Chapter 8 - Suffolk County Community College

Report
Chemistry: A Molecular Approach, 2nd Ed.
Nivaldo Tro
Chapter 8
Periodic
Properties of
the Elements
Roy Kennedy
Massachusetts Bay Community College
Wellesley Hills, MA
Copyright  2011 Pearson Education, Inc.
Nerve Transmission
• Movement of ions across cell membranes
is the basis for the transmission of nerve
signals
• Na+ and K+ ions are pumped across
membranes in opposite directions through
ion channels
Na+ out and K+ in
• The ion channels can differentiate Na+
from K+ by their difference in size
• Ion size and other properties of atoms are
periodic properties – properties whose
values can be predicted based on the
element’s position on the Periodic Table
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Mendeleev
(1834–1907)
• Order elements by atomic mass
• Saw a repeating pattern of properties
• Periodic Law – when the elements are
arranged in order of increasing atomic mass, certain
sets of properties recur periodically
• Put elements with similar properties in the same
column
• Used pattern to predict properties of undiscovered
elements
• Where atomic mass order did not fit other properties,
he re-ordered by other properties
Te & I
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Periodic Pattern
NM
H2O
a/b
H
1.0
H2
M
Li
6.9
M
Li2O
b
M
LiH
9.0
Be
Na2O M
b
Na
23.0
NaH 24.3
M
K2O
b
K
39.1
KH
BeO NM
a/b
B
BeH2 10.8
MgO
b
M
C
BH3 12.0
M
Mg
B2O3 NM
a
Al
MgH2 27.0
CaO
b
Al2O3 M/NM
a/b
Si
AlH3 28.1
CO2 NM
a
N
N2O5 NM
a
O2
NM
O
F
CH4 14.0
NH3 16.0
H2O 19.0
HF
NM
P4O10 NM
a
SO3 NM
a
Cl2O7
a
SiO2
a
SiH4 31.0
P
PH3 32.1
S
Cl
H2S 35.5
HCl
Ca
40.1
CaH2
M = metal, NM = nonmetal, M/NM = metalloid
a = acidic oxide, b = basic oxide, a/b = amphoteric oxide
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Mendeleev's Predictions
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What vs. Why
• Mendeleev’s Periodic Law allows us to predict
•
•
what the properties of an element will be
based on its position on the table
It doesn’t explain why the pattern exists
Quantum Mechanics is a theory that explains
why the periodic trends in the properties exist
and knowing Why allows us to predict What
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Electron Configurations
• Quantum-mechanical theory describes the
•
•
behavior of electrons in atoms
The electrons in atoms exist in orbitals
A description of the orbitals occupied by
electrons is called an electron configuration
principal energy
level of orbital
occupied by the
electron
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1s1
number of electrons
in the orbital
sublevel of orbital
occupied by the
electron
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How Electrons Occupy Orbitals
• Calculations with Schrödinger’s equation show
•
hydrogen’s one electron occupies the lowest energy
orbital in the atom
Schrödinger’s equation calculations for
multielectron atoms cannot be exactly solved
 due to additional terms added for electron-electron
interactions
• Approximate solutions show the orbitals to be
•
hydrogen-like
Two additional concepts affect multielectron atoms:
electron spin and energy splitting of sublevels
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Electron Spin
• Experiments by Stern and Gerlach showed a
•
•
beam of silver atoms is split in two by a
magnetic field
The experiment reveals that the electrons spin
on their axis
As they spin, they generate a magnetic field
spinning charged particles generate a magnetic field
• If there is an even number of electrons, about
half the atoms will have a net magnetic field
pointing “north” and the other half will have a net
magnetic field pointing “south”
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Electron Spin Experiment
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The Property of Electron Spin
• Spin is a fundamental property of all electrons
• All electrons have the same amount of spin
• The orientation of the electron spin is
quantized, it can only be in one direction or its
opposite
spin up or spin down
• The electron’s spin adds a fourth quantum
number to the description of electrons in an
atom, called the Spin Quantum Number, ms
not in the Schrödinger equation
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Spin Quantum Number, ms, and
Orbital Diagrams
• ms can have values of +½ or −½
• Orbital Diagrams use a square to
•
•
represent each orbital and a halfarrow to represent each electron in
the orbital
By convention, a half-arrow pointing
up is used to represent an electron in
an orbital with spin up
Spins must cancel in an orbital
paired
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Orbital Diagrams
• We often represent an orbital as a square and
the electrons in that orbital as arrows
 the direction of the arrow represents the spin of the
electron
unoccupied
orbital
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orbital with
one electron
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orbital with
two electrons
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Pauli Exclusion Principle
• No two electrons in an atom may have the same set
•
•
of four quantum numbers
Therefore no orbital may have more than two
electrons, and they must have with opposite spins
Knowing the number orbitals in a sublevel allows us
to determine the maximum number of electrons in
the sublevel
 s sublevel has 1 orbital, therefore it can hold 2 electrons
 p sublevel has 3 orbitals, therefore it can hold 6 electrons
 d sublevel has 5 orbitals, therefore it can hold 10 electrons
 f sublevel has 7 orbitals, therefore it can hold 14 electrons
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Allowed Quantum Numbers
Quantum
Number
Principal, n
Values
Number Significance
of Values
1, 2, 3, ...
size and
energy of the
orbital
shape of
Azimuthal, l 0, 1, 2, ..., n n
1
orbital
orientation of
Magnetic, -l,...,0,...+ l
2l + 1
orbital
ml
direction of
Spin, m s
-_ , +_
2
electron sp in
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Quantum Numbers of Helium’s Electrons
• Helium has two electrons
• Both electrons are in the first energy level
• Both electrons are in the s orbital of the first energy
level
• Because they are in the same orbital, they must have
opposite spins
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Sublevel Splitting in Multielectron Atoms
• The sublevels in each principal energy shell of
•
•
Hydrogen all have the same energy
 or other single electron systems
We call orbitals with the same energy degenerate
For multielectron atoms, the energies of the
sublevels are split
 caused by charge interaction, shielding and penetration
• The lower the value of the l quantum number, the
less energy the sublevel has
 s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)
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Coulomb’s Law
• Coulomb’s Law describes the attractions and
repulsions between charged particles
• For like charges, the potential energy (E) is positive
and decreases as the particles get farther apart
 as r increases
• For opposite charges, the potential energy is negative
and becomes more negative as the particles get
closer together
• The strength of the interaction increases as the size of
the charges increases
 electrons are more strongly attracted to a nucleus with a 2+
charge than a nucleus with a 1+ charge
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Shielding & Effective Nuclear Charge
• Each electron in a multielectron atom
experiences both the attraction to the nucleus
and repulsion by other electrons in the atom
• These repulsions cause the electron to have a
net reduced attraction to the nucleus – it is
shielded from the nucleus
• The total amount of attraction that an electron
feels for the nucleus is called the effective
nuclear charge of the electron
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Penetration
• The closer an electron is to the nucleus, the
•
•
more attraction it experiences
The better an outer electron is at penetrating
through the electron cloud of inner electrons,
the more attraction it will have for the nucleus
The degree of penetration is related to the
orbital’s radial distribution function
in particular, the distance the maxima of the
function are from the nucleus
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Shielding & Penetration
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Penetration and Shielding
• The radial distribution function shows
•
•
that the 2s orbital penetrates more
deeply into the 1s orbital than does
the 2p
The weaker penetration of the 2p
sublevel means that electrons in the
2p sublevel experience more
repulsive force, they are more
shielded from the attractive force of
the nucleus
The deeper penetration of the 2s
electrons means electrons in the 2s
sublevel experience a greater
attractive force to the nucleus and are
not shielded as effectively
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Effect of Penetration and Shielding
• Penetration causes the energies of sublevels in the
•
•
same principal level to not be degenerate
In the fourth and fifth principal levels, the effects of
penetration become so important that the s orbital
lies lower in energy than the d orbitals of the
previous principal level
The energy separations between one set of orbitals
and the next become smaller beyond the 4s
the ordering can therefore vary among elements
causing variations in the electron configurations of
the transition metals and their ions
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6d
7s
6p
5d
6s
4f
5p
Energy
5s
5f
4d
4p
3d
4s
3p
3s
2p
2s
1s
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Notice the following:
1. because of penetration, sublevels within an
energy level are not degenerate
2. penetration of the 4th and higher energy
levels is so strong that their s sublevel is
lower in energy than the d sublevel of the
previous energy level
3. the energy difference between levels
becomes smaller for higher energy levels
(and can cause anomalous electron
configurations for certain elements)
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Filling the Orbitals with Electrons
• Energy levels and sublevels fill from lowest energy
to high
s → p → d → f
 Aufbau Principle
• Orbitals that are in the same sublevel have the
•
same energy
No more than two electrons per orbital
 Pauli Exclusion Principle
• When filling orbitals that have the same energy,
place one electron in each before completing pairs
 Hund’s Rule
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Electron Configuration of Atoms in their
Ground State
• The electron configuration is a listing of the
sublevels in order of filling with the number of
electrons in that sublevel written as a
superscript.
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
• A short-hand way of writing an electron
configuration is to use the symbol of the
previous noble gas in [] to represent all the inner
electrons, then just write the last set.
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1
= [Kr]5s1
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Order of Sublevel Filling
in Ground State Electron Configurations
Start by drawing a diagram
putting each energy shell on
a row and listing the sublevels,
(s, p, d, f), for that shell in
order of energy (left-to-right)
Next, draw arrows through
the diagonals, looping back
to the next diagonal
each time
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1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
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Electron Configurations
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Example: Write the full ground state orbital diagram
and electron configuration of manganese
Mn
Z = 25, therefore 25 e−
s sublevel holds 2 e−
−
p
sublevel
holds
6
e


      2 
e−
d1ssublevel2s
holds 10 2p
e−
+2 =3s
4e−
1s
  
2s
2p
3s 3p 3p

3d4s
−
+6
+2
=
12e
f sublevel holds 14

  4s
4p
4d
4f
−
+6 +2 = 20e−
+10
=
30e
3d
22s22p63s23p64s23d5
Therefore
the
electron
configuration
is
1s
Based on the order of sublevel filling, we will need the first seven sublevels
e−
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Practice — write the full ground state orbital diagram
and electron configuration of potassium.
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Practice — write the full ground state orbital diagram
and electron configuration of potassium, answer
K
Z = 19, therefore 19 e−
s sublevel holds 2 e−
−
p
sublevel
holds
6
e


      2 
e−
d1ssublevel2s
holds 10 2p
e−
+2 =3s
4e−
1s
  
2s
2p
3s 3p 3p

3d4s
−
+6
+2
=
12e
f sublevelthe
holds
14 configuration is 1s24s
Therefore
electron
2s22p64p
3s23p64d
4s1
+6 +2 = 20e−
e−
4f
Based on the order of sublevel filling, we will need the first six sublevels
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Valence Electrons
• The electrons in all the sublevels with the
•
•
highest principal energy shell are called
the valence electrons
Electrons in lower energy shells are called
core electrons
Chemists have observed that one of the
most important factors in the way an atom
behaves, both chemically and physically, is
the number of valence electrons
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Electron Configuration of Atoms in
their Ground State
• Kr = 36 electrons
1s22s22p63s23p64s23d104p6
 there are 28 core electrons and 8 valence electrons
• Rb = 37 electrons
•
•
1s22s22p63s23p64s23d104p65s1
[Kr]5s1
For the 5s1 electron in Rb the set of quantum
numbers is n = 5, l = 0, ml = 0, ms = +½
For an electron in the 2p sublevel, the set of
quantum numbers is n = 2, l = 1, ml = −1 or (0,+1),
and ms = −½ or (+½)
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Electron Configuration & the
Periodic Table
• The Group number corresponds to the number
•
•
of valence electrons
The length of each “block” is the maximum
number of electrons the sublevel can hold
The Period number corresponds to the
principal energy level of the valence electrons
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s1
1
2
3
4
5
6
7
s2
p1 p2 p3 p4 p5 s2
p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1
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Electron Configuration from
the Periodic Table
8A
1A
1
2
3
4
5
6
7
3A 4A 5A 6A 7A
2A
Ne
P
3s2
3p3
P = [Ne]3s23p3
P has five valence electrons
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Transition Elements
• For the d block metals, the principal energy level is one
less than valence shell
 one less than the Period number
 sometimes s electron “promoted” to d sublevel
Zn
Z = 30, Period 4, Group 2B
[Ar]4s23d10
4s
3d
• For the f block metals, the principal energy level is two less
than valence shell
 two less than the Period number they really belong to
 sometimes d electron in configuration
Eu
Z = 63, Period 6
[Xe]6s24f 7
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6s
38
4f
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Electron Configuration from
the Periodic Table
8A
1A
1
2
3
4
5
6
7
3A 4A 5A 6A 7A
2A
3d10
Ar
As
4s2
4p3
As = [Ar]4s23d104p3
As has five valence electrons
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Practice – Use the Periodic Table to write the short
electron configuration and short orbital diagram for
each of the following
• Na (at. no. 11)
[Ne]3s1
3s
• Te (at. no. 52)
[Kr]5s24d105p4
5s
• Tc (at. no. 43)
[Kr]5s24d5
5s
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5p
4d
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4d
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Irregular Electron Configurations
• We know that because of sublevel splitting, the
4s sublevel is lower in energy than the 3d; and
therefore the 4s fills before the 3d
• But the difference in energy is not large
• Some of the transition metals have irregular
electron configurations in which the ns only
partially fills before the (n−1)d or doesn’t fill at
all
• Therefore, their electron configuration must be
found experimentally
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Irregular Electron Configurations
•
•
•
•
•
•
•
•
•
•
•
•
Expected
Cr = [Ar]4s23d4
Cu = [Ar]4s23d9
Mo = [Kr]5s24d4
Ru = [Kr]5s24d6
Pd = [Kr]5s24d8
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Found Experimentally
Cr = [Ar]4s13d5
Cu = [Ar]4s13d10
Mo = [Kr]5s14d5
Ru = [Kr]5s14d7
Pd = [Kr]5s04d10
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Properties & Electron Configuration
• The properties of the elements
follow a periodic pattern
elements in the same column
have similar properties
the elements in a period show a
pattern that repeats
• The quantum-mechanical
model explains this because
the number of valence
electrons and the types of
orbitals they occupy are also
periodic
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The Noble Gas
Electron Configuration
• The noble gases have eight valence
electrons.
 except for He, which has only two
electrons
• We know the noble gases are
especially non-reactive
 He and Ne are practically inert
• The reason the noble gases are so
non-reactive is that the electron
configuration of the noble gases is
especially stable
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Everyone Wants to Be Like a Noble Gas!
The Alkali Metals
• The alkali metals have one more
•
electron than the previous noble gas
In their reactions, the alkali metals
tend to lose one electron, resulting
in the same electron configuration
as a noble gas
forming a cation with a 1+ charge
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Everyone Wants to Be Like a Noble Gas!
The Halogens
• The electron configurations of the
•
halogens all have one fewer electron
than the next noble gas
In their reactions with metals, the
halogens tend to gain an electron and
attain the electron configuration of the
next noble gas
 Forming an anion with charge 1−
• In their reactions with nonmetals, they
tend to share electrons with the other
nonmetal so that each attains the
electron configuration of a noble gas
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Eight Valence Electrons
• Quantum mechanical calculations show that eight
valence electrons should result in a very
unreactive atom
 an atom that is very stable
 the noble gases have eight valence electrons and are
all very stable and unreactive
 He has two valence electrons, but that fills its valence shell
• Conversely, elements that have either one more or
one less electron should be very reactive
 the halogen atoms have seven valence electrons and
are the most reactive nonmetals
 the alkali metals have one more electron than a noble
gas atom and are the most reactive metals
 as a group
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Electron Configuration &
Ion Charge
• We have seen that many metals and
nonmetals form one ion, and that the charge on
that ion is predictable based on its position on
the Periodic Table
Group 1A = 1+, Group 2A = 2+, Group 7A = 1−,
Group 6A = 2−, etc.
• These atoms form ions that will result in an
electron configuration that is the same as the
nearest noble gas
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Electron Configuration of Anions in
Their Ground State
• Anions are formed when nonmetal atoms gain
enough electrons to have eight valence
electrons
filling the s and p sublevels of the valence shell
• The sulfur atom has six valence electrons
S atom
= 1s22s22p63s23p4
• To have eight valence electrons, sulfur must
gain two more
S2− anion = 1s22s22p63s23p6
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Electron Configuration of Cations in
Their Ground State
• Cations are formed when a metal atom loses
all its valence electrons
resulting in a new lower energy level valence shell
however the process is always endothermic
• The magnesium atom has two valence
electrons
Mg atom
= 1s22s22p63s2
• When magnesium forms a cation, it loses its
valence electrons
Mg2+ cation = 1s22s22p6
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Trend in Atomic Radius – Main Group
• There are several methods for measuring
the radius of an atom, and they give
slightly different numbers
 van der Waals radius = nonbonding
 covalent radius = bonding radius
 atomic radius is an average radius of an atom
based on measuring large numbers of
elements and compounds
• Atomic Radius Increases down group
 valence shell farther from nucleus
 effective nuclear charge fairly close
• Atomic Radius Decreases across period
(left to right)
 adding electrons to same valence shell
 effective nuclear charge increases
 valence shell held closer
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Shielding
• In a multi-electron system, electrons are
•
simultaneously attracted to the nucleus and
repelled by each other
Outer electrons are shielded from nucleus by
the core electrons
screening or shielding effect
outer electrons do not effectively screen for each
other
• The shielding causes the outer electrons to not
experience the full strength of the nuclear
charge
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Effective Nuclear Charge
• The effective nuclear charge is net positive
•
charge that is attracting a particular electron
Z is the nuclear charge, S is the number of
electrons in lower energy levels
electrons in same energy level contribute to
screening, but very little so are not part of the
calculation
trend is s > p > d > f
Zeffective = Z − S
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Screening & Effective Nuclear Charge
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Quantum-Mechanical Explanation for the
Group Trend in Atomic Radius
• The size of an atom is related to the distance the
•
•
•
•
valence electrons are from the nucleus
The larger the orbital an electron is in, the farther
its most probable distance will be from the nucleus
and the less attraction it will have for the nucleus
Traversing down a group adds a principal energy
level
The larger the principal energy level an orbital is
in, the larger its volume
 quantum-mechanics predicts the atoms should
get larger down a column
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Quantum-Mechanical Explanation for the
Period Trend in Atomic Radius
• The larger the effective nuclear charge an electron
•
•
•
experiences, the stronger the attraction it will have
for the nucleus
The stronger the attraction the valence electrons
have for the nucleus, the closer their average
distance will be to the nucleus
Traversing across a period increases the effective
nuclear charge on the valence electrons
 quantum-mechanics predicts the atoms should
get smaller across a period
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Example 8.5: Choose the
Larger Atom in Each Pair
1.
2.
3.
4.
N or F,
F N is farther left
C or Ge
Ge, Ge is farther down
N or Al,
Al Al is farther down & left
Al or Ge? opposing trends
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Practice – Choose the
Larger Atom in Each Pair
••
••
••
••
C or O C is farther left in the period
Li or K K is farther down the column
C or Al Al is farther left and down
Se or I ? opposing trends
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Trends in Atomic Radius
Transition Metals
• Atoms in the same group increase in size
•
down the column
Atomic radii of transition metals roughly the
same size across the d block
much less difference than across main group
elements
valence shell ns2, not the (n−1)d electrons
effective nuclear charge on the ns2 electrons
approximately the same
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Electron Configurations of Main
Group Cations in Their Ground State
• Cations form when the atom loses electrons
from the valence shell
Al atom = 1s22s22p63s23p1
Al3+ ion = 1s22s22p6
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Electron Configurations of Transition
Metal Cations in Their Ground State
• When transition metals form cations, the first
•
•
electrons removed are the valence electrons, even
though other electrons were added after
Electrons may also be removed from the sublevel
closest to the valence shell after the valence
electrons
The iron atom has two valence electrons
Fe atom
= 1s22s22p63s23p64s23d6
• When iron forms a cation, it first loses its valence
electrons
Fe2+ cation = 1s22s22p63s23p63d6
• It can then lose 3d electrons
Fe3+ cation = 1s22s22p63s23p63d5
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Magnetic Properties of
Transition Metal Atoms & Ions
• Electron configurations that result in
unpaired electrons mean that the atom or
ion will have a net magnetic field – this is
called paramagnetism
will be attracted to a magnetic field
• Electron configurations that result in all
paired electrons mean that the atom or ion
will have no magnetic field – this is called
diamagnetism
slightly repelled by a magnetic field
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Transition Metal Atoms and Ions:
Electron Configuration &
Magnetic Properties
• Both Zn atoms and Zn2+ ions are diamagnetic
showing that the two 4s electrons are lost before the
3d
Zn atoms [Ar]4s23d10
Zn2+ ions [Ar]4s03d10
• Ag forms both Ag+ ions and, rarely, Ag2+
Ag atoms [Kr]5s14d10 are paramagnetic
Ag+ ions [Kr]4d10 are diamagnetic
Ag2+ ions [Kr]4d9 are paramagnetic
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Example 8.6c: Write the electron
configuration and determine whether the Fe
atom and Fe3+ ion are paramagnetic or
diamagnetic
• Fe Z = 26
• Previous noble gas = Ar
18 electrons
• Fe3+atom
ion == [Ar]4s
[Ar]4s023d
3d56
• Unpaired electrons
• Paramagnetic
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4s
3d
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Practice – Determine whether the following
are paramagnetic or diamagnetic
• Mn
Mn = [Ar]4s23d5 paramagnetic
4s
3d
• Sc3+
Sc = [Ar]4s23d1
Sc3+ = [Ar] diamagnetic
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Trends in Ionic Radius
• Ions in same group have same charge
• Ion size increases down the column
 higher valence shell, larger
• Cations smaller than neutral atoms; anions larger
•
than neutral atoms
Cations smaller than anions
 except Rb+ & Cs+ bigger or same size as F− and O2−
• Larger positive charge = smaller cation
 for isoelectronic species
 isoelectronic = same electron configuration
• Larger negative charge = larger anion
 for isoelectronic species
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Periodic Pattern – Ionic Radius (Å)
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Quantum-Mechanical Explanation for the
Trends in Cation Radius
• When atoms form cations, the valence electrons are
•
•
•
•
•
removed
The farthest electrons from the nucleus then are the p
or d electrons in the (n − 1) energy level
This results in the cation being smaller than the atom
These “new valence electrons” also experience a
larger effective nuclear charge than the “old valence
electrons,” shrinking the ion even more
Traversing down a group increases the (n − 1) level,
causing the cations to get larger
Traversing to the right across a period increases the
effective nuclear charge for isoelectronic cations,
causing the cations to get smaller
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Quantum-Mechanical Explanation for the
Trends in Anion Radius
• When atoms form anions, electrons are added to the
•
•
•
•
valence shell
These “new valence electrons” experience a smaller
effective nuclear charge than the “old valence
electrons,” increasing the size
The result is that the anion is larger than the atom
Traversing down a group increases the n level,
causing the anions to get larger
Traversing to the right across a period decreases the
effective nuclear charge for isoelectronic anions,
causing the anions to get larger
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Example 8.7: Choose the larger of
each pair
• S or S2−
 S2− is larger because there are more electrons (18 e−)
for the 16 protons to hold
 the anion is larger than the neutral atom
• Ca or Ca2+
 Ca is larger because its valence shell has been lost
from Ca2+
 the cation is always smaller than the neutral atom
• Br− or Kr
 the Br− is larger because it has fewer protons (35 p+) to
hold the 36 electrons than does Kr (36 p+)
 for isoelectronic species, the more negative the charge
the larger the atom or ion
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Practice – Order the following sets by
size (smallest to largest)
Zr4+, Ti4+, Hf4+
same column & charge,
therefore Ti4+ < Zr4+ < Hf4+
isoelectronic,
Na+, Mg2+, F−, Ne
therefore Mg2+ < Na+ < Ne < F−
I−, Br−, Ga3+, In+
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Ga3+ < In+ < Br− < I−
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Ionization Energy
• Minimum energy needed to remove an
electron from an atom or ion
gas state
endothermic process
valence electron easiest to remove, lowest IE
M(g) + IE1  M1+(g) + 1 eM+1(g) + IE2  M2+(g) + 1 efirst ionization energy = energy to remove electron
from neutral atom; 2nd IE = energy to remove from
1+ ion; etc.
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General Trends in 1st Ionization Energy
• The larger the effective nuclear charge on the
•
•
electron, the more energy it takes to remove it
The farther the most probable distance the
electron is from the nucleus, the less energy it
takes to remove it
1st IE decreases down the group
valence electron farther from nucleus
• 1st IE generally increases across the period
effective nuclear charge increases
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Quantum-Mechanical Explanation for the
Trends in First Ionization Energy
• The strength of attraction is related to the most
•
•
•
•
probable distance the valence electrons are from the
nucleus and the effective nuclear charge the valence
electrons experience
The larger the orbital an electron is in, the farther its
most probable distance will be from the nucleus and
the less attraction it will have for the nucleus
 quantum-mechanics predicts the atom’s first
ionization energy should get lower down a column
Traversing across a period increases the effective
nuclear charge on the valence electrons
 quantum-mechanics predicts the atom’s first
ionization energy should get larger across a period
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Example 8.8: Choose the atom in each pair
with the larger first ionization energy
1.
2.
3.
4.
Al or S,
S S is farther right
As or Sb
Sb, As is farther up
N or Si
Si, N is farther up & right
O or Cl? opposing trends
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Practice – Choose the atom with the largest
first ionization energy in each pair
• Mg or P
• Ag or Cu
• Ca or Rb
• P or Se ?
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Exceptions in the 1st IE Trends
• First Ionization Energy generally increases
from left to right across a Period
• Except from 2A to 3A, 5A to 6A
    


N
Be
1s 2s
2p
1s 2s
2p
B   
O     
1s 2s
2p
1s 2s
2p
Which is easier to remove an electron from
from,B,
N or Be?
O? Why?
Why?
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Exceptions in the
First Ionization Energy Trends,
Be and B
Be  
Be+  
1s 2s
2p
1s 2s
2p
To ionize Be you must break up a full sublevel, costs extra energy
  
B+  
1s 2s
2p
1s 2s
2p
When you ionize B you get a full sublevel, costs less energy
B
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Exceptions in the
First Ionization Energy Trends,
N and O
N

1s

2s
  
2p
N+ 
1s

2s
 
2p
To ionize N you must break up a half-full sublevel, costs extra energy
O 
1s
   
2s
2p
O+ 
1s
   
2s
2p
When you ionize O you get a half-full sublevel, costs less energy
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Trends in Successive
Ionization Energies
• Removal of each successive
electron costs more energy
 shrinkage in size due to having
more protons than electrons
 outer electrons closer to the
nucleus, therefore harder to
remove
• Regular increase in energy for
•
each successive valence
electron
Large increase in energy when
start removing core electrons
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Electron Affinity
• Energy released when an neutral atom gains an
electron
gas state
M(g) + 1e−  M1−(g) + EA
• Defined as exothermic (−), but may actually be
endothermic (+)
some alkali earth metals & all noble gases are
endothermic, WHY?
• The more energy that is released, the larger the
electron affinity
the more negative the number, the larger the EA
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Trends in Electron Affinity
• Alkali metals decrease electron affinity down the
column
 but not all groups do
 generally irregular increase in EA from 2nd period to 3rd
period
• “Generally” increases across period
 becomes more negative from left to right
 not absolute
 Group 5A generally lower EA than expected because
extra electron must pair
 Group 2A and 8A generally very low EA because added
electron goes into higher energy level or sublevel
• Highest EA in any period = halogen
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Properties of Metals & Nonmetals
• Metals
malleable & ductile
shiny, lusterous, reflect light
conduct heat and electricity
most oxides basic and ionic
form cations in solution
lose electrons in reactions – oxidized
• Nonmetals
brittle in solid state
dull, non-reflective solid surface
electrical and thermal insulators
most oxides are acidic and molecular
form anions and polyatomic anions
gain electrons in reactions – reduced
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Metallic Character
• Metallic character is how closely an element’s
properties match the ideal properties of a metal
more malleable and ductile, better conductors, and
easier to ionize
• Metallic character decreases left-to-right
across a period
metals are found at the left of the period and
nonmetals are to the right
• Metallic character increases down the column
nonmetals are found at the top of the middle Main
Group elements and metals are found at the bottom
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Quantum-Mechanical Explanation for the
Trends in Metallic Character
• Metals generally have smaller first ionization
•
•
energies and nonmetals generally have larger
electron affinities
except for the noble gases
 quantum-mechanics predicts the atom’s metallic
character should increase down a column because
the valence electrons are not held as strongly
 quantum-mechanics predicts the atom’s metallic
character should decrease across a period
because the valence electrons are held more
strongly and the electron affinity increases
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Example 8.9: Choose the
more metallic element in each pair
1.
2.
3.
4.
Sn or Te,
Te Sn is farther left
P or Sb
Sb, Sb is farther down
Ge or In
In, In is farther down & left
S or Br? opposing trends
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Practice – Choose the
more metallic element in each pair
• Mg or Al
• Si or Sn
• Br or Te
• Se or I ?
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Trends in the Alkali Metals
• Atomic radius increases down the column
• Ionization energy decreases down the column
• Very low ionization energies
 good reducing agents, easy to oxidize
 very reactive, not found uncombined in nature
 react with nonmetals to form salts
 compounds generally soluble in water  found in
seawater
• Electron affinity decreases down the column
• Melting point decreases down the column
 all very low MP for metals
• Density increases down the column
 except K
 in general, the increase in mass is greater than the
increase in volume
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Alkali Metals
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Trends in the Halogens
• Atomic radius increases down the column
• Ionization energy decreases down the column
• Very high electron affinities
 good oxidizing agents, easy to reduce
 very reactive, not found uncombined in nature
 react with metals to form salts
 compounds generally soluble in water  found in
seawater
• Reactivity increases down the column
• React with hydrogen to form HX, acids
• Melting point and boiling point increase down the
•
column
Density increases down the column
 in general, the increase in mass is greater than the
increase in volume
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Halogens
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Reactions of Alkali Metals with Halogens
• Alkali metals are oxidized
•
•
•
to the 1+ ion
Halogens are reduced to
the 1− ion
The ions then attach
together by ionic bonds
The reaction is
exothermic
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Example 8.10: Write a balanced chemical
reaction for the following
• Reaction between potassium metal and
bromine gas
K(s) + Br2(g) 
K(s) + Br2(g)  K+ Br
2 K(s) + Br2(g)  2 KBr(s)
(ionic compounds are all solids at room
temperature)
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Reactions of Alkali Metals with Water
• Alkali metals are oxidized to the 1+ ion
• H2O is split into H2(g) and OH− ion
• The Li, Na, and K are less dense than the
•
•
water – so float on top
The ions then attach together by ionic bonds
The reaction is exothermic,
and often the
heat released
ignites the H2(g)
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Example 8.10: Write a balanced chemical
reaction for the following
• Reaction between rubidium metal and liquid
water
Rb(s) + H2O(l) 
Rb(s) + H2O(l)  Rb+(aq) + OH(aq) + H2(g)
2 Rb(s) + 2 H2O(l)  2 Rb+(aq) + 2 OH(aq) + H2(g)
(alkali metal ionic compounds are soluble in water)
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Example 8.10: Write a balanced chemical
reaction for the following
• Reaction between chlorine gas and solid iodine
Cl2(g) + I2(s) 
Cl2(g) + I2(s)  ICl
write the halogen lower in the column first
assume 1:1 ratio, though others also exist
2 Cl2(g) + I2(s)  2 ICl(g)
(molecular compounds are found in all states at
room temperature, so predicting the state is not
always possible)
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Trends in the Noble Gases
• Atomic radius increases down the column
• Ionization energy decreases down the column
 very high IE
• Very unreactive
 only found uncombined in nature
 used as “inert” atmosphere when reactions with other
gases would be undersirable
• Melting point and boiling point increase down the
column
 all gases at room temperature
 very low boiling points
• Density increases down the column
 in general, the increase in mass is greater than the
increase in volume
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Noble Gases
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