Download Ch8.Periodic properties

Periodic properties
of the elements
Chapter 8
1
Quantum numbers

Principal quantum number, n
Determines size and overall energy of orbital
 Positive integer 1, 2, 3 . . .
 Corresponds to Bohr energy levels

2
Quantum numbers

Angular momentum quantum number, l
Determines shape of orbital
 Positive integer 0, 1, 2 . . . (n–1)
 Corresponds to sublevels

l letter
0
s
1
p
2
d
3
f
3
Quantum numbers

Magnetic quantum number, ml
Determines number of orbitals in a sublevel and
orientation of each orbital in xyz space
 integers –l . . . 0 . . . +l

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
What type of orbital is designated by each set of
quantum numbers?
n = 5, l = 1, ml = 0
 n = 4, l = 2, ml = –2
 n = 2, l = 0, ml = 0


5p
4d
2s
Write a set of quantum numbers for each orbital
4s
 3d
 5p

n = 4, l = 0, ml = 0
n = 3, l = 2, ml = –2, –1, 0, +1, or +2
n = 5, l = 1, ml = –1, 0, or +1
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Electron configurations

Electrons
exist within
orbitals,
given by
three
quantum
numbers n, l,
and ml
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Electron configurations
Configuration shows which
orbitals are occupied
Aufbau principle: e– takes
lowest available energy
 Hund’s rule: if there are 2 or
more orbitals of equal energy
(degenerate orbitals), e– will
occupy all orbitals singly
before pairing

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Electron configurations
Electron has spin, either
“up” or “down”
 Electron spin given by 4th
quantum number, ms
ms   12 or  12


Pauli exclusion principle: no two e– in an atom can
have the same set of 4 quantum numbers ⇒
2 e– per orbital, one up ↑ and one down ↓
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Magnetic properties
Atom or ion with unpaired e– is attracted to a
magnetic field = paramagnetic
 Atom or ion with all e– paired is slightly repelled
by a magnetic field = diamagnetic

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Effective nuclear charge: Zeff
Electron experiences
attraction of nucleus
and repulsion of other
e– in the atom
 Outer e– is partially
shielded from full
charge of nucleus
 Zeff = actual nuclear charge – charge shielded by
other e–

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Effective nuclear charge: Zeff
Core e– effectively shield
outer e– from nuclear charge
 Outer e– do not shield other
outer e– very efficiently
 Thus,

Li outer e– experiences Zeff ≈ 3–2 = +1
 Be outer e– experience Zeff ≈ 4–2 = +2

16
Trends in atomic radius

Atomic radius increases down a group
Same Zeff
 Outer e– in higher principal energy level = larger orbital


Atomic radius decreases across a period
Same principal energy level
 Increasing Zeff pulls in outer e–

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Trends in atomic radius

Transition metal radii stay roughly constant across
a period
Outer e– stay same
 Adding protons to nucleus and electrons to n–1 (core)
orbital, so Zeff stays about constant

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19
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Ions and ionic radii

Elements may lose or gain outer e– to form ions

Metals lose e– → cations
Na 1s2 2s2 2p6 3s1  Na1+ 1s2 2s2 2p6

Nonmetals gain e– → anions
F 1s2 2s2 2p5  F1– 1s2 2s2 2p6

The ions in these examples are isoelectronic
(same e– configuration)
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Transition metal ions

Transition metals lose their ns e– before losing their
(n–1)d e–
V [Ar] 4s2 3d3  V2+ [Ar] 4s0 3d3 = [Ar] 3d3
Zn [Ar] 4s2 3d10  Zn2+ [Ar] 4s0 3d10 = [Ar] 3d10
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Ions and ionic radii

A cation is much smaller than its parent atom
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Ions and ionic radii

An anion is much larger than its parent atom
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Ions and ionic radii

For isoelectronic species, the one with the highest
nuclear charge will have the smallest radius
S2–
Cl1–
K1+
Ca2+
18 electrons
16 protons
18 electrons
17 protons
184 pm
181 pm
18 electrons
19 protons
133 pm
18 electrons
20 protons
99 pm
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Ionization energy

Ionization energy (IE) = energy needed to remove
e– from atom/ion in gaseous state
Mg (g)  Mg1+ (g) + e –
IE1 = 738 kJ/mol
Mg1+ (g)  Mg 2+ (g) + e –
IE 2 = 1450 kJ/mol
Mg 2+ (g)  Mg 3+ (g) + e –
IE 3 = 7730 kJ/mol
IE always positive (endothermic)
 Successive IE values always increase
 IE increases dramatically when begin to remove core e–

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Trends in 1st ionization energy

IE1 decreases down a group
Same Zeff
 Outer e– in higher principal energy level = farther from
nucleus, easier to remove


IE1 generally increases across a period

Increasing Zeff pulls outer e– closer to nucleus, harder to
remove
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Two important exceptions

1st e– in p sublevel
Mg: [Ne] 3s2
Al: [Ne] 3s2 3p1
3s orbital penetrates closer to nucleus
 3p e– somewhat more shielded from
nuclear charge, easier to remove

 e–
begin pairing in p sublevel
P
S
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Electron affinity (EA)

EA = energy change when gaseous atom/ion gains
an e–
Cl (g) + e –  Cl1– (g)

EA = –349 kJ/mol
Trends less regular but generally becomes more
exothermic across a period
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Summary of periodic trends

Across a period
atomic radius decreases
 IE1 (always endothermic) increases
(with two important exceptions)
 EA generally becomes more exothermic


Down a group
atomic radius increases
 IE1 decreases

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