the orbital notation for the following atoms:
N (7 e-):
P ( ___ e-):
____
____
____ ____ ____
1s
2s
2p
Electron Configuration Notation
 Write the energy level and orbital (in increasing energy)
 Write the number of electrons in the sublevel with a superscript
Examples: Write the electron configuration notation for the following atoms:
N:
P:
Tc:
Electron Configuration Notation (Noble Gas Shorthand)
 Start with the noble gas “closet without going over” in brackets
 Continue normal electron configuration
Examples: Write the electron configuration notation for the PREVIOUS atoms (use the noble gas
shorthand)
Lewis Dot Diagrams
Valence electrons (e- in the outermost energy level) are the only electrons that are lost or
gained in ordinary chemical reactions
 Lewis Dot Diagrams only include electrons in valence (highest energy level)
 Write the letter and dots around the outside (2 on each side: maximum of 8)
Examples: Draw the Lewis Dot Diagram for:

N
Ca
Cl
43 61
7
2
58
X
Ion configurations: To form cations from elements remove 1 or more e- from subshell of highest n
[or highest (n + l)].
P [Ne] 3s2 3p3 - 3e- ---> P3+ [Ne] 3s2 3p0
For transition metals, remove ns electrons and then (n - 1) electrons.
Fe [Ar] 4s2 3d5
loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6
loses 3 electrons ---> Fe3+ [Ar] 4s0 3d5
Information Contained on the Periodic Table:
1.
Each group member has the same valence electron configuration
(these electrons primarily determine an atom’s chemistry).
2.
The electron configuration of any representative element. (see order)
3.
Certain groups have special names (alkali metals, halogens, etc).
Noble
gases
02_29
Alkaline
1 earth metals
Halogens
1A
1
Alkali metals
H
8A
2
13
14
15
16
17
2A
3A
4A
5A
6A
7A
2
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
13
14
15
16
17
18
Al
Si
P
S
Cl
Ar
11
12
Na
Mg
3
4
5
6
7
8
9
Transition metals
10
11
12
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
La*
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
104
105
106
107
108
109
110
111
87
88
Fr
Ra
89
Ac†
Unq Unp Unh Uns Uno Une Uun Uuu
*Lanthanides
† Actinides
4.
18
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Metals and nonmetals are characterized by their chemical and physical properties
Periodic Trends: overall change in physical property through row or group on the periodic table.
These are general trends (there are exceptions).
Trend #1: Atomic Radius
Trend #2: Ionic Radius
Cations (positive ions) generally are smaller than their respective atoms
Anions (negative ions) generally are larger than their respective atoms
Trend #3 – Ionization Energy (IE)
Trend #4 – Electronegativity (en)
Redox Reactions:
Why do metals lose electrons in their reactions?
Why does Mg form Mg2+ ions and not Mg3+?
Why do nonmetals take on electrons?
Electron Affinity: A few elements GAIN electrons to form anions.
Electron affinity is the energy involved when an anion loses an electron.
A -(g) → A(g) + eE.A. = ΔE
EA = 141 kJ
ΔE is ENDOthermic because O has an affinity for an e-.
EA = 0 kJ
ΔE is zero for N- due to electron-electron repulsions
Trend #5 Electron Affinity
Affinity for electron increases across a period (EA becomes more negative).
Reason:
Affinity decreases down a group (EA becomes less negative).
Reason:
35 0
30 0
15 0
S4
10 0
S3
Period
50
S2
0
S1
1
2
3
4
5
Group
6
7
Elec tron affinity (kJ/m ol)
25 0
20 0