When substances which are made of ions are dissolved in water, or melted, they can be broken down (decomposed) into simpler substances by passing an electric current through them. This process is called electrolysis. The electrical conducting solution or melt of ions is called the electrolyte.
When an ionic substance is melted or dissolved in water the ions are free to move about and move to the electrical contacts called electrodes. The electrodes are usually inert e.g. carbon or platinum.
The electron rich or negative electrode is called the cathode. The electron deficient or positive electrode is called the anode.
During electrolysis ions move to their oppositely charged electrode.
o Positively charged ions, usually hydrogen or metal ions move to the negative electrode.
Depending on the voltage, the positive ions may be reduced (e.g. Cu2+) by electron gain to deposit a metal (eg. Cu) or release hydrogen gas (H2) from hydrogen ions (H+).
o At the same time negatively charged ions move to the positive electrode. The negative ions may be oxidised by electron loss. This usually results in the release of a non-metallic gas e.g. oxygen (O2) from hydroxide ions (OH-) or chlorine from chloride ions (Cl-)etc.
Electroplating is the process of coating a conducting material with a layer of metal using the process of electrolysis.
The object to be coated is made the negative cathode and dipped into a salt solution of the metal ions of the metal to form the coating. On passing a low d.c. voltage the metal is deposited on the conducting negative cathode. For more details see 4th link below.
Anodising a metal, like aluminium, is done by making it the positive anode in an electrolysis system. When the electrolyte is sulphuric acid solution, the oxygen formed at the anode oxidises the metal surface to make a thicker metal oxide layer.
For more detailed examples see ..
o Electrolysis of sodium chloride solution o Extraction of Aluminium
o Purification of copper o Electroplating
o A more detailed introduction to Electrochemistry
o See also below for oxidation-reduction and electrode equations
Exothermic and Endothermic Reactions or Changes
EXOTHERMIC CHANGES
o Heat is released or given out to the surroundings by the materials involved, so the temperature rises.
o chemical change examples involve a new substance being formed and lots of examples (i) to (vi) below (but they are not always exothermic - see endothermic below).
o physical change examples e.g. condensation, freezing etc. all require the removal of energy from the material e.g. water, to the surroundings to produce the change in state (its the same as releasing heat, but it doesn't seem like it!).
Note: Dissolving substances in water can release heat giving a warm/hot solution e.g.
diluting concentrated sulphuric acid.
At KS3-GCSE level this exothermic 'dissolving' is considered a physical change, but at AS-A2 level they may be considered a chemical change too!
At a higher level of thinking for exothermic chemical changes: The net energy change when the energy needed to break bonds in the reactants is less than the energy released when new bonds are formed in the products.
o See also KS4 Science GCSE/IGCSE Chemistry Notes on Energy Changes for more details
A burning or combustion reaction usually means a very fast exothermic reaction where a flame is observed. It involves a highly energetic oxidation of 'fuels' where the temperature generated is so high the atoms give off light from the luminous flame zone e.g.
o (i) bunsen flame as methane gas fuel burns ...
methane + oxygen ==> carbon dioxide + water
CH4(g) + 2O2(g) ==> CO2(g) + 2H2O(l)
This is complete combustion with a pale blue flame and the products cannot react any further with oxygen.
If the oxygen supply is limited the flame is more yellow and can be 'smokey' due to soot formation (C) and dangerous since carbon monoxide (CO) can be formed.
These are examples of incomplete combustion.
methane + oxygen ==> carbon monoxide + water
2CH4(g) + 3O2(g) ==> 2CO(g) + 4H2O(l) (carbon monoxide formation)
Most people who die in house fires are poisoned by carbon monoxide (and other toxic gases) in the thick smoke rather than from burns.
or
methane + oxygen ==> carbon (soot) + water
CH4(g) + O2(g) ==> C(s) + 2H2O(l) (soot formation)
The sooty carbon particles e.g. in a candle flame, are heated to such a high temperature they become incandescent and give out yellow light, but as far as I know virtually no carbon monoxide is formed!
See also KS4 Science GCSE/IGCSE Chemistry Notes on fossil fuel combustion
o (ii) passing chlorine over hot aluminium metal to make aluminium chloride, the aluminium burns to form the chloride ...
aluminium + chlorine ==> aluminium chloride
2Al(s) + 3Cl2(g) ==> 2AlCl3(s)
o (iii) burning magnesium ribbon with a bright white flame ...
magnesium + oxygen ==> magnesium oxide
2Mg(s) + O2(g) ==> 2MgO(s)
o (i) to (iii) are all oxidation reactions, as are all 'fuel' burning reactions.
Continuous combustion requires the 'fire triangle' of heat + fuel + oxidant (oxidants like
oxygen, air or other reactive gases like chlorine or fluorine and in rockets liquids like hydrogen peroxide) o Very fast or explosive combustion:
A roaring bunsen flame (of methane burning) is an example of fast combustion and when the air (oxygen) - methane (natural gas) mixture is first ignited it is a small explosion! (equation above). It seems contradictory, but a source of ignition is needed because the C-H and O=O bonds are very strong giving a high activation energy. However, once ignited, the heat from the flame keeps the burning going.
Another explosive example is the 'squeaky pop test for hydrogen'. When a lit splint is applied there is a faint blue flame for a fraction of a second as the two gases explode to form water + heat, light and sound energy!
hydrogen + oxygen ==> water
2H2(g) + O2(g) ==> 2H2O(l)
In all these cases the high temperature reaction zone is seen as flame and an initial high energy source for ignition is needed to initiate the reaction e.g. a match or an electrical discharge.
o Slow or smouldering combustion:
In these cases no flame is seen, but a high temperature heat source is still required to start the reaction and the reaction zone is still at a high temperature e.g. the red hot slow burning of charcoal (mainly carbon), but the main combustion product is still carbon dioxide. You can only get this slow/smouldering combustion with solid combustible reactants.
Gases will tend to explode unless controlled in a burner and liquids will vaporise in the heat from the flame and so will also burn very fast with a flame.
carbon + oxygen ==> carbon dioxide
C(s) + O2(g) ==> CO2(g)
This is an example of complete combustion.
BUT quite often, with limited air/oxygen supply, carbon monoxide is readily formed,
carbon + oxygen ==> carbon monoxide
2C(s) + O2(g) ==> 2CO(g)
This is an example of incomplete combustion.
o Spontaneous combustion:
This is when combustion occurs without any application of a high energy ignition source, sometimes described as self-ignition, though in some cases heat is generated in some way which triggers the reaction.
For example, it is possible to prepare a very finely divided black powder form of iron(II) oxide. When the powder is dropped through air lots of tiny flashes of light are seen as it burns to form another iron oxide (probably Fe3O4). The reaction is triggered by heat from friction. The powder has such a large surface area that the friction caused by just falling in air produces enough heat to initiate the reaction. Powdered coal dust or very fine flour can behave in the same way and both have been responsible for serious accidents in industry.
Other substances can spontaneously ignite in air because the activation energies required are so low and the kinetic energy of the particles is sufficient for the reaction to happen without help! e.g.
the highly reactive Group 1 Alkali Metal caesium and the silicon-hydrogen compounds called silanes (SiH4, Si2H6 etc. which are like organic alkanes with the C's replaced by Si and far less stable).
Potassium and all the alkali metals below it ignite in water (Rb and Cs explosively) because the reaction is so exothermic and ignites metal vapour and the hydrogen gas produced.
See other web page for more detailed notes on energy changes and calculations.
BUT many exothermic reactions are not as dramatic as burning with a flame! e.g.
o (iv) Respiration: the relatively slow 'burning' of carbohydrates in animals/plants, but it releases plenty of energy at 37oC!
glucose + oxygen ==> carbon dioxide + water + energy
C6H12O6(aq) + 6O2(g) ==> 6CO2(g) + 6H2O(l) + energy o (v) Neutralisation: acid + alkali ==> salt + water
e.g. hydrochloric acid + sodium hydroxide ==> sodium chloride + water
which is one of the fastest reactions in water, but the mixture only warms up by 5 to 10oC! A bunsen flame reaches 1200oC in the main combustion zone!
More details further down.
o (vi) Rusting in which iron slowly reacts with water and oxygen (from air) to form the orange-brown hydrated iron oxide we call rust.
ENDOTHERMIC CHANGES
o Heat is absorbed or taken in by the materials involved from the surroundings, the system cools or has to be heated to effect the change.
o Chemical change examples of endothermic reactions e.g. thermal decomposition of limestone, cracking oil fractions, decomposition by electrolysis etc.
(i) Photosynthesis: input of energy from sunlight needed
carbon dioxide + water ==> glucose + oxygen
6CO2 + 6H2O + sunlight energy ==> C6H12O6 + 6O2
(ii) Making lime by heating limestone to over 900oC where a net input/absorption of energy is needed to bring about this thermal decomposition ...
limestone ==> quicklime + carbon dioxide
calcium carbonate ==> calcium oxide + carbon dioxide
CaCO3(s) ==> CaO(s) + CO2(g)
(iii) Cracking hydrocarbon molecules from oil to make smaller molecules, also requires this absorption of heat by the reactant molecules to break em' up', or to put it 'poshly', another example thermal decomposition e.g.
hexane => ethene + butane
C6H14 ==> C2H4 + C4H10
o Physical change examples e.g. melting, boiling, evaporation etc. all require the input of energy to effect the change of state of the material.
Note: Dissolving substances in water can absorb heat giving a cool solution e.g. dissolving ammonium nitrate salt in water.
At KS3-GCSE level this endothermic 'dissolving' is considered a physical change, but at AS-A2 level they may be considered a chemical change too!
At a higher level of thinking for endothermic chemical changes. The net energy change when the energy needed to break bonds in the reactants is more than the energy released when new bonds are formed in the products.
See also KS4 Science GCSE/IGCSE Chemistry Notes on fossil fuel combustion
Decomposition and Thermal Decomposition
Decomposition in general means to break down into small species e.g. natural organic matter decomposes with enzymes into carbon dioxide, water and nitrogen etc.
o Fermentation is form of biological degradation, catalysed by enzymes, to break down glucose sugar into the smaller molecules of ethanol ('alcohol') and carbon dioxide ...
C6H12O6(aq) ==> 2C2H5OH(aq) + 2CO2(g)
Light can cause decomposition e.g. in photography is a sort of photo-decomposition.
o silver chloride + light ==> silver + chlorine o 2AgCl ==> 2Ag + Cl2
Thermal decomposition means to break down substances into two or more substances by heat (usually endothermic reactions at temperatures well above room temperature) e.g.
o The decomposition of calcium carbonate (limestone) into calcium oxide (lime) and carbon dioxide in a high temperature lime kiln.
o calcium carbonate ==> calcium oxide + carbon dioxide o CaCO3(s) ==> CaO(s) + CO2(g)
For more details see the Extra Industrial Chemistry notes.
o The breaking down of hydrocarbons into smaller ones using a catalyst as well as a high temperature. This reaction is also known as cracking.
For more details see the Oil and its useful Products notes.
o e.g. octane ==> hexane + ethene
C8H18 ==> C6H14 + C2H4
o Other thermal decompositions which are examples of reversible reactions.
OXIDATION and REDUCTION - REDOX REACTIONS
OXIDATION - definition and examples REDUCTION - definition and examples The gain or addition of oxygen by an atom, molecule or
ion e.g. ...
(1) S + O2 ==> SO2 [burning sulphur - oxidised to sulphur dioxide]
(2) CH4 + O2 ==> CO2 + 2H2O [burning methane to water and carbon dioxide, methane oxidised as the C and H atoms gain O]
(3) 2NO + O2 ==> 2NO2 [nitrogen monoxide is oxidised to nitrogen dioxide by gaining oxygen]
(4) SO3
2- + [O] ==> SO4
2- [oxidising the sulphite ion to the sulphate ion]
The loss or removal of oxygen from a compound etc. e.g. ...
(1) CuO + H2 ==> Cu + H2O [loss of oxygen from copper(II) oxide shows it to be reduced to copper atoms]
(2) Fe2O3 + 3CO ==> Fe + 3CO2 [iron(III) oxide ore is reduced to iron metal by oxygen loss in the blast furnace]
(3) 2CO + 2NO ==> CO2 + N2 [nitrogen monoxide reduced to nitrogen by losing oxygen]
(4) CuO + Mg ==> Cu + MgO [loss of oxygen from copper(II) oxide shows it to be reduced to copper atoms]
The loss or removal of electrons from an atom, ion or molecule e.g. form chlorine molecules in electrolysis of chlorides or
The gain or addition of electrons by an atom, ion or molecule e.g. ...
(1) Cu2+ + 2e- ==> Cu [the copper(II) ion gains 2 electrons to form neutral copper atoms e.g. in electrolysis or metal displacement reactions)
(2) Fe3+ + e- ==> Fe2+ [the iron(III) ion gains an electron and is reduced to the iron(II) ion]
(3) 2H+ + 2e- ==> H2 [hydrogen ions gain electrons to form
halogen displace] neutral hydrogen molecules]
(4) Cl2 + 2e- ==> 2Cl- [chlorine molecules gain electrons to form chloride ions
An oxidising agent is the species that gives the oxygen or removes the electrons
A reducing agent is the species that removes the oxygen or acts as the electron donor
REDOX REACTIONS - in a reaction overall, reduction and oxidation must go together
Redox reaction analysis based on the oxygen definitions 1. copper(II) oxide + hydrogen ==> copper + water
o CuO(s) + H2(g) ==> Cu(s) + H2O(g)
o copper oxide reduced to copper, hydrogen is oxidised to water, hydrogen is the reducing agent (removes O from CuO) and copper oxide is the oxidising agent (donates O to hydrogen)
2. iron(III) oxide + carbon monoxide ==> iron + carbon dioxide o Fe2O3(s) + 3CO(g) ==> 2Fe(l) + 3CO2(g)
o the iron(III) oxide is reduced to iron, the carbon monoxide is oxidised to carbon dioxide, CO is the reducing agent (O remover from Fe2O3) and the Fe2O3 is the oxidising agent (O donator to CO)
3. 2NO(g) + 2CO(g) ==> N2(g) + 2CO2(g)
o nitrogen monoxide is reduced to nitrogen, carbon monoxide is oxidised to carbon dioxide, CO is the reducing agent, NO is the oxidising agent
4. iron(III) oxide + aluminium ==> aluminium oxide + iron o Fe2O3(s) + 2Al(s) ==> Al2O3(s) + 2Fe(s)
o iron(III) oxide is reduced and is the oxidising agent, aluminium is oxidised and is the reducing agent - incidentally, this is the 'thermit' reaction!
5. For more details of some of these and similar reactions see ...
o Metal Reactivity notes and Metal Extraction notes.
Redox reaction analysis based on the electron definitions 1. magnesium + iron(II) sulphate ==> magnesium sulphate + iron
o Mg(s) + FeSO4(aq) ==> MgSO4(aq) + Fe(s)
o [this is the 'ordinary' equation but this does not really show what happens in terms of atoms, ions and electrons, so we use ionic equations like the one shown below. The sulphate ion SO4
2-(aq) is called a spectator ion, because it doesn't change in the reaction and can be omitted from the ionic equation. No electrons show up in the full equations because electrons lost by x = electrons gained by y!!]
o Mg(s) + Fe2+(aq) ==> Mg2+(aq) + Fe(s)
o the magnesium atom loses 2 electrons (oxidation) to form the magnesium ion, the iron(II) ion gains 2 electrons (reduced) to form iron atoms. Mg is the reducing agent and the Fe2+ is the oxidising agent](2)(i) zinc + hydrochloric acid ==> zinc chloride + hydrogen
2. zinc + hydrochloric acid ==> zinc chloride + hydrogen o Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g)
o the chloride ion Cl- is the spectator ion]
o Zn(s) + 2H+(aq) ==> Zn2+(aq) + H2(g)
o Zinc atoms are oxidised to zinc ions by electron loss and zinc is the reducing agent, hydrogen ions are the oxidising agent (gaining the electrons) and are reduced to form hydrogen molecules]
3. copper + silver nitrate ==> silver + copper(II) nitrate o Cu(s) + 2AgNO3(aq) ==> 2Ag(s) + Cu(NO3)2(aq)
o the nitrate ion NO3
is the spectator ion o Cu(s) + 2Ag+(aq) ==> 2Ag(s) + Cu2+(aq)
o copper atoms are oxidised by the silver ion, electrons are transferred from the copper atoms to the silver ions, which are reduced. Silver ions are the oxidising agent and copper atoms are the reducing agent 4. chlorine + potassium iodide ==> potassium chloride + iodine
o Cl2(aq) + 2KI(aq) ==> 2KCl(aq) + I2(aq)
o or Cl2(aq) + 2I-(aq) ==> 2Cl-(aq) + I2(aq)
o a halogen displacement reaction, the more reactive chlorine displaces the less reactive iodine.
o Chlorine is reduced by electron gain and the iodide ions are oxidised by electron loss.
5. For more details of similar reactions see the Metal Reactivity, Metal Extraction and Group 7 The Halogens notes.
POLYMERISATION
means joining many small molecules called monomers into a long molecules of many units called a polymer and there are two principal types of polymerisation process
(1) Addition polymers are formed by (e.g. alkene) monomers adding together and forming no other products except the polymer e.g. two examples of addition polymerisation are
ethene ==> poly(ethene)
phenylethene ==> poly(phenylethene), old name polystyrene
(2) Condensation polymers are formed by one or more monomers add together, forming the polymer BUT in forming the polymer small molecules are eliminated 'between' the monomers e.g. two examples of condensation polymerisation are ...
dicarboxylic acid + diol ==> polyester + water diamine + dicarboxylic acid ==> nylon + water (1) Example equations showing addition polymerisation
(Ex. 1a) formation of poly(ethene) or 'polythene' from polymerising ethene to form
an addition polymer. No other molecule is formed - just simple addition polymerisation.
(Ex. 1b) formation of poly(chloroethene) or 'PVC' from polymerizing chloroethene to form
an addition polymer. No other molecule is formed - just simple addition polymerization.
For more examples and details of addition polymers see Useful Oil Products Part 7
(2) Example equation illustrating condensation polymerisation
+ small molecules eliminated
In the case of Nylon, for each 'red' monomer - 'blue' monomer, a link is formed at each end of each monomer molecule by eliminating a water molecule e.g. where [R] = 'rest of molecule' a single link formation reaction can be shown as
[R]-COOH + HO-[R] ==> [R]-CO-O-[R] + H2O
(Example of 2) representation of a Nylon made from two different monomers (shown as red
and green + linking atoms) joining by eliminating a small molecule between the two monomers, therefore Nylon is a condensation
polymer.
For more examples and details of condensation polymers see Useful Oil Products Part 11
NEUTRALISATION
NEUTRALISATION usually involves mixing an acid (pH <7 if soluble) with a base or alkali (pH > 7 if soluble) which react to form a neutral salt solution of around pH7
Two situations are common:
(1) Water soluble bases, called alkalis and often hydroxides, are mixed with a soluble acid such as hydrochloric, citric, sulphuric or nitric acid.
o acid + base/alkali ==> salt + water
o e.g. sodium hydroxide + hydrochloric acid ==> sodium chloride + water
NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l)
o certain carbonates like sodium carbonate, are also soluble to form alkaline solutions, and they will be similarly neutralised with 'fizzing' as carbon dioxide is formed as a 3rd product
o e.g. sodium carbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide
Na2CO3(aq) + 2HCl(aq) ==> 2NaCl(aq) + H2O(l) + CO2(g)
(2) Dissolving a water insoluble base (often an oxide) in an acid
o e.g. copper oxide + sulphuric acid ==> copper sulphate + water
CuO(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l)
o the acid can also be neutralised with a metal or a carbonate to give a salt solution o metal + acid ==> salt + hydrogen (this is also a redox reaction)
o e.g. zinc + hydrochloric acid ==> zinc chloride + hydrogen
Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g)
o insoluble carbonate + acid ==> (often soluble) salt + water + carbon dioxide
o e.g. magnesium carbonate + sulphuric acid ==> magnesium sulphate + water + carbon dioxide
MgCO3(s) + H2SO4(aq) ==> MgSO4(aq) + H2O(l) + CO2(g)
More details of these types of reactions involving acids and bases-alkalis