Download CHEM 120 WEEK 11 LECTURES (INORGANIC WEEK 2) Dr. MD

CHEM 120
WEEK 11 LECTURES
(INORGANIC WEEK 2)
Dr. MD BALA
The Periodic table
Alkali metals
Group 1A
s- block
Valence shell electronic structure =
ns1
Alkali metals (Group 1)
• Soft, metallic solids.
• Name comes from Arabic
word for ashes.
• Produce bright colors when placed in flame.
Li
Na
K
Rb
Cs
Alkali metals (Group 1)
• Have low densities and melting points.
• Also have low ionization energies.
• potential reducing agents due to their capacity to form
stable cations (M+).
• Found only as compounds in nature and not as pure metals.
• Lithium reacts with the nitrogen in the air to give lithium
nitride. Lithium is the only element in this Group to form a
nitride in this way.
6 Li + N2
2 Li3N
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
and reactivity increases down the group.
Alkali metals: Reactions with oxygen
• Alkali metals (except Li) react with oxygen to form peroxides or
superoxides.
• Na and K can react with oxygen to form peroxides.
• Rb and Cs can react with oxygen to form superoxides.
4Li(s) + O2(g)
2Li2O(s)
(oxide, all conditions)
4Na(s) + O2(g)
2Na(s) + O2(g)
2Na2O(s)
Na2O2(s)
(oxide, limited O2)
(peroxide, excess O2)
2K(s) + O2(g)
K(s) + O2(g)
K2O2(s)
KO2(g)
(peroxide, limited O2)
(superoxide, excess O2)
Rb(s) + O2(g)
Cs(s) + O2(g)
RbO2(s)
CsO2(s)
white
white
yellow
(superoxide, all conditions)
(superoxide, all conditions) orange
Peroxides
• Here oxygen has an
oxidation state of -1
• The O-O bond is
very weak
– Decomposition of
peroxides can be
dangerously
exothermic
Superoxides
• Oxygen has oxidation state of -½
• The most active metals (K, Rb,
Cs) form superoxides through
reaction with O2
• React with H2O to form O2
– Source of O2 in self-contained
breathing devices.
Alkali metals: Reaction of KO2
4 KO2 + 4 CO2 +2 H2O → 4 KHCO3 + 3 O2
Purify air in submarines
Emergency breathing apparatus
Uses of alkali metals
• Lithium
– Alloys of Li-Al-Mg for aircraft and space
applications
– Battery anodes
• Sodium
– Heat-transfer medium in
nuclear reactors
– Sodium vapour lamps
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)
(Recall NaCl electrolysis topic)
A pile of salt at Uppington,
South Africa.
Sodium compounds
= Heat
 Try example A and B page 875 – Petrucci 9th Edition
Carbonates
 Li2CO3 is unstable relative to the oxide.
- Used to treat manic depression (1-2 g/day).
 Na2CO3 primarily used to manufacture glass.
- Currently mined from rich U.S. resources, but can
be manufactured by the Solvay process.
- Used in the lab as a primary standard to standardize
acids.
 NaHCO3 baking powder,
- as an antacid to treat acid indigestion and heartburn.
- cleaning and scrubbing.
- Buffers, because it is amphoteric, reacting with both
acids and bases
The Periodic table
Alkaline earth metals
Main Group
s- block
Valence shell electronic structure =
ns2
Alkaline earth metals (Group 2)
• Have higher densities and melting points than alkali
metals.
• Have low ionization energies, but not as low as alkali
metals.
• potential reducing agents due to their capacity to form
stable cations (M2+)
Alkaline earth metals: Reaction with water
• Be does not react with
water, Mg reacts only
with steam, but others
react readily with water.
Mg + 2 H2O → Mg(OH)2 + H2
• Reactivity tends to
increase as go down
group.
Group 2 : The alkaline earth metals
• Principle forms:
– 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 I (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) (M = Be)
• 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:
– Therefore they are covalent rather than ionic solids.
Beryllium Chloride
 Group 2 elements form di-cations (M2+) therefore one can predict that these
elements will form ionic compounds with non-metals.
 The exception is beryllium which forms covalent compounds
Why ?
 Due to the undesirable high charge density on Be2+ (charge density is the charge
per unit volume and the volume of the Be2+ ion is very small).
Decomposition of CaCO3 (lime)
In the lime kiln:
CaCO3 Δ
→
In the lime slaker:
CaO + H2O → Ca(OH)2
CaO + CO2
slaked lime
burnt lime
or
quicklime
Ca(OH)2
Calcium hydroxide
Stalactites and Stalagmites
CO2 + H2O → H3O+ + HCO3Ka = 4.410-7
HCO3- + H2O → H3O+ + CO32Ka = 4.710-11
CaCO3(s) + H2O(l) + CO2(g) → Ca(HCO3)2(aq)
(Limestone)
( Recall the solubility properties – How acids dissolves insoluble CaCO3 )
The Periodic table
Group 13
Valence shell electronic structure =
ns2np1
Group 13
 Contains only metals, apart from boron.
 Boron is also the only element which does not form a stable
trication (B3+) again will have too high a charge density to be
stable.
Why do the other elements form tri-cations (M3+ )?
Soln. √ Because they have the valence electronic configuration
ns2np1 and when three electrons are lost they have a an
electron configuration of the closest noble gas.
1.
 The metals, however, can also form covalent compounds
with non-metals depending on the difference in
electronegativity. AlF3 for instance is ionic but AlCl3 is
covalent.
Nb. A very large electronegativity difference between two atoms is an indication
of a probable ionic bond formation. (EN) ; > 1.7 ionic;
0.4 – 1.7 polar covalent; 0 to 0.4 nonpolar covalent:
The Boron Family
Borax
Boric acid
Try example A and B page 889 – Petrucci 9th Edition
Uses of Group 13 metals
• Aluminium is 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) + 3/2 O2(g) → Al2O3(s) ΔH = -1676 kJ
The Thermite reaction
(used in on-site welding of large objects)
2 Al(s) + Fe2O3(s) → Al2O3(s) + Fe(s)
Uses of Group 13 metals
• Indium
– Makes low melting alloys.
– Low-temperature transistors and
photoconductors.
• Thallium
– Extremely toxic. Few industrial uses.
– Tl2Ba2Ca2Cu3O8+x exhibits superconductivity up
to 125 K.
Oxidation states of Group 13 metals
• Al almost exclusively 3+
• In and Ga both 3+ and 1+
…(but 3+ is favoured)
• Tl both 1+ and 3+
…(but 1+ is favoured)
– Tl = [Xe]4f145d106s26p1
– Tl+ resembles Group 1 (looses 6p1)
– [Xe]4f145d106s2 – the inert pair effect.
(electron pair due to 6s2)
 Small bond and lattice energies associated with large atoms and ions
at the bottom of a group are not sufficiently great to offset the
ionization energies of the ns2 electrons.
Aluminium Halides
Adduct
Lewis acid
Al2Cl6
dimer
Aluminium and Alums
Anodized aluminum
Alum crystals
Electrolysis – half-reaction at the anode
2Al(s) + 3H2O(l)
Al2O3(s) + 6H+ + 6e-
Aluminium hydroxide is also 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)
The Periodic table
Group 14
Valence shell electronic structure =
ns2np2
Group 14
o Properties vary through this group but
• all exhibit +4 oxidation state.
o Tin and Lead are metallic and form ionic bonds
• +2 and +4 oxidation states.
o Germanium and Silicon are semiconductors and metalloids.
• they form covalent bonds.
o Carbon is a nonmetal and forms covalent bonds.
Carbon
Synthetic
diamonds
Carbon nanotubes
The Periodic table
Group 15
Valence shell electronic structure =
ns2np3
Group 15
• Nitrogen and phosphorus are non-metals.
• Arsenic (As) and antimony (Sb) are metalloids and bismuth
(Bi) is a metal.
• They all exhibit the maximum oxidation state (+5) in the
group, and their chemistry is largely covalent.
• Most of the chemistry of these elements is in the +3 or +5
oxidation states although they form gaseous compounds with
hydrogen in the -3 oxidation state:
i.e. ammonia NH3, phosphine PH3, arsine AsH3,
stibine SbH3, and bismuthine BiH3;
Properties of Nitrogen
• It is a colourless, odourless, and tasteless gas composed of N2
molecules.
• It is unreactive because of the strong triple (N≡N) bond.
• Exception: Burning Mg or Li in air (78% nitrogen) forms
nitrides:
• 3Mg(s) + N2(g) → Mg3N2(s); also
• 6Li(s) + N2(g) → 2Li3N(s)
• N3– is a strong Brønsted-Lowry base (forms NH3 in water):
• Mg3N2(s) + 6H2O(l) → 2NH3(aq) + 3Mg(OH)2(s)
• Nitrogen exhibits all formal oxidation states from –3 to +5.
• The most common oxidation states are +5, 0, and –3.
The Periodic Table
Group 16
Valence shell electronic structure =
ns2np4
Group 16
• Oxygen (O), sulfur (S), and selenium (Se) are nonmetals.
• Tellurium (Te) is a metalloid.
• The radioactive polonium (Po) is a metal and has no stable isotopes.
• Their chemistry is dominated by the formation of the -2 oxidation
state but sulfur in particular can exhibit a variety of oxidation states.
Oxygen
• Two allotropes:
– O2, dioxygen
– O3, ozone
• Three anions:
– O2−, oxide
– O22−, peroxide
– O21−, superoxide
• Tends to take electrons
from other elements
(oxidation)
Oxides
Oxides
• Oxygen is the second most electronegative element.
• Oxides are compounds with oxygen in the –2 oxidation state.
Nonmetal oxides are covalent.
• Most metal oxides combine with water to give oxyacids.
• Oxides that react with water to form acids are called acidic anhydrides, or
acidic oxides.
• Anhydride means without water.
• Example:
SO2(g) + H2O(l) → H2SO3(aq)
• Metal oxides are ionic.
• Oxides that react with water to form hydroxides are called basic anhydrides,
or basic oxides.
• Example: BaO in water produces Ba(OH)2.
BaO(s) + H2O(l) → Ba(OH)2(aq)
• Oxides that exhibit both acidic and basic properties are said to be amphoteric
40
(e.g. Cr2O3).
Sulfur
• Weaker oxidizing
agent than oxygen.
• Most stable allotrope
is S8, a ringed
molecule.
Refer to prescribed textbook by Petrucci Chapter 22; page 929
The Periodic Table
Valence shell electronic structure =
ns2np5
Group 17
( Halogens )
Group 17: Halogens
• Prototypical nonmetals
• Name comes from the Greek halos and gennao:
“salt formers”
Group 17: Halogens
• Formation of stable mono-anions is
the most common chemistry displayed
by these elements.
•
•
However, they also display the +7,
+5, +3 and +1 oxidation states, when
bonded to oxygen.
Large, negative electron affinities
 Therefore, tend to oxidize other
elements easily
•
React directly with metals to form
metal halides
Properties and Preparation of the Halogens
• The properties of the halogens vary regularly with their
atomic number.
• Each halogen is the most electronegative element in its row.
• Halogens exist as diatomic molecules.
• In solids and liquids, the molecules are held together by
weak London-dispersion forces
• Iodine has the highest melting point and the strongest
intermolecular forces.
• At room temperature, I2 is a solid, Br2 is a liquid, and Cl2
and F2 are gases.
• Hence, fluorine is very reactive.
• The reduction potential of fluorine is very high.
46
Interhalogen Compounds
• Diatomic molecules containing two different halogens are called
interhalogen compounds.
• Most higher interhalogen compounds have Cl, Br, or I as the
central atom surrounded by 3, 5, or 7 F atoms.
• The larger the halogen, the more interhalogen compounds it
can form.
• The compound ICl3 is unique.
• The large size of the I atom allows it to accommodate the
three Cl atoms.
• No other halogen is large enough to accommodate three Cl
atoms.
• Interhalogen compounds are very reactive; they are powerful
oxidising agents.
47
The Periodic Table
Valence shell electronic structure =
Group 18
( Noble gases )
ns2np6
Group 18: Noble Gases
• The already have a full valence shell and therefore don’t get
themselves involved in too much chemistry!
• Very large ionization energies
• Positive electron affinities
– Therefore, relatively unreactive
• Monatomic gases
Group 18: Noble Gases
• Xe forms three
compounds:
– XeF2
– XeF4 (at right)
– XeF6
• Kr forms only one stable
compound:
– KrF2
• The unstable HArF was
synthesized in 2000.
Noble Gases
• Extremely stable and
unreactive.
• Liquid He (boiling point 4.2
K) used as a coolant.
• Ne used in electric signs.
• Ar used in light bulbs and
as insulating gas between
panes in thermal windows.