Unit 3 - Inorganic Chemistry

Report
Periodicity
Period 3 - elements
Can you write one fact about each of the following
elements?
 Na







Mg
Al
Si
P
S
Cl
Ar
Candidates should be able to:
•describe qualitatively (and indicate the periodicity in) the
variations in atomic radius, ionic radius, melting point and
electrical conductivity of the elements (see the Data Booklet)
•explain qualitatively the variation in atomic radius and ionic
radius
•interpret the variation in melting point and in electrical
conductivity in terms of the presence of simple molecular, giant
molecular or metallic bonding in the elements
•explain the variation in first ionisation energy.
Na
[Ne] 3s1
Mg
[Ne] 3s2
Al
[Ne] 3s2 3px1
Si
P
S
Cl
Ar
[Ne] 3s2 3px1 3py1
[Ne] 3s2 3px1 3py1 3pz1
[Ne] 3s2 3px2 3py1 3pz1
[Ne] 3s2 3px2 3py2 3pz1
[Ne] 3s2 3px2 3py2 3pz2
Atomic radius
0.25
Atomic radius (nm)
0.2
0.15
0.1
0.05
0
Na
Mg
Al
Si
P
Element
S
Cl
Ar
Atomic radius – what about argon?
www.chemguide.co.uk/atoms/properties/atradius.html
What is atomic radius?
www.rjclarkson.demon.co.uk/found/period3_atomrad.gif
Explaining the trend
An atomic radius is a measure of the distance from the
nucleus to the bonding pair of electrons.
From sodium to chlorine, the bonding electrons are all in
the 3rd shell being screened by the electrons in the first
and second levels, i.e. the screening remains fairly
constant.
The increasing nuclear charge as you go across the period
pulls the bonding electrons more tightly towards it.
You have to ignore the noble gas at the end of
each period. Because neon and argon don't form
bonds, you can only measure their van der Waals
radius - a case where the atom is pretty well
"unsquashed". All the other atoms are being
measured where their atomic radius is being
lessened by strong attractions. You aren't
comparing like with like if you include the noble
gases.
Leaving the noble gases out,
atoms get smaller as you go
across a period.
IONIC RADIUS
Ions aren't the same size as the atoms they come from. Compare the sizes of sodium
and chloride ions with the sizes of sodium and chlorine atoms.
Positive ions
Positive ions are smaller than the atoms they come from. Sodium is 2,8,1; Na+
is 2,8. You've lost a whole layer of electrons, and the remaining 10 electrons
are being pulled in by the full force of 11 protons.
Negative ions
Negative ions are bigger than the atoms they come from. Chlorine is 2,8,7; Clis 2,8,8. Although the electrons are still all in the 3-level, the extra repulsion
produced by the incoming electron causes the atom to expand. There are still
only 17 protons, but they are now having to hold 18 electrons.
Pauling electronegativity value
Electronegativity
3.5
3
2.5
2
1.5
1
0.5
0
Na
Mg
Al
Si
P
Element
S
Cl
Ar
First ionisation energy (kJ mol-1)
First ionisation energy
1600
1400
1200
1000
800
600
400
200
0
Na
Mg
Al
Si
P
Element
S
Cl
Ar
Electronegativity
Bonding, Structure and Properties
Structure
Type of
element
Bonding
Formula
Type of
force
broken on
melting/
boiling
Does the
element
conduct
electricity
?
Na
Mg
Al
Si
P
S
Cl
Ar
Giant
metallic
Metal
Giant
metallic
Metal
Giant
metallic
Metal
Giant
Simple
Simple
Simple
Monatomic
covalent
molecule
molecule
molecule
Non-metal Non-metal Non-metal Non-metal Non-metal
Metallic
Metallic
Metallic
Covalent
Metallic
bond
Metallic
bond
Metallic
bond
Covalent
bond
Yes
Yes
Yes
No
Covalent
P4
vdW
Covalent
S8
vdW
Covalent
Cl2
vdW
Ar
vdW
No
No
No
No
Structure and Properties
Electrical conductivity
Melting and boiling points
3000
2500
2000
1500
1000
500
0
Periodicity
Period 3 - oxides
Starter activity
Complete these sketches to show how these
properties change as you go along Period 3 from
Na to Ar.
Candidates should be able to:
describe the reactions, if any, of the elements with
oxygen to give Na2O, MgO, Al2O3, P4O10, SO2 and SO3.
state and explain the variation in oxidation number of
the oxides.
describe the reactions of the oxides with water.
describe and explain the acid/base behaviour of oxides
and hydroxides, including, where relevant, amphoteric
behaviour in reaction with NaOH and acids.
Reactions with oxygen
Group number
Element in
Period 3
Nuclear charge
[Ne] electronic
configuration
Trend in Atomic
radius
Trend in 1st
ionisation energy
Trend in
electronegativity
Formula of
oxide/s
Oxidation state
1
2
Na
Mg
11+
3s1
12+
3s2
3
Al
4
Si
5
P
6
S
13+
14+
15+
16+
3s2 3p1 3s2 3p2 3s2 3p3 3s2 3p4
decreases
increases
Na2O
MgO
Al2O3
SiO2
P4O10
+1
+2
+3
+4
+5
increases
SO2/
SO3
+4/+6
Reactions with oxygen
Sodium
Sodium burns in oxygen with an orange flame to
produce the white solid sodium oxide.
Magnesium
Magnesium burns in oxygen with an intense white flame
to give white solid magnesium oxide.
Reactions with oxygen
Aluminium
Aluminium will burn in oxygen if it is powdered, otherwise
the strong oxide layer on the aluminium tends to inhibit
the reaction. If you sprinkle aluminium powder into a
Bunsen flame, you get white sparkles. White aluminium
oxide is formed.
Silicon
Silicon will burn in oxygen if heated strongly enough.
Silicon dioxide is produced.
Reactions with oxygen
Phosphorus
White phosphorus catches fire spontaneously in air,
burning with a white flame and producing clouds of
white smoke - a mixture of phosphorus(III) oxide and
phosphorus(V) oxide. The proportions of these depend
on the amount of oxygen available. In an excess of
oxygen, the product will be almost entirely
phosphorus(V) oxide:
Reactions with oxygen
Sulphur
Sulphur burns in air or oxygen on gentle heating with a
pale blue flame. It produces colourless sulphur dioxide
gas.
In an excess of pure oxygen, some SO3 is also formed.
This utilises the highest oxidation state of sulphur.
Reactions with oxygen
Which is which?
Melting points of oxides
Na2O
MgO
Al2O3
SiO2
P4O10
SO2
Tm/K
1548
3125
2345
1883
573
200
Bonding
Ionic
Ionic
Ionic
Covalent
Covalent
Covalent
Giant
lattice
Giant
lattice
Giant
lattice
Giant
lattice
Structure
Simple
Simple
molecular molecular
Giant ionic and covalent solids contain only strong bonds and have high melting
points. Simple molecules have weak vdW forces between molecules and have
low melting points.
Acid/base properties of the oxides
Reaction with water
pH
Na2O + H2O  2Na+ + 2OH-
14
MgO + H2O  Mg2+ + 2OH-
9
Al2O3 insoluble – no reaction
7
SiO2 insoluble – no reaction
7
P4O10 + 6H2O  4H3PO4
0
SO2
+
H2O 
H2SO3
3
SO3
+
H2O 
H2SO4
0
Acid/base behaviour of aluminium oxide
BASE:
Al2O3 + 3H2SO4 →
Al2(SO4)3 + 3H2O
ACID:
Al2O3 + 2NaOH + 3H2O →
2NaAl(OH)4
Periodicity
Period 3 - chlorides
Candidates should be able to:
•describe the reactions, if any, of the elements with
chlorine to give NaCl, MgCl2, Al2Cl6, SiCl4, and PCl5.
•describe and explain the reactions of the chlorides
with water.
Reaction with chlorine
Element
Na
Mg
Description of
reaction with Very vigorous
Vigorous
chlorine
Formula of
NaCl
MgCl2
chloride/s
Oxidation
state of
+1
+2
period 3
element
State of
chloride at
Solid
Solid
r.t.p.
b.pt. of
1465
1418
chloride (oC)
Structure of
Giant lattice
chloride
Bonding in
Ionic
chloride
Al
Si
P
Vigorous
Slow
Slow
Al2Cl6
SiCl4
PCl3 / PCl5
+3
+4
+3 / +5
Solid
Liquid
liquid / solid
423
57
74 / 164
Simple molecular
Covalent
Structure of Al2Cl6
Acid/base properties of the chlorides
Reaction with water
NaCl(s)  Na+(aq) + Cl-(aq)
MgCl2(s)  Mg2+(aq) + 2Cl-(aq)
pH
7
6/7
Al2Cl6(s) + 12H2O(l)  2[Al(H2O)6]3+ 6Cl-(aq)
3
SiCl4(l) + 2H2O(l)  SiO2(s) + 4HCl(g)
0
PCl5(l) + 4H2O(l)  H3PO4(aq) + 5HCl(g)
0
Periodicity
Group 7 - Halogens
Starter activity
Can you complete task 1 – IGCSE revision on the
halogens
Candidates should be able to:
•describe the trends in volatility and colour of chlorine,
bromine and iodine.
•interpret the volatility of the elements in terms of van der
Waals’ forces.
•describe the relative reactivity of the elements as oxidising
agents.
•describe and explain the reactions of the elements with
hydrogen.
•describe and explain the relative thermal stabilities of the
hydrides.
•interpret these relative stabilities in terms of bond
energies.
The Halogens
Ionic
Covalent
The Halogens
Trend in colour and physical state
Like dissolves like!
Periodicity
Group 7 – the halides
Starter activity
Can you complete the table ‘Reducing power of the
halides’?
Trend in reducing ability
NaX
NaF
NaCl
NaBr
NaI
Observations
steamy fumes
steamy fumes
steamy fumes
colourless gas
Products
HF
HCl
HBr
SO2
brown fumes
steamy fumes
colourless gas
Br2
HI
SO2
yellow solid
S
smell of bad eggs
H2S
Grey solid, purple I2
fumes
Type of reaction
acid-base (F- acting as a base)
acid-base (Cl- acting as a base)
acid-base (Br- acting as a base)
redox (reduction product of
H2SO4)
redox (oxidation product of Br-)
acid-base (I- acting as a base)
redox (reduction product of
H2SO4)
redox (reduction product of
H2SO4)
redox (reduction product of
H2SO4)
redox (oxidation product of I-)
Candidates should be able to:
•describe and explain the reactions of halide ions with
aqueous silver ions followed by aqueous ammonia
concentrated sulphuric acid.
•describe and interpret in terms of changes of oxidation number
the reaction of chlorine with cold, and with hot, aqueous sodium
hydroxide.
•explain the use of chlorine in water purification.
•recognise the industrial importance and environmental
significance of the halogens and their compounds, (e.g. for
bleaches; PVC; halogenated hydrocarbons as solvents, refrigerants
and in aerosols).
Summary of reducing power
Fluoride
Relative
reducing
power
Chloride
Bromide
Iodide
Testing for halide ions
ion present
observation
F-
No precipitate
Cl-
White precipitate
Br-
Cream precipitate
I-
Yellow precipitate
Silver fluoride is soluble, and
so you don't get a precipitate.
Confirmatory tests
original precipitate
E.g.
observation
AgCl
Soluble in dilute NH3(aq)
AgBr
Soluble in concentrated NH3(aq)
AgI
Insoluble in concentrated NH3(aq)
AgCl(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) +
Cl-(aq)
Reactions of chlorine with NaOH
Cl2(g) + 2NaOH(aq) → NaClO(aq) + NaCl(aq) + H2O(l)
0
+1
-1
3Cl2(g) + 6NaOH(aq) → NaClO3(aq) + 5NaCl(aq) + 3H2O(l)
Uses of halogens and their compounds
bleach
refrigerants
PVC
solvents
Aerosol propellants
Periodicity
Nitrogen and Sulphur
Starter activity
Can you use information in your textbooks to
complete task 9?
Candidates should be able to:
•explain the lack of reactivity of nitrogen.
•describe the
oformation, and structure of, the ammonium ion
othe displacement of ammonia from its salts.
•understand the environmental consequences of the uncontrolled use of
nitrate fertilisers.
•understand and explain the occurrence, and catalytic removal, of oxides
of nitrogen.
•explain why atmospheric oxides of nitrogen are pollutants, including their
catalytic role in the oxidation of atmospheric sulphur dioxide.
•describe the formation of atmospheric sulphur dioxide from the
combustion of sulphur contaminated carbonaceous fuels.
•state the role of sulphur dioxide in the formation of acid-rain and
describe the main environmental consequences of acid-rain.
•understand the industrial importance of sulphuric acid.
•describe the use of sulphur dioxide in food preservation.
Gases in the air
Structure of nitrogen
Bond dissociation enthalpy = +946 kJ/mol
Nitrogen
Enzymes
High energy
The Haber Process
The ammonium ion
Ammonium NH4+
Add sodium
hydroxide solution Ammonia gas is
to a solution of the
given off.
substance and
gently heat.
Fertiliser
n. Any of a large number of natural and synthetic
materials, including manure and nitrogen, phosphorus,
and potassium compounds, spread on or worked into soil
to increase its capacity to support plant growth.
Fertilisers
Ammonium nitrate
Ammonium phosphate
Potassium chloride
Nitric acid
NH3(g) + 2O2(g)  HNO3(l) + H2O(l)
Pt/Rh catalyst 900oC
The environment and fertilisers
fertilisers applied to farm land
if
too much used, at the wrong time of year, during
wet weather,
excess
washed into
rivers and lakes
causes
Excessive growth of aquatic plants.
The bacteria which live on dead plants
thrive and use up the oxygen in the water.
The lack of oxygen causes death of fish.
This is called eutrophication.
excess
contaminates
underground
drinking water
supplies
causes
Harm to infants - called
‘blue baby’ syndrome
Natural sources of SO2 and NOx
Sulphur dioxide - SO2
Used as a food preservative:
Inhibits growth of moulds,
yeasts and aerobic bacteria.
Acts as a reducing agent
and retards the oxidation of
foodstuffs.
Acid rain
SO2(g) + NO2(g)  SO3(g) + NO(g)
SO3(g) + H2O(l)  H2SO4(l)
Uses of sulphuric acid
Detergents
Paints , pigments
and dyestuffs
Fertilisers
AS Chemistry
Periodicity
Group 2
Candidates should be able to interpret, and make
predictions from, the trends in physical and chemical
properties of the elements and their compounds.
Trends in Atomic Radius
The radius of an atom is governed by:
the distance of the outer electrons from the nucleus
the nuclear charge, and
the amount of shielding
Trends in First Ionisation Energy
Trends in First Ionisation Energy
Ionisation energy is also governed by:
the charge on the nucleus,
the amount of shielding by the inner electrons,
the distance between the outer electrons and the
nucleus.
Increased nuclear charge ‘off-set’ by increased shielding.
Outermost electron increasingly distant from pull of
nucleus, less energy needed to remove it.
Trends in Electronegativity
Trends in Electronegativity
Trends in Electronegativity
Trend in Electronegativity
As the metal atoms get bigger, any bonding
pair gets further and further away from the
metal nucleus, and so is less strongly attracted
towards it.
Increasing nuclear charge ‘off-set’ by more
shielding.
In other words, as you go down the Group,
the elements become less electronegative.
Trend in Melting Point
Trend in Melting point
Going down Group 2:
•the number of delocalised electrons remains the same ...
•the charge on each metal cation stays the same at 2+,
but ...
•the ionic radius increases ...
•so the attraction between the delocalised electrons and
the metal cations decreases.
AS Chemistry
Periodicity
Group 2 – chemical
properties
Candidates should be able to:
•describe the reactions of the elements with oxygen
and water.
•describe the behaviour of the oxides with water.
•describe the thermal decomposition of the nitrates
and carbonates.
•interpret, and make predictions from, the trends in
chemical properties of the elements and their
compounds.
•explain the use of magnesium oxide as a refractory
lining material and calcium carbonate as a building
material.
•describe the use of lime in agriculture.

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