Acids, alkalis, and salts

Example questions are taken from past papers.


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

Indicator	Colour in Acid	Colour in Neutral Solution	Colour in Alkali
Universal indicator	pH 0-3: Red (Strong) pH 4-6: Yellow (Weak)	Green (pH 7)	pH 8-10: Blue (Weak) pH 11-14: Purple(Strong)
Red litmus	Red	Red	Blue
Blue litmus	Red	Blue	Blue
Methyl orange	Red (Sometimes pink is accepted)	Yellow	Yellow
Phenolphthalein	Colourless	Colourless	Pink

Example question: State the colour of methyl orange in water contaminated with a small amount of nitric acid. (1m)

Mark scheme: Red (allow pink)

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

pH 0-3: strongly acidic

pH 4-6: weakly acidic

pH 7: neutral

pH 8-10: weakly alkaline

pH 11-14: strongly alkaline

Example question: Suggest a possible pH value for water contaminated with a small amount of nitric acid. (1m)

Mark scheme: Any value in range 1-6.9. 

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

pH 0-3: Red  (strongly acidic)

pH 4-6: Yellow  (weakly acidic)

pH 7: Green  (neutral)

pH 8-10: Blue  (weakly alkaline)

pH 11-14: Purple  (strongly alkaline)

Example question: Suggest why universal indicator is more suitable than methyl orange for comparing the acidities of samples of water. (1m)

Mark scheme: more than one colour in acid / indicates pH/ shows strongly or weakly acidic / shows how acidic the water is

Do not accept just more than one colour

My answer: Universal indicator can measure across the full pH scale and has colour change depending on acidity of the sample; but methyl orange is red in all acidic solutions, regardless of how acidic it is. 

4.4 define acids as source of hydrogen ions, H+, and alkalis as source of hydroxide ions, OH-

Example question: Hydrogen bromide and hydrogen chloride have similar chemical properties. 

(i) A sample of hydrogen bromide is dissolved in water.
A piece of blue litmus paper is placed in the solution. State, with a reason, the final colour of the litmus paper. (2m) 

Markscheme: 

Colour: Red/pink

Reason:  Hydrobromic acid formed/ H+ ions present

My answer: Red. They hydrogen atoms in hydrogen bromide dissociate in water to from H+ ions, which give the solution its acidity, thus forming hydrobromic acid. 

NB: Remember, it is the  H+ ions that are responsible for acidic properties. Hydrogen bromide by itself does not have  H+ ions, so it is not an acid.   

(ii) A sample of hydrogen bromide is dissolved in methylbenzene. 
A piece of blue litmus paper is placed in the solution. State, with a reason, the final colour of  

the litmus paper. (2m)

Markscheme: 

Colour: Blue

Reason:  No acid formed / no reaction / no H+ ions

My answer: Blue. Hydrogen bromide does not dissociate in organic solvents to from  H+ ions, so the solution is not acidic, and no acid is formed. 

4.5 predict the products of reactions between dilute hydrochloric, nitric and sulphuric acids; and metals, metal oxides and metal carbonates (excluding the reactions between nitric acid and metals)

Acid + alkali à salt + water (NEUTRALISATION REACTION)

• Hydrochloric acid + sodium hydroxide à sodium chloride + water

Acid + base (metal oxide) à salt + water (NEUTRALISATION REACTION)

• Sulphuric acid + copper oxide à copper sulphate + water

Acid + metal carbonate à salt + water + carbon dioxide

• Nitric acid + sodium carbonate à sodium nitrate + water + carbon dioxide

Acid + metal à salt + hydrogen

• Hydrochloric acid + magnesium à Magnesium chloride + hydrogen

4.6 recall 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 sulphates are soluble, except those of barium and calcium

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

NB: any group 1 metal compound is soluble--they are alkali metals--dissolve in water to form alkalis!

4.7 describe how to prepare soluble salts from acids

acid + metal/metal oxide/metal carbonate

• Add acid to the excess metal/oxide/carbonate (NB: By using an excess of the metal or metal compound, you ensure that there is no acid left in the solution after reaction, so the filtrate will be a pure solution of the salt.)
• Stir
• Filter to remove excess solid
• Gently heat to evaporate the water to saturation point
• Dry the crystals with a paper towel/oven or in a dessicator 

4.8 describe how to prepare insoluble salts using precipitation reactions

Preparation of an insoluble salt:

1. Add the sodium salt solution of the anion to the nitrate salt solution of the cation until no more precipitate forms.
2. Filter to collect the residue
3. Wash the residue with cold water
4. Leave residue to dry on filter paper/dry in a warm oven

Example: Describe how to prepare a dry solid sample of silver chloride, AgCl, a salt which is insoluble in water. 

Sodium chloride + silver nitrate à silver chloride + sodium nitrate

NaCl (aq) + AgNO₃(aq) à AgCl (s) + NaNO₃ (aq)

1. Add dilute sodium chloride (chloride is the anion) solution to dilute silver nitrate (silver is the cation) solution until no more precipitate forms. 
2. Filter.
3. Wash residue with cold water.
4. Dry in a warm oven. 

4.9 describe how to carry out acid-alkali titrations

Titration:

1. fill the acid up to the mark in the burette
2. pipette 25.0cm³ sodium hydroxide into a conical flask
3. add a few drops of methyl orange indicator
4. add acid from the burette drop wise with swirling of flask
5. stop when colour change is permanent (turns pink/red)
6. note burette readings
7. repeat until concordant results are obtained (results are within 0.1 of each other)
8. take average of results

If they ask you how to prepare a soluble salt using an acid and an alkali, titration must be used. You first carry out a normal titration, and find out the exact amount of acid needed to neutralise the alkali. Then you repeat it without an indicator so that the salt is not contaminated with its colour. You remove the salt from the neutralised solution by evaporation, then you dry it.