National Chemistry





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

      of metals.

      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

                           (OXIDATION)                                            (REDUCTION)

               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


example 1

example 2




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

of  Fe2+(aq).

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.


Corrosion summary

National 4 and 5



These notes are for the Scottish National Chemistry course taught in fourth year in most Scottish schools.


The notes available here are concise notes. They are NOT to be considered as material to learn from - they are for revision.


 The notes are arranged under the following topics


National 3          Unit 1

                          Unit 2

                          Unit 3


National 4 and 5  Unit 1            Rates of Reaction

                                                  Atomic Structure

                                                  Bonding and Properties

                                                  Acids and bases


National 4 and 5  Unit 2       Fuels and Homologous Series

                                                Consumer Products



National 4 and 5  Unit 3         Metals



                                                   Nuclear Chemistry

                                                  Chemical analysis