Module 9: Chemical from the Earth

Formulae of elements and compounds

Chemical equations
copper II oxide + carbon ---> carbon monoxide + copper
CuO(s) + C(s)  --->  CO(g) + Cu(s)
iron III oxide + carbon monoxide ---- iron + carbon dioxide
Fe₂O₃(s) + 3CO(g) ---> 2Fe(s) + 3CO₂(g)
aluminium oxide ---> aluminium + oxygen
2Al₂O₃(s) ---> 2Al(s) 3O₂(g)
sodium chloride + water ---> sodium hydroxide + hydrogen + chlorine
2NaCl(aq) + 2H₂O(l) ---> 2NaOH(aq) + H₂(g) +Cl₂(g)
sodium + water ---> sodium hydroxide + hydrogen
2Na(s) + 2H₂O(l) ---> 2NaOH(aq) + H₂(g)
nitrogen + hydrogen ---> ammonia
N₂(g) + H₂(g) ---> 2NH₃(g)
sulfuric acid + ammonia ---> ammonium sulfate
H₂SO₄(aq) + 2NH₃(aq) ---> (NH₄)₂SO₄(aq)

Ionic equations
aluminium ion + electrons ---> aluminium atoms
Al³⁺ + 3e^- ---> Al
copper  atoms - electrons ---> copper ion
Cu - 2e^- ---> Cu²⁺

9.01 Uses of iron copper and aluminium

Task 9.01

9.02 Metal ores
An ore is a metal compound mixed with rock which is found in the ground.  
See illustrations:  Examples include:  
haematite which is mostly iron III oxide Fe₂O₃
bauxite which is mostly aluminium oxide Al₂O₃
malachite which is mostly copper carbonate CuCO₃

9.03 Reduction
Reduction is a chemical reaction when there is a loss of oxygen from a substance.  e.g.
copper oxide --> copper + oxygen
Copper has lost oxygen so copper is reduced.

9.04 Oxidation reduction and electrons
Oxidation is a gain of oxygen. e.g.
copper + oxygen --> copper oxide
Copper has gained oxygen so copper is oxidised.
Oxidation and reduction can be described as loss or gain of electrons.  Remember:
Oxidation
Is
Loss of electrons
Reduction
Is
Gain of electrons
Oxidation  Mg - 2e^- --> Mg²⁺
Reduction Cu²⁺ +2e^- --> Cu

9.05 The extraction of metals
An ore is a material found in the ground which contains a metal.  An ore is often a metal oxide.  When a metal is extracted its ore is reduced.    Metal loses oxygen from its oxide. E.g.
iron oxide + carbon monoxide --> iron + carbon dioxide
(iron ore)                                     (the ore is reduced to iron)

9.06 Order of reactivity of metals and extraction
The method of extraction of a metal depends on the reactivity of the metal.  The reactivity series is a list of metals (and some other elements) in order of their reactivity:
potassium
sodium
lithium
calcium
magnesium
aluminium
(carbon)
zinc
iron
tine
lead
(hydrogen)
copper
silver
gold
platinum

Unreactive metals near the bottom, like gold, can be found in the ground as lumps of metal.
 
More reactive metals like iron can be reduced with carbon because carbon is more reactive than iron.  Carbon is more reactive than iron as carbon is above iron in the reactivity series.

Very reactive metals like aluminium can only be extracted by electrolysis. Carbon is not used as it is less reactive than aluminium.  Carbon is below aluminium in the series.

9.07 The extraction of iron

A blast furnace is used in the process of extracting iron. The raw materials iron ore, coke and limestone are put in at the top.  Hot air is blasted into this furnace at the bottom making the coke (carbon) burn faster and the temperature rises to about 1500º. When the coke burns, carbon dioxide is produced.
C + O₂ ---> CO₂
CO₂ reacts with the unburnt coke to form carbon monoxide CO
CO₂ + C ---> 2CO
Iron oxide Fe₂O₃  in the ore is reduced to iron by the reaction with the carbon monoxide.
3CO + Fe₂O₃ ---> 3CO₂ + 2Fe
Molten iron is a dense liquid, so runs to the bottom of the furnace and is tapped off.
Limestone CaCO₃ helps remove impurities during the extraction by forming calcium oxide CaO.
CaCO₃ ---> CaO + CO₂
The rock impurities silicon dioxide SiO₂ are then removed by the following reaction.
CaO + SiO₂ ---> CaSiO₃
CaSiO₃ is known as slag and can be used in making cement and road building.

9.08 Electrolysis
This is a chemical reaction caused by electricity.
It needs a liquid chemical that conduct electricity (an electrolyte).
It needs conductors to carry the electricity into and out of the electrolyte (electrodes)
A positive electrode is called an anode.  A negative electrode is called a cathode.

9.09 Explaining electrolysis

 
Electrolytes like lead bromide contain ions.  These ions have a charge and they are attracted to the charged electrodes because opposite charges attract. 
Negative bromide ions go to the positive anode.  The bromide ions lose electrons and turn into bromine atoms.
2Br^- - 2e^- --> Br₂
Positive lead ions go to the negative cathode.  The lead ions gain electrons and are reduced to lead atoms.
Pb²⁺ +2e^- --> Pb

9.10 The extraction of aluminium

Aluminium is found in the ground in an ore called bauxite.  Bauxite is aluminium oxide  (Al₂O₃) with iron oxide impurities. After purification aluminium oxide is mixed with cryolite to lower the melting point from 2000º to 1000º, which saves money. This mixture is heated and the molten liquid used as the electrolyte. Both electrodes are made of graphite (carbon).  The anode (+ve) is graphite and the cathode (-ve) is a graphite lining to a steel case.

At cathode - positive aluminium ions attracted, gain electrons and become atoms.
Al³⁺ + 3e^- ---> Al

At anode - negative oxide ions attracted, lose electrons and become atoms.
2O^2- ----> O₂ + 4e^-

9.11 The purification of copper

Very pure copper is needed for copper wires. Electrolysis is needed to purify copper. The anode is a mass of impure copper and the cathode is pure copper. The electrolyte is sulphuric acid. The impurities drop at the anode as sludge during electrolysis.
At anode
Cu ---> Cu²⁺ + 2e^-
At cathode
Cu²⁺ + 2e^-  ---> Cu

9.12 Transition metals in the periodic table


9.13 The physical properties transition metals

Transition metals have high melting and boiling points.  They have high densities, that is, they are very heavy.  They are good conductors of heat and electricity
Task 9.13

9.14 Transition metals and coloured compounds
All compounds of transition metals are coloured.  Some examples are:
copper sulphate - blue, iron III oxide - brown, cobalt chloride - pink.

9.15 Uses of transition metals as catalysts

Transition metal or compound	Catalyst use
manganese (IV) oxide	decomposing hydrogen peroxide
platinum	catalytic converter in car exhaust
nickel	making margarine
vanadium (V) oxide	making sulphuric acid
platinum with rhodium	making nitric acid
iron	making ammonia

9.16 Names of alkali metals
The alkali metals are
Lithium Li
sodium Na
potassium K

9.17 Physical properties of alkali metals
Alkali metals have low melting points and are very soft compared to other metals.
Task 9.17

9.18 The reactions of lithium, sodium and potassium with water

Lithium is stored under oil to stop it reacting with oxygen.  Lithium is soft and can just be cut with a knife.  When cut its surface is shiny but is slowly changes when exposed to the air.  On water lithium floats because it has a lower density than water.  Lithium moves around slowly on the surface fizzing because a gas called hydrogen is formed.  The solution left behind turn universal indicator blue.  Lithium hydroxide is formed which is an alkali.  The word equation is  lithium + water --> lithium hydroxide + hydrogen
The balanced equation is:
2Li(s) + 2H₂O(l) --> 2LiOH(aq) + H₂(g)

Sodium is stored under oil to stop it reacting with oxygen.  Sodium is soft and can easilybe cut with a knife.  When cut its surface is shiny but is quickly changes when exposed to the air.  On water sodium floats because it has a lower density than water.  Lithium moves around quickly on the surface fizzing because a gas called hydrogen is formed.  The sodium may burn with a yellow flame. The solution left behind turn universal indicator blue.  Sodium hydroxide is formed which is an alkali.  The word equation is:
sodium + water --> sodium hydroxide + hydrogen
The balanced equation is:
2Na(s) + 2H₂O(l) --> 2NaOH(aq) + H₂(g)  
See a large piece of sodium dropped in water.

Potassium is stored under oil to stop it reacting with oxygen.  Potassium is very soft and can very easily be cut with a knife.  When cut its surface is shiny but is very quickly changes when exposed to the air.  On water potassium floats because it has a lower density than water.  Potassium moves around vigorously on the surface fizzing because a gas called hydrogen is formed.  The potassium burns with a lilac flame.  The solution left behind turn universal indicator blue.  Potassium hydroxide is formed which is an alkali.  The word equation is: 
potassium + water --> potassium hydroxide + hydrogen
The balanced equation is:
2K(s) + 2H₂O(l) --> 2KOH(aq) + H₂(g)
Task 9.18

9.19 The pattern in reactivity of the alkali metals
Alkali metals get more reactive with water as you go down the group. e.g.
Lithium at the top, fizzes in water and is less reactive than potassium below it in the group, which has a violent reaction leading to it bursting into a lilac flame.  Lower still in the group   rubidium and caesium are even more violent.

9.20 The electrolysis of concentrated aqueous sodium chloride (rock salt)
When the electrolysis of aqueous sodium chloride takes place, hydrogen and chlorine are given off as gases and sodium hydroxide is left. Aqueous sodium chloride contains hydrogen ions H⁺ and hydroxide ions OH^- (from the water) and sodium ions Na⁺ and chloride ions Cl^-. The positive sodium and hydrogen ions go to the cathode and the negative chloride and hydroxide ions go to the anode. Hydrogen is formed at the cathode and chlorine is formed at the anode. 


9.21 The uses of sodium chloride, hydrogen and sodium hydroxide

9.22 understand that the crystalline nature of igneous rocks and the fact that they do not contain fossils are evidence for their formation from hot, molten magma
Igneous rocks are formed when liquid rock or magma pushes up into the crust and cools.  It is made up of crystals  It does not contain any fossils. Any living thing falling into the molten rock would be burnt and leave no trace. Igneous rock forms as magma cools slowly under the surface e.g. Granite.  Magma reaching the surface through a volcano cools quickly e.g. basalt.  See some igneous rocks.

9.23 understand that crystal size in igneous rocks depends on the rate of cooling
Igneous rock forms as magma cools slowly under the surface e.g. Granite.  Magma reaching the surface through a volcano cools quickly e.g. basalt.  Igneous rocks which cool slowly have large crystals e.g. granite but rock forming quickly has smaller crystals e.g. balsalt.

9.24 understand that the presence of fossils in a rock is evidence that it has been formed from sediments

Archioptrix was the first fossil bird.  
It was found in sedimentary rock.  The bird lay in a sediment when it died.  Layers of sediment formed on top.  These layers turned into sedimentary rock.

9.25 The age of sedimentary rocks
Sedimentary rock forms in layers. New sediments are layered on top of old ones so the age of sedimentary rock increases with depth.  Species become extinct and new species take their place so fossils of different species are of known age can be used to date the rock.


9.26 Metamorphic rocks
Metamorphic rocks are formed by the action of heat and pressure over long periods of time on rocks that are already formed. Earth movements push all types of rock underground, here they are compressed and heated and the mineral structure and texture may change. For example marble is formed from limestone.  See the rock cycle (needs powerpoint)

9.27 The composition of metamorphic rocks
Metamorphic rock has the same chemicals in it as the rock that it was made from.  Limestone is made of calcium carbonate.  Marble is a metamorphic rock made from limestone.  It is also made of calcium carbonate.  Mudstone in a sedimentary rock which turns into the metamorphic rock called slate.  Both mudstone and slate contain the same clay like minerals.  See some metamorphic rocks

9.28 The current composition of the atmosphere
The atmosphere is made up of the following gases, Nitrogen 78%, Oxygen 21%, Argon 1%, Carbon dioxide 0.03%, water - variable

9.29 The early atmosphere and volcanic activity
The primary atmosphere of the Earth was hydrogen and helium. These light gases were slowly lost. They were replaced by a secondary atmosphere produced by the action of volcanoes.

9.30 The composition of the early atmosphere
The Earth's secondary atmosphere was made up of some left over hydrogen, carbon dioxide, water vapour, nitrogen, carbon monoxide, sulphur dioxide ammonia and methane.

9.31 The origins of the oceans and their effect on carbon dioxide
As the Earth cooled to below 100^oC oceans were formed when water vapour condensed and formed liquid water. Oceans are reservoirs for carbon dioxide because they can store the gas when it dissolves in  them.  The new oceans dissolved a great deal of the carbon dioxide in the atmosphere.  The oceans still play a part in keeping carbon dioxide levels constant.  If there is a lot of carbon dioxide in the air then more can dissolve.  If there is less carbon dioxide in the air then some comes out of solution back into the air.

9.32 Primitive plants and the release of oxygen
As the temperature of the Earth cooled simple green plants evolved in the oceans to use the carbon dioxide in the environment.  These green plants steadily removed carbon dioxide and produced oxygen by photosynthesis.  Oxygen levels in the atmosphere slowly increased.

9.33 The atmosphere in a state of approximate balance
The carbon cycle helps to keep the atmospheric composition constant by adding carbon dioxide to the atmosphere and also taking carbon dioxide away from the atmosphere.  
Carbon dioxide is taken away from the atmosphere or out of the cycle by photosynthesis, dissolving in water and by chemical reactions, for example with rock. 
It is brought into the atmosphere or into the cycle by respiration, combustion, volcanic activity and decay.

9.34 The conditions for the Haber process
The optimum (best) conditions for the Haber process that turns nitrogen and hydrogen into ammonia are:

• 350 atmospheres; high pressure increases yield
• about 450ºC ; high temperature cuts yield but increases rate
• and the use of a catalyst, which is usually iron; increases rate

9.35 Dynamic equilibrium
The chemical reaction used to make ammonia is:
nitrogen + hydrogen --> ammonia
N₂         +      3H₂    ---> 2NH₃
Ammonia can also decompose to form nitrogen and hydrogen:
2NH₃ ---> N₂ +3H₂
This reaction is reversible.  The reaction is never complete but does reach a state when no more change can be seen.  This state is called equilibrium.

Although no change is seen at equilibrium the reaction still carries on with some ammonia molecules being made and some decomposed.  This is called dynamic equilibrium.

9.36 Explaining the choice of conditions used in the Haber process
Temperature
The reaction below is exothermic, energy is given out when ammonia forms, but energy is taken in if ammonia breaks up.
N₂(g)  +   3H₂(g)    ---> 2NH₃(g)  ;  DH = - 92kJ/mol
Reactions resist changes.

If the temperature goes up the reaction tries to prevent this by taking in energy.
The reaction can take in energy by breaking up ammonia. 
The position of the equilibrium moves to the left.
So if the temperature goes up then ammonia breaks up which is not helpful.

If the temperature goes down the reaction tries to prevent this by giving out energy.
The reaction can give out energy by forming ammonia. 
The position of the equilibrium moves to the right.
So if the temperature goes down then ammonia form which is helpful.

Pressure
On the left hand side of the equation there are 4 molecules
On the right hand side there are only 2 molecules which take up less space than 4 molecules
Pressure is reduced if the position of the equilibrium moves to the right, so an increase in pressure causes a shift to the right so more ammonia is formed.  As high pressure favours a big yield of ammonia so 200 atmospheres pressure is used.

9.37 Nitrogenous fertilisers and growth in plants
Plants grow well when they can obtain fixed nitrogen from the soil.  Only a few plants like peas and beans can make use of nitrogen in the air.  Most plants need fixed nitrogen in compounds like nitrates.  Plants need nitrogen to make proteins which gives them strong stems and healthy leaves.  Some nitrates find their way into the soil naturally but intensive farming removes a lot when crops are harvested.  Fertilisers containing nitrogen are used to replace nitrogen lost from the soil during farming.

9.38  Manufacturing nitrogenous fertiliser
When nitric or sulfuric acid reacts with ammonia (an alkali) the acid is neutralised and a salt is formed.
acid + alkali ---> salt + water
nitric acid + ammonium hydroxide ---> ammonium nitrate + water
HNO₃(aq) + NH₃OH(aq) ---> NH₄NO₃(aq) + H₂O(l)
Ammonium nitrate contains fixed (chemically combined) nitrogen and so is a good fertiliser.

9.39 The environmental consequences of the over-use of fertilisers
1. Fertilisers are very soluble.
2. Fertilisers dissolve in rain water.
3. Fertilisers are leached from the soil and washed into rivers.
4. Water plants grow very well in fertilised river water.
5. The over growth of plants like algae at the surface cuts out the light to plants below.
6. Plants without light stop growing and die.
7. Dead plants rot due to bacteria that use a lot of oxygen
8 The amount of oxygen in the water drops.
9 Fish and other animals start to die because of a lack of oxygen.
10 The process is called eutrophication.

9.40 Noble gases as chemically inert
Noble gases are monatomic. This means they exist only as single atoms. Their atoms cannot combine with other atoms because in their electronic structure, the outer shell is always full and this makes noble gases unreactive.

9.41 Lack of reactivity and the electronic arrangement
He 2;                Ne 2,8;              Ar 2,8,8
All of these noble gases have 2 or 8 outer electrons.  These are the maximum numbers of electrons for the outer shells.  These atom cannot now lose gain or share electrons so do not form chemical bonds.  They are therefore unreactive.

9.42 Uses of the noble gases

Task 9.42