Download lecture slides of chap8

Periodic Relationships Among
the Elements
Chapter 8
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Development of the Periodic Table
• Based on chemical and physical properties of the elements
• In 1869, ordered according to relative atomic mass
• Today, ordered according to atomic number (the of protons)
• Electron configurations help explain recurring properties
Classification of the Elements
8.2
Outermost subshell being filled with electrons
7.8
Electron Configurations of the Elements
Valence Electrons
• the outer electrons of an atom, the ones involved in
bonding
• for the representative elements (main group elements), the
number of valence electrons is equal to the A Group Number
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
4f
5f
8.2
Which of the following electron configurations do represent
similar chemical properties of their atoms?
(i) 1s22s22p63s2
(ii) 1s22s22p3
(iii) 1s22s22p63s23p64s23d104p5
(iv) 1s22s1
(v) 1s22s22p6
(vi) 1s22s22p63s23p3
(ii), (vi)
Practice problem of chap8:Q2
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al
[Ne]3s23p1
Al3+
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
[Ne]
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
Remember: electrons are first
removed from orbitals with the
highest principal quantum number.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
8.2
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
8.2
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: have the same number of electrons
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
Which of the following pairs are not isoelectronic
with Ar?
(a) K+ and Cl(b) Ca2+ and S2(b) Al3+and N3(d) P3- and Sc3+
(c) Rh3+ and Ir3+
Practice problem of chap8:Q5
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
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
8.2
A metal ion with a net +3 charge has five electrons in the
3d subshell. What is this metal?
(a) Cr (b) Mn (c) Fe (d) Co (e) Ni
This species has +3 charges, which indicates that it
has three more protons than the electrons. According
to the question that it has five electrons in the 3d
subshell, and thus the total electrons in valence
shells for its atomic type will be 8. Note that the
transition metals (with d electrons) losing its s
electrons prior to its d electrons and thus the valence
electron configuration will be 4s23d6, which is Fe.
Practice problem of chap8:Q3
Periodic Variation in Physical Properties
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
Na
Z
Core
Zeff
11
10
1
Radius (pm)
186 The larger the
Mg
12
10
2
Al
13
10
3
effective nuclear
charge, the
160 stronger hold of
the nucleus on
143 electrons.
Si
14
10
4
132
8.3
Shielding Effect in Many-Electron Atoms
Electrons close to the nucleus shield electrons with higher
principal quantum numbers.
Therefore the outer electrons do not feel the entire positive
charge of the nucleus.
Electrons will not exert shielding effect on the electrons in
the same subshell.
7.8
Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
The larger the effective nuclear charge, the stronger hold of
the nucleus on electrons.
8.3
In metal, radius is
defined as one half the
distance between the
centers of adjacent
atoms.
Radius is defined as
one half the distance
between the centers of
atoms in the molecule.
8.3
8.3
Atomic Radii
8.3
Sizes of atoms: Trends
• size increases while going down a Group
• Why?
• Because orbital size increases with increasing n
number, the size of atom increases. The highest energy
electrons can be farther away from the nucleus.
• size decreases going across Period
• Why?
• Zeff increases. The larger the effective nuclear charge,
the stronger hold of the nucleus on electrons. The
electron clouds shrink.
• Remember n does NOT increases--e- cannot be farther
from nucleus.
Ionic Radius
Cation is always smaller than atom from which it is formed.
This is because the nuclear charge remains the same but
the reduced electron repulsion resulting from removal of
electrons make the electron clouds shrink.
Anion is always larger than atom from which it is formed.
This is because the nuclear charge remains the same but
electron repulsion resulting from the additional electron
enlarges the electron clouds.
8.3
Comparison of Atomic Radii with Ionic Radii
8.3
The Radii (in pm) of Ions of Familiar Elements
8.3
•From top to bottom, both the atomic and ionic radius
increase within group.
•Across a period the anions are usually larger than
cations.
• For ions derived from different groups, size comparison
is meaningful only if the ions are isoelectronic.
E.g. Na+ (Z=11) is smaller than F- (Z=9).
The larger effective charge results in a smaller radius.
Radius of tripositive ions < dipositive ions<unipositive ions
Al3+ <Mg2+ <Na+
Radius of uninegative ions < dinegative ions
O2->F-
Which of the following is a correct order of atomic radii?
(a) Ar < P < Na < Ca < Cs (b) Cs < Rb < K < Na < Li
(c) Na < Mg < Al < Si < P (d) Rb < Mg < C < F < He
(e) Li < H < Al < K < Ar
Which of the following is a correct order of ionic radii?
(a) Na+ < Mg2+ < Al3+ < O2- < F- < N3(b) Al3+ < Mg2+ < Na+ < F- < O2- < N3(c) F- < O2- < N3- < Na+ < Mg2+ < Al3+
(d) Al3+ < Mg2+ < Na+ < N3- < O2- < F(e) Na+ < Mg2+ < Al3+ < O2- < N3- < F-
Practice problem of chap8:Q6,7
Ionization energy (I.E.) is the minimum energy (kJ/mol)
required to remove an electron from a gaseous atom in its
ground state.
I1 + X (g)
I2 + X+(g)
I3 +
X2+
(g)
X+(g) + e-
I1 first ionization energy, remove 1st electron
X2+(g) + e- I2 second ionization energy, remove 2nd electro
X3+(g)
+
rd electron
I
third
ionization
energy,
remove
3
3
e
I1 < I2 < I3
8.4
8.4
General Trend in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
8.4
First IE:
Trends down a Group
• remember that Cs is more reactive than Na
• Easier to remove an e- from Cs than from Na
• As size increases, I.E. decreases
Trends across a Period
• remember that Cl gains rather than loses an e• Easier to remove an e- from Al than from Cl
• As size decreases, I.E. increases
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
8.4
Exceptions: Ionization Energy
1. Occur between Group 2A and 3A
Group 3A: ns2np1. one single electron in p orbital.
• removing the first p electron is less than expected
•Why? This p electron is shielded by inner ns2 electrons.
Less energy is needed to remove a single p electron than
to remove a pair of s electron of the same n level.
2. Occur Group 5A and 6A
5A: ns2np3 6A: ns2np4
• removing the fourth p electron is less than expected
•Why? 2nd electron in a p orbital increases the
electron-electron repulsion, which makes it easier to
ionize an atom of Group 6A.
Arrange the following in order of the increasing
first ionization energy: Na, Cl, Al, S, Cs
Hint: Ionization energy increases across a row of
the periodic table and decreases down a column
or group.
Cs < Na < Al < S < Cl
Group practice problem:Q10
Electron affinity (EA) is the negative of the energy change
that occurs when an electron is accepted by an atom in the
gaseous state to form an anion.
X (g) + e-
X-(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
8.5
8.5
Variation of Electron Affinity With Atomic Number (H – Ba)
Overall trend: increases
from left to right across the
period. The values varied
little in the group. The EA
of metals are generally
lower than those of
nonmetals.
8.5
Variation in Chemical Properties
of
Representative Elements
Group 1A Elements (ns1, n  2)
8.6
General trends in chemical properties
Hydrogen
• Non-representative member of GroupIA
• Nonmetal
• Diatomic gas
• Relative high I.E.
• Forms H+ (aq): resembles Group IA
• Forms H- : resembles Group VIIA (halogens)
• Reactions:
+ oxygen --> water
+ Group IA (and Group IIA) --> metal hydride
Metal hydride + water --> metal hydroxide + H2
Group 1A Elements (ns1, n  2)
M+1 + 1e-
2M(s) + 2H2O(l)
4M(s) + O2(g)
2MOH(aq) + H2(g)
2M2O(s)
Increasing reactivity
M
• Alkali Metals
• Very Reactive: not found as elements in nature
• Form M+ ions
• Hydroxides are strong bases
• Reactions with oxygen:
Li --> lithium oxide [Li2O (s)]
Na --> sodium peroxide [Na2O2 (s)]
K --> potassium superoxide [KO2 (s)]
8.6
Group 2A Elements (ns2, n  2)
8.6
Group 2A Elements (ns2, n  2)
M
M+2 + 2e-
• Alkaline Earth Metals
• Form M2+ ions
• Be and some Mg compounds are molecular
• Hydroxides are strong bases
• Rxn + water:
Be = no reaction
M(s) + 2H2O(l)
Mg(OH)2(aq) + H2(g)
M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba
• Rxn + oxygen:
Be and Mg + heat --> oxides
Other Group IIA --> oxides
Increasing reactivity
Mg(s) + 2H2O(g)
8.6
Group 3A Elements (ns2np1, n  2)
8.6
Group 3A Elements (ns2np1, n  2)
• Boron = metalloid
+ oxygen = NR
+ water = NR
• Aluminum
Forms Al3+ ion
2Al2O3(s)
+ oxygen --> Al2O3 (s) 4Al(s) + 3O2(g)
+ acid H2
2Al(s) + 6H+(aq)
2Al3+(aq) + 3H2(g)
• Other Group IIIA:
Form M+ (more stable ion) and M3+ ion
• Molecular compounds also formed
Al + H2 --> AlH3 (resembles rxn of Be + H2)
8.6
Group 4A Elements (ns2np2, n  2)
8.6
Group 4A Elements (ns2np2, n  2)
• Carbon = nonmetal
Molecular compounds formed
Oxidation state + 4 is more stable
CO2 is more stable than CO
• Silicon & Germanium = metalloids
Oxidation state of Si = +4 only [SiO2]
• Tin & lead = metals:
Oxidation state of Sn = +4 more stable than +2
Oxidation state of Pb = +2 more stable than +4
Reaction: + acids --> M2+ ions + H2
Sn(s) + 2H+(aq)
Sn2+(aq) + H2 (g)
Pb(s) + 2H+(aq)
Pb2+(aq) + H2 (g)
8.6
Group 5A Elements (ns2np3, n  2)
8.6
Group 5A Elements (ns2np3, n  2)
• Nitrogen & phosphorus = nonmetals
Elements = N2 (g) and P4 (s)
Reactions:
N2 + Group IA --> nitrides (N3-) ions
N2 (or P) + oxygen --> many oxides
oxides + water --> acids
N2O5(s) + H2O(l)
2HNO3(aq)
P4O10(s) + 6H2O(l)
4H3PO4(aq)
• Arsenic & antimony = metalloids
• Bismuth = metal:
less reactive than metals of Group IVA
8.6
Group 6A Elements (ns2np4, n  2)
8.6
Group 6A Elements (ns2np4, n  2)
• Oxygen, sulfur, & selenium = nonmetals
Elements = O2 (g), S8 (s) and Se8 (s)
Reactions:
O2 forms oxides (O2-2 O2-) ions
[-2 ions also formed by S, Se, and Te]
nonmetals + Oxygen --> molecular compounds
sufur oxides + water --> acids
SO3(g) + H2O(l)
H2SO4(aq)
• Tellurium & polonium = metalloids
• Polonium = radioactive
8.6
Group 7A Elements (ns2np5, n  2)
8.6
Group 7A Elements (ns2np5, n  2)
X-1
Increasing reactivity
• Nonmetals called halogens
Elements = F2 (g), Cl2 (g), Br2 (l), I2 (s)
• Reactions:
Ions formed: -1 with metals X + 1e+ Nonmetals --> Molecular compounds
+ H2 --> HX X2(g) + H2(g)
2HX(g)
HX + water --> acids
HF = weak acid
others are strong acids
Reactivity with nonmetals:
F2 (explosive) ----I2 (least reactive)
8.6
Group 8A Elements (ns2np6, n  2)
Completely filled ns and np subshells.
Highest ionization energy of all elements.
No tendency to accept extra electrons.
• Noble gases
• Exist as atoms
Un-reactive
• Some Reactions under special conditions:
Xe (and Kr) + O2 (and F2) --> compounds
8.6
Properties of Oxides Across a Period
basic
acidic
Across a period, oxides change from basic to amphoteric to
acidic. Going down a group, the oxides become more basic. 8.6