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Periodic Table
Mendeleev
• ordered 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 patterns to predict properties of
undiscovered elements
• where atomic mass order did not fit other
properties, he reordered by different properties
Chapter 8
Periodic
Properties of
the Elements
The beginning of a periodic Trend
NM
H
1.0
H2
M
Li
6.9
Li2O
b
M
LiH
9.0
Be
Na2O M
b
M
Na
23.0
NaH 24.3
M
K2O
b
K
39.1
What vs. Why
H2O
a/b
KH
BeO NM
a/b
B
BeH2 10.8
MgO
b
Al
MgH2 27.0
M
Ca
40.1
CaO
b
CaH2
C
BH3 12.0
M
Mg
B2O3 NM
a
Al2O3
a/b
M/NM
Si
AlH3 28.1
?
CO2 NM
a
N
N2O5 NM
a
CH4 14.0
NH3 16.0
NM
P4O10 NM
a
SiO2
a
P
NM
F
H2O 19.0
S
SO3
a
HF
NM
Cl
Cl2O7
a
PH3 32.1
H2S 35.5
HCl
M/NM As2O5 NM
a/b
SeO3 NM
a
Br2O7
a
SiH4 31.0
?
O2
O
As
74.9
AsH3 79.0
As
H2Se 79.9
Br
HBr
M = metal, NM = nonmetal, M/NM = metalloid
• 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 and its theory of
electron configurations is what explains why
the periodic trends in the properties exist.
a = acidic oxide, b = basic oxide, a/b = amphoteric oxide
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4
First: the Electron Spin
•
Electron pairs residing in the same orbital are
required to have opposing spins.
– This causes electrons to behave like tiny bar magnets.
Electron Configurations
• An electron configuration of an atom is a
particular distribution of electrons
among available subshells.
– The notation for a configuration lists the
subshell symbols sequentially with a
superscript indicating the number of
electrons occupying that subshell.
– For example, lithium (at. # 3) has two
electrons in the “1s” subshell and one
electron in the “2s” sub shell is:
Lithium: 1s2 2s1
Pauli Exclusion Principle
• No two electrons can have the same
four quantum numbers.
– In other words, an orbital can hold at most
only two electrons, and then only if the
electrons have opposite spins.
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Summary of Orbital Types
• The maximum number of electrons and
their orbital shells are:
Maximum
Number of Number of
Sub shell
Orbitals
Electrons
s
1
2
p
3
6
d
5
10
f
7
14
8
Order for Filling Atomic
Subshells
The Aufbau Principle
• Every atom has an infinite number of
possible electron configurations.
– The configuration associated with the lowest
energy level of the atom is called the “ground
state.”
– Other configurations correspond to “excited
states.”
– The Aufbau Principle is a scheme used to
reproduce the ground state electron
configurations of atoms by following the
“building up” order.
– You need to remember the order of filling
the orbitals (next slide)
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Another Way to Remember the
Filling Order
1s
2s
3s
4s
5s
6s
or
2p
3p
4p
5p
6p
3d
4d 4f etc
5d 5f etc
6d 6f etc
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s etc
Example of Subshell Filling Order
• With sodium (Z = 11), the 3s subshell begins
to fill.
Z=11
Sodium
1s22s22p63s1
or
[Ne]3s1
Z=12
Magnesium
1s22s22p23s2
or
[Ne]3s2
• Then the 3p subshell begins to fill.
• Progressing from left to right across
each consecutive row will give the
correct filling order… 1s,2s,2p,3s,3p…
11
Z=13
:
:
Z=18
Aluminum
1s22s22p63s23p1
or [Ne]3s23p1
Argon
1s22s22p63s23p6
or [Ne]3s23p6
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Abbreviated form
Using the Aufbau Principle
• The “building up” order corresponds, for the
most part, to increasing energy of the
subshells.
– By filling orbitals of the lowest energy first, you
usually get the lowest total energy (“ground
state”) of the atom.
– Remember, the number of electrons in the
neutral atom equals the atomic number, Z.
“Abbreviated form” of electron configurations uses the
closet noble gas to indicate that its inner electrons have filled
that particular noble gas configuration.
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Electronic config. can be written in 1) order of filling or
2) in order of increasing principal quantum #.
The table above follows the latter.
Ground State Electron
Configuration Examples
Sample Problem
Valid and Invalid Electron Configurations
• Which ground-state electron
configurations are INCORRECT?
WHY?
Z=4 Beryllium 1s22s2 or [He]2s2
Z=3 Lithium
Z=5
Z=6
Z=7
Z=8
Z=9
Z=10
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
1s22s1 or [He]2s1
1s22s22p1
1s22s22p2
1s22s22p3
1s22s22p4
1s22s22p5
1s22s22p6
or
or
or
or
or
or
[He]2s22p1
[He]2s22p2
[He]2s22p3
[He]2s22p4
[He]2s22p5
[He]2s22p6
1.
2.
3.
4.
5.
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Abbreviated forms
15
Cr: [Ar] 3d6
Ca: [Ar] 4s2
Na: 1s2 2s2 2p6 3s1
Zn: [Ar] 3d10 4s2
Kr: [Ar] 3d10 4s2 4p6
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Electron Configurations
of elements and their orbital diagram
Orbital Filling Diagrams
• An orbital filling diagram is used to show how
the orbitals of a subshell are occupied by
electrons.
– Each orbital is represented by a box or circle.
– Each group of orbitals is labelled by its subshell
notation. Electrons are represented by arrows (up or
down)
– : up for spin= +1/2 and down for spin = -1/2
For He
1s
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Sample Problem
Hund’s Rule
• Give the expected ground-state
electron configurations (full as well as
abbrev.) and the orbital-filling diagrams
for:
1. P (Z=15)
2. Zn (Z=30)
3. Ca (Z=20)
• Consider carbon (Z = 6) with the ground
state configuration 1s2 2s2 2p2.
• Hund’s rule states that the lowest
energy arrangement (the “ground state”)
of electrons in a subshell is obtained by
putting electrons into separate orbitals
of the subshell with the same spin
before pairing electrons.
Carbon's orbital
diagram is
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↑↓
↑↓
↑ ↑
1s
2s
2p
20
20
Valence Electrons
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• The electrons in all the subshells 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.
The # of valence electrons = group #
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Electron Configuration and the
Periodic Table
• The group number corresponds to the
number of valence electrons.
• The number of columns in each “block” is
the maximum number of electrons that
sublevel can hold.
• The period number corresponds to the
principal energy level of the valence
electrons.
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Writing electron configurations from
their position in the Periodic Table
• Progressing from left to right across
each consecutive row will give the
correct filling order… 1s,2s,2p,3s,3p…
Anomalous 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
anomalous electron configurations in which
the (n)s only partially fills before the (n−1)d
or doesn’t fill at all.
• Therefore, their electron configurations must
be found experimentally.
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Anomalous Electron Configurations
Expected
• Cr = [Ar]4s23d4
• Cu = [Ar]4s23d9
• Mo = [Kr]5s24d4
• Pd = [Kr]5s24d8
Found experimentally
• Cr = [Ar]4s13d5
• Cu = [Ar]4s13d10
• Mo = [Kr]5s14d5
• Pd = [Kr]5s04d10
That is, if the “d” orbital can be half filled or
completely filled , it will take from the “s” orbital
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Trends in Periodic Properties
• The periodic law states that when the
elements are arranged by atomic
number, their physical and chemical
properties vary periodically.
• We will look at three periodic properties:
A) Atomic radius
B) Ionization energy
C) Electron affinity
Effective Nuclear Charge
Why the Trends?
Shielding and Effective Nuclear Charge
• In a multielectron system, electrons are
simultaneously attracted to the nucleus and
repelled by each other.
• Outer electrons are shielded from the nucleus by
the core electrons.
– Thus a screening effect
– But outer electrons do not effectively screen for
each other.
• Because of this shielding, the outer electrons do
not experience the full strength of the nuclear
charge.
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Screening and
Effective Nuclear Charge
• Effective nuclear charge (Zeffective) is the
net positive charge that is attracting a
particular electron.
• Z is the nuclear charge; S is the charge due
to electrons in lower energy levels.
– Electrons in same energy level do 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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A) Atomic Radius Trend
Figure
Representation
of atomic radii
(covalent radii)
of the maingroup elements
• Within each period (horizontal row), the
atomic radius tends to decrease with
increasing atomic number (or increasing
effective nuclear charge).
• Within each group (vertical column), the
atomic radius tends to increase with the
period number.
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B) Ionization Energy Trend
i.e, Na (g) → Na +(g) + 1 e-
Ionization Energy Trend
IE1= 496 kJ/mol
• Ionization energy trends
– There is a general trend that ionization
energies increase with atomic number
within a given period.
– This follows the trend in size, as it is more
difficult to remove an electron that is closer
to the nucleus.
– For the same reason, we find that
ionization energies, again following the
trend in size, decrease as we descend a
column of elements.
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C) Electron Affinity
Electron Affinity Trend
• The electron affinity is the energy
change for the process of adding an
electron to a neutral atom in the
gaseous state to form a negative ion.
• The more negative the electron affinity,
the more stable the negative ion that is
formed.
• The general trend goes from lower left
to upper right as electron affinities
become more negative.
• The next Table gives the electron
affinities of the main-group elements.
– For a chlorine atom, the first electron
affinity is illustrated by:
Cl (g) + e-
→
Cl - (g)
EA= -349 kJ/mol
Also by showing the species” electron configuration
Cl([Ne]3s 2 3p 5 ) + e − → Cl − ([Ne]3s 2 3p 6 )
Electron Affinity = -349 kJ/mol
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The Main-Group Elements
• The physical and chemical properties of
the main-group elements clearly display
“periodic behaviour’.
e-
higher
affinities
– Variations of metallic-non-metallic
character.
– Basic-acidic behaviour of the oxides
- atomic radii, electron affinity, etc
– Reaction with water
– Readily oxidized or reduced to form ions
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40
Properties of Groups
IA and IIA
• Group IA (Alkali Metals)
– Largest atomic radii
– React violently with water to form H2
– Readily oxidized to 1+
– Oxides dissolved in water form basic
solutions
• Group IIA (Alkaline Earth Metals)
– Readily ionized to 2+
– React with water to form H2
– Oxides dissolved in water form basic
solutions
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 their extra electron,
resulting in the same electron
configuration as a noble gas.
– forming a cation with a 1+ charge
Also the case for The Alkali Earth Metals
which tend to loss two electrons, resulting in a
noble gas electron configuration i.e., Ca 2+
Properties of Groups
VIA, VIIA, VIIIA
Properties of Groups
IIIA, IVA and VA
• Group VI A
• Group III A
– Metals (except for boron)
– Several oxidation states (commonly 3+)
• Group IV A
(i.e., C)
– Form the most covalent compounds
– Oxidation numbers vary between 4+ and 4-
• Group V A
(i.e., N, P)
– Form anions generally ( 3-), though positive
oxidation states are possible
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(i.e., O, S)
– anions generally, though positive oxidation
states are possible
– React vigorously with alkali and alkali earth
metals and nonmetals
• Group VII A (halogens)
–
–
–
–
Form mono anions
High electronegativity (electron affinity)
Diatomic gases
Most reactive of the nonmetals (F2 in particular)
• Group VIII A (noble gases)
– Minimal reactivity
– Monatomic gases;
Closed (fill) shell
44
Everyone Wants to Be Like a Noble Gas!—
The Halogens
• 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.
Noble Gas
Electron Configuration
• The noble gases have eight valence
electrons.
– Except for He, which has only two electrons.
• We know that the noble gases are
especially nonreactive.
– forming an anion with charge 1−.
– He and Ne are practically inert.
• 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.
• The noble gases are so nonreactive
because the electron configuration of the
noble gases is especially stable.
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Electron Configuration and
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.
48
Electron Configuration of
Anions in Their Ground State
Electron Configuration of
Cations in Their Ground State
• Anions are formed when atoms gain enough
electrons to have eight valence electrons.
• Cations are formed when an atom loses all its
valence electrons.
– filling the s and p sublevels of the valence shell
– resulting in a new lower energy level valence shell
– However, the process is always endothermic.
• The sulfur atom has six valence electrons.
S atom
= 1s22s22p63s23p4
• The magnesium atom has 2 valence electrons.
• In order to have eight valence electrons, it
must gain two more.
Mg atom
= 1s22s22p63s2
• When it forms a cation, it loses its valence
electrons.
S2− anion = 1s22s22p63s23p6
Mg2+ cation = 1s22s22p6
50
Magnetic Properties of
Transition Metal Atoms and Ions
Electron Configuration of
Cations in Their Ground State
• Electron configurations that result in
unpaired electrons mean that the atom or
ion will have a net magnetic field. This is
called paramagnetism.
• Cations form when the atom loses
electrons from the valence shell.
• For transition metals, electrons may also be
removed from the sublevel closest to the
valence shell.
– 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.
1s22s22p63s23p1
Al atom =
Al3+ ion = 1s22s22p6
Fe atom =1s22s22p63s23p64s23d6
Fe2+ ion = 1s22s22p63s23p63d6
Fe3+ ion = 1s22s22p63s23p63d5
=1s22s22p63s23p64s13d10
Cu atom
Cu+ ion = 1s22s22p63s23p63d10
– slightly repelled by a magnetic field
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Practice—Determine whether the following
are paramagnetic or diamagnetic.
another Term to know:
isoelectronic = ions with same electron configuration
i.e., Ca2+ is isoelectronic with K+
D) Trends in Ionic Radius
• Mn
Mn = [Ar]4s23d5
4s
paramagnetic !
3d
• Sc3+
Sc = [Ar]4s23d1 paramagnetic!
Sc3+ = [Ar] diamagnetic !
• 1) ion size increases down the group
– higher valence shell = larger
•
•
•
•
2) cations are smaller than the neutral atom
3) anions are bigger than the neutral atom
4) cations smaller than anions
5) larger positive charge = smaller cation
– for isoelectronic species i.e., Na+ > Mg2+ > Al3+
• 6) larger negative charge = larger anion
– for isoelectronic species N3 - > O2 - > F -
Main Group 1A, 2A &3A
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Main group 6A & 7A
55
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Ionic vs. neutral atom radii
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 electrons have
been lost to form Ca2+.
– The cation is always smaller than the neutral atom.
• Br− or Kr
– 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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57
Ionic Radii
Ionic Radii
• Within an isoelectronic group of ions, the
one with the greatest nuclear charge will be
the smallest.
• In this group, calcium has the greatest
nuclear charge and is, therefore, the
smallest.
– For example, look at the ions listed below:
20 Ca
2+
19 K
+
20 Ca
18 Ar
17 Cl
-
16S
All have 18 electrons
– Note that they all have the same number of
electrons, but different numbers of protons.
2-
2+
< 19K + < 18 Ar < 17Cl - < 16S 2All have 18 electrons
– Sulfur has only 16 protons to attract its 18
electrons and, therefore, has the largest radius.
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Practice—Order the following sets by size
(smallest to largest).
Zr4+, Ti4+, Hf4+
same column and charge;
therefore, Ti4+ < Zr4+ < Hf4+
isoelectronic;
Na+, Mg2+, F−, Ne
therefore, Mg2+ < Na+ < Ne < F−
I−, Br−, Ga3+, In+
Ga3+ < In+ < Br− < I−