10.1 Oxidation and reduction
The electrochemical series (page 7 of data book) lists metals down the right hand side in order of losing electrons and down the left hand side are metal ions in order of ease of gaining electrons therefore a metal from the right will react with a substance below it and on the left
3rd top Ca2+ + 2e ------ > Ca
7th top Zn2+ + 2e------ > Zn so Ca(s) reacts with Zn2+
8th top Fe2+ + 2e------ > Fe
15th top Cu2+ + 2e------ > Cu so Fe(s)reacts with Cu2+
(a) displacement as another substance is pushed out
eg. Ca(s) + Zn2+(aq) -----> Zn(s) + Ca2+(aq)
ZINC IS DISPLACED BY CALCIUM
(b) redox as one substance is OXIDISED (loses electrons)
Ca(s)------ > Ca2+ (aq) + 2e
(equation in opposite direction from page 7 of data book)
and another substance is REDUCED (gains electrons)
Zn2+(aq) + 2e------ > Zn(s) (equation in right direction as in data book)
Some metals displace hydrogen from acids showing where hydrogen appears in the electrochemical series.
10.2 Redox/displacement reactions where electrons are lost and gained can be carried
out in a CELL where two different metals are placed in solutions of their own ions
Example of copper-zinc cell (ignore anode and cathode words)
Zinc loses electrons to Cu2+(aq) via the electrons moving through the metal wire - an electric current.
a zinc-copper cell in sulphuric acid
10.3 Different pairs of metals produce different voltages and this produced the electrochemical series - the bigger the voltage the further apart in the electrochemical series.
Cells of non-metals (or metal / non-metal can also produce electricity)
The solutions used allow ions to flow and are called an electrolytes - helps complete the circuit
The ion bridge (ions travel on it) completes the circuit.
10.4 In a battery the electricity comes from a chemical reaction producing a flow of negatively
charged electrons along a metal wire. A battery stops producing electricity when one of the chemicals in the reaction is used up. Some batteries are rechargeable e.g. lead-acid car battery.
Batteries are more portable and safer but more expensive and use up more resources than mains electricity.
11.1 Properties of metals
Electrical conductors when solid or molten - electrical wiring
Heat (thermal) conductors - radiators
Strength - anything being supported - electricity pylons
Density - (low density) for aeroplanes
Properties of metals can be improved by mixing metals together. A mixture of metals is called an alloy.
11.2 Metals react with substances like oxygen, water and acids at different rates of reaction - if a metal reacts faster with oxygen then it will react faster with other substances - this produces an order of reactivity of metals.
Very reactive - Na, Li, K, Ca - fast, explosive - even at low temperatures (Rb and Cs)
Mid reactive - Mg, Al, Zn, Fe, Sn - readily, slower - heat helps.
Unreactive - Cu, Ag, Au - slowly, do not react.
Page 7 data book gives useful order of reactivity.
If a metal reacts with:
(i) oxygen then the metal oxide is produced + hydrogen gas
(ii) water then a hydroxide or oxide is produced + hydrogen gas
(iii) acid then a salt is produced + hydrogen gas
When metals react they are oxidised (i.e.lose electrons).
11.3 The unreactive metals were the first to be discovered as they are found uncombined
(as elements) in nature.
Ores are the naturally occurring compounds (usually oxides, sulphides or carbonates
The more reactive the metal in a compound, the harder it is to decompose the
ore to produce the metal (reduction)
EASY TO MAKE COMPOUND HARD TO BREAK COMPOUND
M ------ > M2+ + 2e M2+ + 2e------ > M
Some metals (unreactive) can be obtained by heating ore.
Other metals (mid reactive) require an agent to remove the oxygen from the ore e.g., heat with carbon or hydrogen or carbon monoxide.
Very reactive metal ore requires energy from electricity to break up the compound by electrolysis into the elements.
As metals and metal ores are a finite resource, scrap metal needs to be recycled (this saves energy and money, as well as resources).
11.4 Iron metal is produced in a blast furnace by
C + O2 ------ > 2CO carbon burned to produce carbon monoxide
CO + FeO ------ > CO2 + Fe carbon monoxide reduces iron oxide to iron
11.5 Alloys are a mixture of metals (brass in copper and zinc) or metals and non-metal (iron and carbon is steel) and are produced to improve properties or achieve properties for a specific purpose.
11.6 Calculate the empirical formula
Unit 12 CORROSION
12.1 Corrosion is a chemical change involving the surface of a metal reacting with water and oxygen in the presence of an electrolyte (dissolved C02) to produce a compound - usually an oxide. Different metaIs corrode (oxidise) at different rates.
Corrosion is an example of OXIDATION
12.2 The corrosion of iron (oxidation of iron by water and air) is called rusting.
Iron loses two electrons Fe ------ > Fe2+ + 2e
The extent of rusting can be shown using ferroxyl indicator which turns blue in presence
Fe2+(aq) can be oxidised to Fe3+(aq)
The electrons are gained by the oxygen and water to produce hydroxide ions
Fe(s) becoming Fe2+(aq) and showing up as blue with ferroxyl indicator can indicate where oxidation occurred in a cell - this electrode would lose electrons.
12.3 The presence of salt or acid rain in water increases the rate of corrosion as the
dissolved ions produce an electrolyte. Heat also speeds up all chemical reactions, including corrosion.
12.4 Iron does not rust when attached to the negative terminal of a battery supplying electrons to the iron.
More reactive metals close by also supply electrons to the iron (while being oxidised themselves) preventing corrosion by sacrificial protection.
Iron will corrode faster if a less reactive metal, such as tin, is close by - losing
electrons from the iron to the other metal.
12.5 A surface barrier to water and air (oxygen) can prevent corrosion e.g. grease,
paint,plastic or another less corrosive metal.
electroplate - plate one metal with another (e.g. iron with nickel) to act as a barrier (i.e. nickel plate or tin plating)
galvanising - covering iron with zinc as barrier
Aluminium and zinc have strong oxide layers which prevent further corrosion.