Download File

Periodic Relationships Among the
Elements
Chapter 8
8.1
8.2
8.3
8.4
8.5
8.6
Development of the Periodic Table
Periodic Classification of the Elements
Periodic Variation in Physical Properties
Ionization Energy
Electron Affinity
Variation in Chemical Properties of the Representative Elements
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
8.1 Development of the Periodic Table
J. W. Dobereiner grouped elements into triads:
– Three elements with similar properties;
– Properties followed a pattern;
– The same element was in the middle of all trends
BUT not all elements belongs to a triad.
By mid 19th century, the atomic masses of many of the
elements had been determined.
John Newlands:
– hypothesized that the chemistry of the elements
might be related to their masses;
– arranged the known elements in order of increasing
atomic masses;
– discovered that every 7th element had similar
properties (noble gases were still unknown during
this time);
– Proposed the arrangement of elements into
‘octaves’.
BUT Newland’s ‘Law of Octaves’ did not work for some elements.
Julius Lothar Meyer & Dimitri Mendeleev focused on the relationship
between atomic mass & various physical and chemical properties of
elements.
Meyer aligned the elements according to
‘atomic volumes’.
Mendeleev:
- Credited with the origin of the
modern periodic table;
- Assumed that some elements
have not been discovered yet,
and deliberately left blanks in
his table.
H.G.J. Moseley:
- Discovered that the foundation to the
order of the elements was the atomic
number, not atomic mass.
The Modern Periodic Table:
- Elements are arranged in order of increasing atomic number;
- There are 7 horizontal rows called Period;
- There are 18 vertical columns called Group (or families);
- Classification of elements: metals, non-metals & metalloids;
- Electron configurations help explain the recurrence of physical &
chemical properties;
- Very important: understanding of general properties and trends within
a group or a period helps in predicting the properties of any element.
Metal: Elements that are usually
solids at room temperature. Most
elements are metals.
Non-Metal: Elements in the upper
right corner of the periodic Table.
Their chemical and physical
properties are different from metals.
Metalloid: Elements that lie on a diagonal line between the Metals
and non-metals. Their chemical and physical properties are
intermediate between the two.
Classification of the Elements
6
Periodic Patterns:
The chemical behavior of elements is determined by its electron
configuration.
Energy levels are quantized so roughly correspond to layers of
electrons around the nucleus.
A shell is all the electrons with the same value of n. n is a row in
the periodic table.
Each period begins with a new outer electron shell.
Each period ends with a completely filled outer
shell that has the maximum number of electrons
for that shell.
The number identifying the families identifies the
number of electrons in the outer shell, except
helium.
The outer shell electrons are responsible for
chemical reactions.
8.2 Periodic Classification of the Elements
According to the type of subshells being filled, the elements can
be divided into:
a) Representative Elements (Main Group Elements)
 Group 1A – 7A (Group 1, 2, 13, 14, 15, 16, 17);

All elements have incompletely filled s or p subshells of
the highest principal quantum number;
b) Noble Gases – Group 8A or Group 18; all have completely
filled p subshell.
c) Transition Elements (d block transition elements)
• Groups 1B, 3B – 8B (Group 3 – 11), with incompletely filled
d subshells.
d) Special metals – Zn, Cd, Hg.
e) Lanthanides – f block transition elements, have incompletely
filled f subshells.
f)
Actinides – also called f block transition elements.
4f
5f
9
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
Valence and Core Electrons
Valence Electrons - Those electrons outside of a closed
electron shell. These electrons take part in chemical
reactions.
Core Electrons - The electrons in the closed shells. They
cannot take part in chemical reactions.
Sodium 11 electrons:
Valence electrons
Core electrons
[Ne] 3s 1 --- one
1s 2 2s 2 2p 6 --- ten
Chlorine 17 electrons:
Valence electrons
Core electrons
[Ne] 3s 2 3p 5---- seven
1s2 2s 2 2p 6 ---- ten
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]
11
-1
-2
-3
+3
+1
+2
Cations and Anions Of Representative Elements
12
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
Isoelectronic is the condition where ions and/or
atoms possess the same number of electrons and
hence the same ground state electronic
configurations.
What neutral atom is isoelectronic with H− ?
8.2
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
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
8.2
8.3 Periodic Variation in Physical Properties
• Atomic size/radius – how big the atoms are;
• Ionic size/radius – how big ions are.
• Ionization energy – how much energy to remove an
electron;
• Electron afinity – the ability to accept one or more
electrons.
Periodic trends – How does the four properties above
vary as you go across a period?
Group trends – How does the four properties above
vary as you go down a group
Why? Explain why they vary…
8.3
Effective nuclear charge (Zeff) is the “positive charge”
felt by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Z
Core
Zeff
Radius
Na
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
8.3
Inner shell electrons shielded the nuclear charge
effectively from the outer shell electrons.
Example: Li 1s2 2s1
Electrons in 1s shield the nuclear charge and the 2s
electron feels or sees the 3+ nucleus as having 1+
charge. The net charge or Zeff is 1+.
Electrons in the same shell do not shield each other
effectively because, in average, they are at the same
distance.
Example: Be 1s2 2s2
The two electrons in 2s see the nucleus as having a 2+
charge.
Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
18
Atomic radius – variation in size.
Atomic size is usually represented by atomic radius.
Since an atom are usually bonded to another atom,
atomic radius is taken as half the distance between two
closest identical atoms.
One-half the distance
between the nuclei
centers of two adjacent
atoms.
(covalent radius)
(metallic radius)
Periodic variation in atomic size:
The atomic size become larger
going from top to bottom in a group
and become smaller from left to
right across a period.
Within a group, the Zeff felt by the
outer electrons is the same.
However, the core size is increasing
due to increasing value of n, thus,
increasing the atomic size.
Within a period, electrons are added
into the same shell. As Zeff is
increasing from left to right, the
attraction force is stronger and the
electrons are drawn inward.
The size variation is less pronounce
for the transition element.
Fig. 8.12
21
Trends in Atomic Radii
22
Question #1: Ranking Elements by Size
Problem: Rank the following elements in each group
according to decreasing size ( largest first!):
a) Na, K, Rb
b) Sr, In, Rb
c) Cl, Ar, K
d) Sr, Ca, Rb
Question #2: Ranking Elements by Size
Using only a periodic Table, arrange these atoms in
order of decreasing atomic size:
Br , Cl , Ge , K , S
Comparison of Atomic Radii with Ionic Radii
25
Ionic radius – variation in size.
Anion is always larger than their neutral atom. When
electrons are added, mutual repulsions increase which cause
the electrons to occupy larger volume.
Cation is always smaller than their neutral atom. When
electrons are remove, mutual repulsion decrease, electrons
are pulled closer together around the nucleus.
Periodic variation in ionic size
is similar to periodic variation
in atomic size.
For isoelectronic cations: M3+ < M2+ < M+
Example: Al3+ < Mg2+ < Na+
Explanation: all cations have the same number
of electrons (10 electrons) but Al has 13
protons, Mg has 12 protons whereas Na has
only 11 protons. Thus, the electron cloud in Al
will be pulled closer to the nucleus as
compared to Mg or Na.
For isoelectronic anions: X3– > X2– > X–
Example: N3– > O2– > F–
Explanation: all anions have the same number of electrons
(10 electrons) but the repulsion is much stronger in N3– than
in O2– or F–.
8.3
The Radii (in pm) of Ions of Familiar Elements
28
Question #3: Ranking Ions According to Size
Rank each set of Ions in order of increasing size:
a) K+, Rb+, Na+
b) Na+, O2-, F c) Fe+2, Fe+3
8.4 Ionization Energy
Ionization energy (IE) is the minimum energy (kJ/mol) required
to remove an electron from a gaseous atom in its ground state.
IE1 + X (g)
X+(g) + e-
IE1 first ionization energy
IE2 + X+(g)
X2+(g) + e-
IE2 second ionization energy
IE3 + X2+(g)
X3+(g) + e-
IE3 third ionization energy
IE1 < IE2 < IE3
Periodic trend of IE:
• IE generally decrease from top to bottom within a Group and
increase from left to right across a Period.
• Within a Group, n increases going down a Group, shells become
larger and outer electrons are farther away from the nucleus and
less tightly bound.
• Going across a Period, atomic size become smaller as Zeff increase.
The electrons are held more tightly and more difficult to be
removed.
8.4
Example:
Li  Li+ + e
IE1 = 520 kJ/mol
Li+  Li2+ + e
IE2 = 7 297 kJ/mol
Li2+  Li3+ + e
IE3 = 11 810 kJ/mol
IE1 < IE2 < IE3
The 1st electron removed is from 2s orbital and the product
gains stable noble gas configuration.
More energy is required to remove electron from an ion that
has reached noble gas configuration. Therefore, a sudden
increase (a large jump) in value of IE can be seen.
Ionization is always endothermic, because atom (or ion)
need to absorb energy to remove the electron(s).
8.4
32
Variation of the First Ionization Energy with Atomic Number
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
33
General Trends in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
34
Successive Ionization Energies
# Valence
Z Element Electrons IE1
IE2
IE3
IE4
IE5
IE6
IE7
3
Li
1
0.52 7.30 11.81
4
Be
2
0.92 1.76 14.85 21.01
5
B
3
0.80 2.43 3.66 25.02 32.82
6
C
4
1.09 2.35
7
N
5
1.40 2.86 4.58 7.48 9.44 53.27 64.36
8
O
6
1.31 3.39 5.30 7.47 10.98 13.33 71.33
4.62 6.22 37.83 47.28
Question #4:
Using the Periodic table only, rank the following elements
in each of the following sets in order of increasing 1st IE.
a) Ar, Ne, Rn
c) Be, Na, Mg
b) At, Bi, Po
d) Cl, K, Ar
Question #5:
Given the following series of ionization energies (in kJ/mol)
for an element in period 3, name the element and write its
electron configuration:
IE1
IE2
IE3
IE4
580
1,815
2,740
11,600
8.5 Electron Affinity
Electron affinity(EA) is the negative value of the energy
change that occurs when an electron is accepted by an
atom in the gaseous state to form an anion.
EA is a measure of the attraction the atom has for an
additional electron.
X (g) + eX-(g)
F (g) + e-
X-(g) DH = -328 kJ/mol
EA = +328 kJ/mol
O (g) + e-
O-(g) DH = -141 kJ/mol
EA = +141 kJ/mol
The large negative value for Fluorine
indicates that Fluorine atoms readily accept
an electron.
8.5
Trend in Electron Affinity
EA decreases when
going down a
Group (less
exothermic) and
increases when
going across a
period from left to
right (more
exothermic).
40
Variation of Electron Affinity With Atomic Number (H – Ba)
41
Summary of Trend
Periodic Table and Periodic Trends
1. Electron Configuration
3. Ionization Energy: Largest toward NE of PT
4. Electron Affinity: Most favorable NE of PT
2. Atomic Radius: Largest toward SW corner of PT
8.6 Variation in Chemical Properties of the Representative Elements
Elements in the same group resembles one another in
chemical behaviour because they have similar outer
electron configuration.
Diagonal relationships:
Similarities between pairs of elements in different
groups and periods in the Periodic Table.
8.3
Li, Be and B (elements form 2nd period) exhibits many
similarities to those elements located diagonally below
them in the Periodic Table.
The reason:
a) Closeness of the charge densities of their cations;
b) Similar measurement of atomic and ionic radii;
Li = 152 pm
Li+ = 76 pm
Mg = 160 pm
Mg2+ = 72 pm
Both cations will react similarly with the same anion &
form the same type of compound:
Li + N2  Li3N, lithium nitride
Mg + N2  Mg3N2 , magnesium nitride
8.6
Group 1A Elements (ns1, n  2)
M
M+1 + 1e-
2M(s) + 2H2O(l)
2M2O(s)
Increasing reactivity
4M(s) + O2(g)
2MOH(aq) + H2(g)
45
Group 1A Elements (ns1, n  2)
46
Group 2A Elements (ns2, n  2)
M
M+2 + 2e-
Be(s) + 2H2O(l)
Mg(s) + 2H2O(g)
Mg(OH)2(aq) + H2(g)
M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba
Increasing reactivity
M(s) + 2H2O(l)
No Reaction
47
Group 2A Elements (ns2, n  2)
48
Group 3A Elements (ns2np1, n  2)
4Al(s) + 3O2(g)
2Al(s) + 6H+(aq)
2Al2O3(s)
2Al3+(aq) + 3H2(g)
49
Group 3A Elements (ns2np1, n  2)
50
Group 4A Elements (ns2np2, n  2)
Sn(s) + 2H+(aq)
Sn2+(aq) + H2 (g)
Pb(s) + 2H+(aq)
Pb2+(aq) + H2 (g)
51
Group 4A Elements (ns2np2, n  2)
52
Group 5A Elements (ns2np3, n  2)
N2O5(s) + H2O(l)
P4O10(s) + 6H2O(l)
2HNO3(aq)
4H3PO4(aq)
53
Group 5A Elements (ns2np3, n  2)
54
Group 6A Elements (ns2np4, n  2)
SO3(g) + H2O(l)
H2SO4(aq)
55
Group 6A Elements (ns2np4, n  2)
56
Group 7A Elements (ns2np5, n  2)
X2(g) + H2(g)
X-1
2HX(g)
Increasing reactivity
X + 1e-
57
Group 7A Elements (ns2np5, n  2)
58
Group 8A Elements (ns2np6, n  2)
Completely filled ns and np subshells.
Highest ionization energy of all elements.
No tendency to accept extra electrons.
59
Compounds of the Noble Gases
A number of xenon compounds XeF4, XeO3,
XeO4, XeOF4 exist.
A few krypton compounds (KrF2, for example)
have been prepared.
60
Comparison of Group 1A and 1B
The metals in these two groups have similar outer
electron configurations, with one electron in the
outermost s orbital.
Chemical properties are quite different due to difference
in the ionization energy.
Lower IE1, more reactive
61
Properties of Oxides Across a Period
Metallic characters of the elements decreases from left
to right across a Period and increases from top to
bottom within a Group.
Most oxides of representative metals are ionic
compounds; in water they act as bases, producing OH
ions and reacting with acids.
Example: Barium forms the basic oxide BaO:
Reacts with water:
BaO (s) + H2O (ℓ)  Ba2+ (aq) + 2 OH– (aq)
Reacts with acid:
BaO (s) + H+ (aq)  Ba2+ (aq) + H2O (ℓ)
Other examples of basic oxides:
Li2O , Na2O , K2O , MgO , CaO , SrO
8.6
Non-metal oxides are covalent compounds; in water they
act as acids, producing H+ ions and reacting with bases.
Example: Sulfur forms the acidic oxide SO2:
Reacts with water:
SO2 (g) + H2O (ℓ) ⇌ H2SO3 (aq) ⇌ H+ (aq) + HSO3– (aq)
Reacts with base:
SO2 (g) + 2 OH– (aq)  SO32– (aq) + H2O (ℓ)
Other examples of acidic oxides:
Cl2O7 , SO3 , N2O5 , N2O3 , P4O10 , CO2 , B2O3
8.6
Some metals and many metalloids form oxides that are
amphoteric: they can act as acids and as bases.
Example: Al2O3 is an amphoteric oxide:
Reacts with acid:
Al2O3 (s) + 6 H+ (aq)  2 Al3+ (aq) + 3 H2O (ℓ)
Reacts with base:
Al2O3 (s) + 2 OH– (aq) + 3 H2O ()  2 Al(OH)4– (aq)
Other examples of amphoteric oxides:
BeO , Ga2O3 , SnO2 , SnO , PbO2
8.6
Properties of Oxides Across a Period
basic
acidic
8.6
The Trends in Acid-Base Behavior of Elemental Oxides:
As the elements become more metallic down a Group,
their oxides become more basic.
As the elements become less metallic across a period,
their oxides become more acidic.
Fig. 8.22