Chemistry of the
Main - Group
Elements
CHEM 110/ 2013
Main Group Elements
Groups 1-2, & 13-18 (Old numbering IA - 8A)
1. Physical Properties
2. Uses of Elements
3. Chemical Properties
4. Compounds: Preparation, Reactions &
Uses
Alkali metals
Group 1A
s- block
Periodic Table
Valence shell electronic structure = ns1
Group 1: The Alkali Metals
Chemical Properties
 They have low ionization energies
 Are potential reducing agents due to their capacity to
form stable cations (M+)
 Found only as compounds in nature and not as pure
metals
 Reaction with nitrogen
Li reacts the N2 in the air to form lithium nitride,
Li3N.
6 Li + N2  2 Li3N
Group 1: The Alkali Metals
Reactions with Water
Li
2 M(s) + 2 H2O(l)
Na
K
2 M+(aq) + 2 OH-(aq) + H2(g)
Alkaline
2 Na(s) + 2 H2O(l)
2 Na+(aq) + 2 OH-(aq) + H2(g)
Their reactions with water are famously exothermic &
reactivity increases down the group.
Group 1: The Alkali Metals
Reactions with Oxygen
Lithium - forms only the oxide
4 Li(s) + O2(g)  2 Li2O(s)
Sodium & Potassium - both oxide and peroxide
4 Na(s) + O2(g)  2 Na2O(s) (white oxide, limited O2)
2 Na(s) + O2(g)  Na2O2(s) (peroxide, excess O2)
Potassium - both peroxide and superoxide
2 K(s) + O2(g)  K2O2(s)
(peroxide, limited O2)
K(s) + O2(g)  KO2(g) (yellow superoxide, excess O2)
Rubidium & Cesium - only superoxides
M(s) + O2(g)  RbO2(s) (superoxide, all conditions)
Due to the cation /anion size ratio
Group 1 Compounds
Halides
• NaCl 50 million
tons/year in U.S.
• Preservative, used
on roads, water
softener regeneration.
• KCl from natural brines.
• Plant fertilizers, feed stock.
• Feed stock for other chemicals.
2 NaCl (aq) + 2 H2O(l)  2 NaOH (aq) + H2(g) + Cl2(g)
2 NaCl(l)  2 Na(s) + Cl2(g)
A pile of salt at
Uppington,
South Africa
Group 1 Compounds
Carbonates
Li2CO3 (unstable relative to the oxide)
- Used to treat manic depression (1-2 g/day)
Na2CO3 (manufactured by the Solvay process)
- Manufacture glass
- Primary standard to standardize acids
NaHCO3 (baking powder)
- Antacid to treat acid indigestion & heartburn.
- Cleaning and scrubbing
- In buffers, because it is amphoteric, reacting with
both acids and bases
Periodic Table
Alkaline Earth Metals
Main Group
s- block
Valence shell electronic structure = ns 2
Group 2: Alkaline Earth Metals
Physical & Chemical Properties
 Higher densities & melting points than alkali metals
 Low ionization energies, but not as low as alkali metals
 Potential reducing agents due to their capacity to form
stable cations (M2+)
Group 2: Alkaline Earth Metals
Reaction with water
 Be does not react with water, Mg reacts only with
steam, but others react readily with water
M(s) + 2 H2O  M(OH)2 + H2(g)
 Reactivity tends to increase
as go down group
Group 2: Alkaline Earth Metals
 Principal compounds:
– carbonates, sulfates and silicates
 Oxides and hydroxides only sparingly soluble
– Basic or “alkaline”
 Compounds that do not decompose on heating
– Therefore named “earths”
 Heavier element compounds are more reactive and are
similar to Group 1 (also in other respects)
Beryllium
 Unreactive toward air and water
 BeO does not react with water
– All other Group 2 oxides form hydroxides
MO(s) + H2O(l)  M2+(aq) + 2 OH-(aq)
 BeO dissolves in strongly acidic or basic solutions.
Therefore is an amphoteric oxide
Acid: BeO + 2 HCl + H2O  BeCl2 + 2 H2O
Base: BeO + 2 NaOH + H2O  Na2Be(OH)4
 BeCl2 and BeF2 melts are poor conductors,  they are
covalent rather than ionic solids
Beryllium Chloride
Valence electon configuration: ns 2
 Di-cations (M2+)
 Mg  Ba: form ionic compounds
with non-metals, while
 beryllium - forms covalent compounds
Decomposition of CaCO3
CaCO3 - lime/limestone/marble
Building material; mfr quick lime & slaked lime;
in glass; & as a flux in metallurgical processes
Calcination - In the lime kiln:
CaCO3  CaO + CO2
Δ
burnt lime or quicklime
Slaking - In the lime slaker:
CaO + H2O  Ca(OH)2
slaked lime
Carbonation
Ca(OH)2 + CO2  CaCO3
purified limestone
Calcium hydroxide
Periodic Table
Group 13
Valence shell electronic structure = ns2np1
Group 13
Valence electron configuration: ns 2 np 1
 Boron - nonmetal, Al  Tl - metals
Bonding in Groups 13 elements:
 Boron - does not form a stable ion (B3+/B3-) due to
high charge density - covalent compounds

Aluminium - AlF3 - ionic, but AlCl3 - covalent

Gallium  Thallium - ionic compounds
Group 13: Oxidation States
• Boron - both +3 & -3
• Aluminium - exclusively +3
• Gallium & Indium - both +1 & +3 (but +3 is favoured)
• Thallium - both +1 and +3
(but +1 is favoured)
Inert Pair Effect:
In the +1 oxidation state only the np 1 electron is lost:
ns 2 np 1  ns 2 + 1 eInert Pair - behave as part of the core
Small bond and lattice energies associated with large atoms & ions
at the bottom of a group are not sufficiently great to offset the
ionization energies of the ns2 electrons.
The Boron Family
Borax
Boric acid
B-H-B
bond
Uses of Group 13 metals
Aluminium - most important
 Third most abundant element, 8.3% by mass of
earth’s crust
 Lightweight alloys
 Easily oxidized to Al3+
2 Al(s) + 6 H+(aq)  2 Al3+(aq) + 3 H2(g)
2 Al(s) + 1½ O2(g)  Al2O3(s)
ΔH = -1676 kJ
Thermite Reaction
- used in on-site welding of large
objects
2 Al(s) + Fe2O3(s) → Al2O3(s) + Fe(s)
Aluminium Halides
Adduct
Lewis acid
Al2Cl6
dimer
Aluminium and Alums
Anodized aluminum
Alum crystals
Electrolysis – half-reaction at the anode
2 Al(s) + 3 H2O(l)  Al2O3(s) + 6 H+ + 6e-
Aluminium Hydroxide - Amphoteric
Reactions with acid: 2Al(OH)3(s) + 3H3O+(aq)  [Al(H2O)6]3+(aq)
Reactions with base: 2Al(OH)3(s) + OH-(aq)  [Al(OH)4]-(aq)
Periodic Table
Group 14
Valence shell electronic structure = ns2np2
Group 14
Valence electron configuration: ns 2 np 2
Properties vary through this group
 Carbon - nonmetal & forms covalent bonds, with
oxidation states of +4 & -4
 Silicon - mostly nonmetallic, forms covalent bonds,
+4 oxidation state - are semiconductors
 Germanium - metalloid & forms covalent bonds, +4
oxidation state - are semiconductors eg of bond
 Tin & Lead - metals & form ionic bonds, with
oxidation states of +2 and +4 eg of bond
Periodic Table
Group 15
Valence shell electronic structure = ns2np3
Group 15
Valence electron configuration: ns 2 np 3
Metallic Character:
• Nitrogen & Phosphorus - nonmetals
• Arsenic & Antimony - metalloids
• Bismuth - metal
Oxidation states & Bonding Character:
• All exhibit a max oxidation state of +5, and covalent
bonding.
• Most occur in the +3/+5 oxidation states but the
nonmetal can exist in the -3 oxidation state
Nitrogen Oxides
Nitrogen oxides range in oxidation state from +1 to +5
Chemical Oxidation
Formula
State
Use
N2 O
+1
‘Laughing gas’ - anaesthetic in dentistry
NO
+2
Biologically important
- Long term memory
- Regulation of blood pressure
N2 O 3
+3
Acid anhydride of HNO2
NO2
+4
N2O 4
+4
N2O 5
+5
Acid anhydride of HNO3
Nitrogen Oxides
Preparation
Nitrogen(II) Oxide
Common name: Nitric oxide




N O
Colourless gas
Toxic
Paramagnetic - unpaired electron  reactive
Product of internal combustion engines
Synthesis
3Cu(s) + 8HNO3(aq)  3Cu(NO3)2(aq) + 4H2O(l) + 2NO(g)
Dilute
Key Reaction
2NO(g) + O2(g)  2NO2(g)
Colourless
Brown
NO oxidises in the air to form NO2,
a main component of
photochemical SMOG.
Nitrogen(III) Compounds
N2O3
 Anhydride of nitrous acid
 Pale blue solid  bright blue liquid
O
N O
N2O3 + H2O(l)  2HNO2
N
HNO2
 Exists in aqueous solutions
H O
 Thermally unstable
Δ
3HNO2(aq)  HNO3(aq) + NO(g) + H2O(l)
O
 Synthesis - ‘milkshake’ reaction
Ba(NO2)2(s) + H2SO4(aq)  2HNO2(aq) + BaSO4(s)
Nitrogen(V) Compounds
N2O5
 Anhydride of nitric acid
N2O5 + H2O(l)  2HNO3
HNO3




Colourless, pungent, syrupy liquid
Available as a 55-68 %(m/V) solution
Industrial preparation - Ostwald Process
Decomposes on exposure to light
4 HNO3(aq)  4 NO2(g) + 2 H2O(l) + O2(g)
 Uses:
• Strong oxdizing agent
• Preparation of dyes, fertilizers, drugs & explosives
• Metallurgy
• Reprocessing spent nuclear fuel
• Aqua Regia (33% HNO3 & 67% HCl) - dissolves gold
Phosphorus(V) Compounds
P4O10
Acid anhydride
P4O10(s) + 6H2O(l)  4H3PO4(aq)
Dessicant – dehydrating agent
H3PO4
 Orthophosphoric acid
 Concentrated is 98.3 %(m/V)
 Dehydrate  pyrophosphoric acid (H4P2O7) 
triphosphoric acid (H5P3O10)
O
H O P O H
O
H
Orthophosphoric
Acid
H2O
O
O
H O P O P O H
O
O
H
H
Pyrophosphoric
Acid
H2O
O
O
O
H O P O P O P O H
O
O
O
H
H
H
Triphosphoric
Acid
Phosphorus in Living Organisms
adenosine diphosphate
adenosine triphosphate
Phosphate Fertilizers
• Fluorapatite , Ca3(PO4)2.CaF2 mineral
(NB: this poses a fluoride waste problem)
• Ca3(PO4)2 has a low solubility  poor uptake by plant 
react it with H2SO4 to form (soluble) superphosphate and
gypsum
Ca3(PO4)2(s) + 2H2SO4(aq)  Ca(H2PO4)2(s) + 2CaSO4(s)
superphosphate
gypsum
Periodic Table
Group 16
Valence shell electronic structure = ns2np4
Group 16
Valence electron configuration: ns 2 np 4
Metallic Character:
 Oxygen, sulfur, and selenium - nonmetals.
 Tellurium - metalloid.
 Polonium - radioactive metal with no stable isotopes.
Oxidation States:
 Oxygen: superoxide (-½); peroxide (-1), oxide(-2)
 Sulfur: sulfide (-2), +4 and +6
 Tellurium & Po: +4 and +6
Sulfur
Pyrite- iron disulfide (FeS2)
Sulfur crystals
Sulfur(IV) Oxide
Chemical formula: SO2
 Colourless gas, with a pungent smell
 Oxidation: Dry SO2 can acts as
reducing agent. It is used as a food
preservative by slowing down the oxidation of food by
O2 in air.
SO2(g) + PbO2(s)  PbSO4(s)
 Reduction: With a strong reducing agent, eg Mg
SO2(g) + 2Mg(s)  2MgO(s) + S(s)
 Dissolves in water to form H2SO3 - thermally unstable
SO2(g) + H2O(l)  H2SO3(aq)
 Produced by the burning of sulfur:
S(s) + O2(g)  SO2(g)
Sulfur(VI) Compounds
SO3
 Volatile white solid, low bp 45 oC
H2SO4
 Dense, colourless oily liquid
 Oxidizing agent: Concentrated H2SO4 is a:
•
•
mild oxidizing agent when cold
fairly strong ox. agent when hot
 hot H2SO4 is needed to dissolve copper:
Cu(s) + 2H2SO4(aq)  CuSO4(aq) + SO2(g) + 2H2O(l)
 Dehydrating agent: Conc. H2SO4 is a strong dehydr. agent
C12H22O11(s)  12 C(s) + 11H2O(g) Exothermic
 Dilute H2SO4
•
•
Reacts with metals to produce H2(g) - REDOX
Dissolves carbonates to producs CO2(g) – Acid-Base
Periodic Table
Group 17
(Halogens)
Valence shell electronic structure = ns2np5
Group 17: Halogens
Chemical & Physical Properties
 Nonmetals
 Diatomic molecules symbolized by X2.
 Oxidation States:
• Fluorine only -1
• Chlorine  Astatine: -1, and +1, +3, +5 & +7 when
bonded with oxygen
 Fluorine:
• Most electronegative element.
• Forms strong bonds (both ionic and covalent)
 Large, negative electron affinities  tend to oxidize
other elements easily
 React directly with metals to form metal halides (MX)
Periodic Table
Group 18
(Noble gases)
Valence shell electronic structure = ns2np6
Group 18: Noble Gases
Valence electron configuration: ns 2 np 6in summary
 Noble - very unreactive due to full valence shell
 Monatomic gases
 Very large ionization energies & positive electron
affinities  neither want to accept nor donate electrons
Group 18: Noble Gases-oxides
Noble Gas Compounds (F or O)
 Xenon - forms three compounds:
+2: XeF2
+4: XeF4(), XeOF2
+6: XeF6, XeO3
+8: XeO4, H4XeO6
 Krypton - forms only one stable
compound:
+2: KrF2
 Argon - The unstable HArF was
synthesized in 2000.