4.9 describe experiments to carry out acid-alkali titrations

• Have a known volume of acid in a beaker with methyl orange, the solution will be red as it is very acidic.
• Set up a burette with alkali in it. Open the tap very slightly, so that it drips very slowly into the acid.
• With each drop stir the contents of the beaker.
• The more alkali that is added the more neutral and closer to orange it gets.
• When the solution is neutral the it will be completely orange, at this point close the tap on the burette.
• The level in the burette will have dropped, showing the volume of alkali used.
• This shows you how much alkali you need to use to neutralise the acid.

4.8 describe experiments to prepare insoluble salts using precipitation reactions

Silver nitrate and sodium chloride are added together, the product, silver chloride is made, this salt is insoluable and so will form a white precipitate in the solution.

AgNO3 + NaCl > AgCl + NaNO3

4.7 describe experiments to prepare soluble salts from acids

Dilute sulphuric acid is added to an excess of magnesium.
Mg + H2SO4 > MgSO4 + H2
The left over magnesium is filtered off and the mixture boiled down slowly to concentrate.
When it is cooled, crystals will form, these can be blotted dry with a piece of paper.

This can be done with other metals and acids:

Nitric, sulphuric and hydrochloric acids make soluble salts with most metals.

ammonium, potassium and sodium make soluble salts with acids

4.6 understand the general rules for predicting the solubility of salts in water

i) all common sodium, potassium and ammonium salts are soluble

ii) all nitrates are soluble

iii) common chlorides are soluble, except silver chloride

iv) common sulfates are soluble, except those of barium and calcium

v) common carbonates are insoluble, except those of sodium, potassium and ammonium

4.5 predict the products of reactions between dilute hydrochloric, nitric and sulfuric acids; and metals, metal oxides and metal carbonates

Metals

• Hydrochloric acid + metal > metal chloride salt + hydrogen
• Nitric acid + metal > You don't need to know
• Sulphuric acid + metal > metal sulphate + hydrogen

Metal oxides

• Hydrochloric acid + metal oxide > metal chloride salt + water
• Nitric acid + metal oxide > metal nitrate salt + water
• Sulphuric acid + metal oxide > metal sulphate + water

Metal carbonates

• Hydrochloric acid + metal carbonate > metal chloride salt + water + carbon dioxide
• Nitric acid + metal carbonate > metal nitrate salt + water + carbon dioxide
• Sulphuric acid + metal carbonate > metal sulphate + water + carbon dioxide

4.4 define acids as sources of hydrogen ions, H+, and alkalis as sources of hydroxide ions, OH¯

Essentially if something is acidic it contains positive hydrogen ions, if something is alkaline it contains negative hydroxide ions.

4.3 describe the use of universal indicator to measure the approximate pH value of a solution

Neutral is green

The more red, the more acidic.

The more purple, the more alkaline.

4.2 understand how the pH scale, from 0–14, can be used to classify solutions as strongly acidic, weakly acidic, neutral, weakly alkaline or strongly alkaline

4.1 describe the use of the indicators litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions

Red litmus paper turns blue in the presence of alkali.

Blue litmus paper turns red in the presence of acid.

Phenolphthalein goes pink in alkalis.

Methyl orange is yellow for alkali, red for acid.

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