2.28 describe a physical test to show whether water is pure.

Pure water boils at 100 ^OC and freezes at 0^OC.  You can check these to confirm the identity of the liquid.

2.27 describe the use of anhydrous copper(II) sulfate in the chemical test for water

Anhydrous copper(II) sulfate is white. It turns blue, forming the hydrated salt, when water is added.  This is an exothermic reaction.

CuSO4(s) + 5H2O(l) → CuSO4.5H2O(s)

See also section 4.23 (reversible reactions)

2.39 describe simple tests for the gases:

Gas	Test	Results
I) hydrogen (H2)	introduce lighted splint	explodes with squeaky pop
I    i) oxygen (O2)	introduce glowing splint	reignites the splint
I    ii) carbon dioxide (CO2)	bubble through lime water	goes milky white*
I    v) ammonia (NH3)	introduce moist red litmus paper	turns blue
V   ) chlorine (Cl2)	introduce moist indicator paper (any colour)	indicator paper is bleached white

Note:    The limewater goes milky because a white precipitate of insoluble calcium carbonate is formed. The equation for limewater (calcium hydroxide) reacting with carbon dioxide is:

Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

2.38 describe simple tests for the anions:

2.38      describe simple tests for the anions:

i           Cl−, Br− and I−, using dilute nitric acid and silver nitrate solution

Silver chloride is insoluble in water (see Section 4.6).  Silver bromide and silver iodide are also insoluble.  Adding silver nitrate solution to a solution containing these ions will produce a precipitate of the insoluble silver halide.  The solution must be acidified in case there are any hydroxide or carbonate ions present that would also give a precipitate with the silver ions.

Take a few cm3 of the sample to be tested. Add a similar volume of dilute nitric acid. Then add a few drops of silver nitrate solution and look for a precipitate.
chloride (Cl-)	white precipitate of silver chloride Ag+(aq) + Cl-(aq) → AgCl(s)
bromide (Br-)	cream precipitate of silver bromide Ag+(aq) + Br-(aq) → AgBr(s)
iodide (I-)	yellow precipitate of silver iodide Ag+(aq) + I-(aq) → AgI(s)

ii          SO42−, using dilute hydrochloric acid and barium chloride solution

Barium sulfate is insoluble (see Section 4.6).  A white precipitate of barium sulfate will form if a solution of barium chloride is added to a solution containing sulfate ions.  The solution must be acidified in case there are any carbonate ions present that would also give a precipitate with the barium ions.  Take a few cm3 of the sample to be tested.  Add a similar volume of dilute hydrochloric acid.  Then add a few drops of barium chloride solution.

Ba2+(aq)   +   SO42−(aq)  ®     BaSO4(s)

iii         CO32−, using dilute hydrochloric acid and identifying the carbon dioxide evolved

If carbonate ions are present, carbon dioxide gas will be given off when dilute HCl is added. Take a few cm3 of the sample to be tested.  Add a similar volume of dilute hydrochloric acid.  Bubble the gas evolved through limewater and the limewater will turn milky.

This test will also work on a solid sample of a carbonate.  (See Section 4.5)

CO32-(aq) + 2H+(aq)  → CO2(g)   +   H2O(l)

2.37 describe simple tests for the cations:

Analysis of Ions and testing for gases

i           Li+, Na+, K+, Ca2+ using flame tests

To do a flame test, you need to use a piece of nichrome wire, dipped in concentrated hydrochloric acid.  Dip the wire in concentrated hydrochloric acid then place it in a roaring Bunsen flame to clean it.  Dip into the acid again, and then into the sample to be tested, and then hold in a blue Bunsen flame.  (Alternatively, use a wet wooden splint to put the sample into the flame.)  The metal ions in the compounds give different colours to the Bunsen flame:

Metal	Flame colour
lithium (Li+)	red
sodium (Na+)	yellow
potassium (K+)	lilac
calcium (Ca2+)	brick red (an orangey red)

Ii          NH4+ using sodium hydroxide solution and identifying the ammonia evolved

Add some dilute sodium hydroxide solution to a sample of the substance in a test tube and warm the mixture. Test any gas given off with moist red litmus paper.  If the substance contains ammonium ions, then ammonia gas will be given off.  The red litmus paper will turn blue, because ammonia is an alkaline gas.

NH4+(aq) + OH-(aq)  → NH3(g)   +   H2O(l)

iii         Cu2+, Fe2+ and Fe3+ using sodium hydroxide solution

Add a few drops of sodium hydroxide solution to a few cm3 of the sample to be tested.  Look for a coloured precipitate of the metal hydroxide to form.
copper(II) (Cu2+)	blue precipitate of copper (II) hydroxide Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s)
iron(II) (Fe2+)	green precipitate iron (II) hydroxide Fe2+(aq) + 2OH-(aq) → Fe(OH)2(s)
iron(III) (Fe3+)	brown precipitate iron (III) hydroxide Fe3+(aq) + 3OH-(aq) → Fe(OH)3(s)

2.36 understand the sacrificial protection of iron in terms of the reactivity series.

If iron is covered with a more reactive metal such as zinc, even if the coating is damaged to reveal the iron, the zinc will continue to react in preference to the iron because the zinc is more reactive, and the exposed iron is not affected. The more reactive metal is sacrificed to protect the iron.

A good example is bolting blocks of Mg or Zn to ships’ hulls. These react with the saltwater, but are easily replaced. The steel hulls are protected.

2.35 describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

Grease, oil, paint and plastic are called “barrier methods”.  They stop air and water from accessing the iron.  Remember a good example of each method, e.g., grease for axles on cars, oil for a bicycle chain, paint for iron railings and plastic for the shelves in a fridge.

Galvanising involves covering the iron completely with zinc.  This is a form of “sacrificial protection” as well.  It is suitable for smaller objects, e.g., car bodies, buckets and nails.  See the next Section 2.36.

2.34     describe the conditions under which iron rusts

Iron rusts in the presence of oxygen (air) and water.  Both of these substances need to be present.

Like any other reaction, this is speeded up if the temperature is raised.  Salt (sodium chloride) also speeds up the rusting process.

Note that “iron” here can refer to all types of steel except stainless steel.

5.5        explain the uses of aluminium and iron, in terms of their properties

Aluminium has a high strength for its low density. This makes it suitable for aircraft manufacture. It is also malleable and resistant to corrosion as it is covered in an oxide layer. This makes it suitable for making drinks cans.  Aluminium is an excellent conductor of heat and can be used for saucepans and cooking foil.  It is an excellent conductor of electricity and is used for overhead power cables.

Iron is generally used in the form of steel (iron with a small amount of carbon). Steel is strong and relatively cheap, so it is suitable for construction of buildings and bridges, car bodies, railway lines, tools and machinery, etc.

5.4        describe and explain the main reactions involved in the extraction of iron from iron ore (haematite), using coke, limestone and air in a blast furnace

The raw materials needed to make iron in a blast furnace are the iron ore itself, coke (which is an impure form of carbon made from coal), limestone and air.  The main ore of iron is called haematite and contains the compound iron (III) oxide, Fe2O3.

The ore contains other impurities like sand and clay.  Sand is silicon dioxide.  This would clog up the furnace if it were not removed.  Limestone (which is calcium carbonate) is used to react with the sand inside the furnace and remove it as slag (calcium silicate).

1)   The carbon (coke) burns to provide the heat required in the blast furnace

C(s) + O2(g) → CO2(g)

2)   The carbon dioxide reacts with more carbon to produce the reducing agent, carbon monoxide

CO2(g) + C(s) → 2CO(g)

3)   The reducing agent, carbon monoxide, removes the oxygen from the iron oxide (haematite)

Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g)

4)   Meanwhile the calcium carbonate (limestone) decomposes in the high temperatures

CaCO3(s) → CaO(s) + CO2(g)

5)   Calcium oxide is a basic oxide (see Section 2.3) and reacts with the acidic silicon dioxide (sand impurity) in a neutralisation reaction. This forms calcium silicate (slag) which is easily removed from the iron, because it floats on top of it at the bottom of the furnace.

CaO(s) + SiO2(s) → CaSiO3(l)