GCSE Chemistry Acids & salts questions by topic

Complete the sentences about acids and alkalis.

1. (a)Acidic solutions have a pH … than 7.[1]
2. (b)Acids release … ions in water.[1]
3. (c)Aqueous alkalis release … ions.[1]
4. (d)Name one indicator used to estimate pH.[1]

Solutions A to D are tested with universal indicator and a pH probe.

1. (a)Which solution is strongly acidic?[1]
2. (b)Which solution is neutral?[1]
3. (c)Which solution could be oven cleaner?[1]
4. (d)State the colour universal indicator shows in solution A.[1]

Name the salt formed in each reaction.

1. (a)Hydrochloric acid + sodium hydroxide[1]
2. (b)Sulfuric acid + copper oxide[1]
3. (c)Nitric acid + potassium hydroxide[1]
4. (d)Hydrochloric acid + zinc carbonate[1]

Write word equations for these reactions.

1. (a)Zinc + hydrochloric acid →[2]
2. (b)Calcium carbonate + sulfuric acid →[2]

A student makes copper sulfate crystals by reacting copper oxide with warm sulfuric acid.

1. (a)Why is the copper oxide added in excess?[2]
2. (b)Why is the mixture filtered before evaporation?[2]
3. (c)State one safety precaution.[1]

A student adds 2 cm of cleaned magnesium ribbon to dilute sulfuric acid and collects the gas given off.

1. (a)Name the gas produced and state the test for it.[2]
2. (b)Write the word equation, naming the salt and the gas.[2]
3. (c)State one observation, other than gas production, that shows a reaction is happening.[1]

A farmer tests a field and finds the soil is acidic, with a pH of 5.2. Crops grow best near pH 6.8.

1. (a)Name a chemical the farmer could spread on the field, and state what type of substance it is.[2]
2. (b)After treatment the soil pH is 6.8. Explain why the farmer should not simply add much more of the chemical “to be safe”.[3]

A student compares hydrochloric acid of pH 1 with ethanoic acid of pH 3, both 50 cm³.

1. (a)Compare the hydrogen ion concentrations of the two acids.[2]
2. (b)The student adds equal masses of magnesium ribbon to each. Describe and explain the difference in the fizzing.[3]
3. (c)Explain why both acids can eventually neutralise the same amount of alkali if given time.[1]

Write balanced symbol equations with state symbols.

1. (a)Neutralisation of sodium hydroxide solution by hydrochloric acid.[2]
2. (b)Copper(II) oxide powder reacting with warm sulfuric acid solution.[2]
3. (c)The ionic equation for reaction (a).[2]

Two indigestion remedies are compared: tablets containing calcium carbonate, and tablets containing magnesium hydroxide.

1. (a)Write the equation for calcium carbonate neutralising hydrochloric acid.[2]
2. (b)Explain why the carbonate tablet may cause a feeling of bloating.[2]
3. (c)Explain why both tablets raise the stomach pH.[2]

A class investigates the practical: copper oxide + sulfuric acid → copper sulfate.

1. (a)State the dependent variable if the class measures how well the method purifies the product.[1]
2. (b)State two control variables.[2]
3. (c)A student heats the acid strongly until it boils. Explain two things that could go wrong.[2]
4. (d)The crystals come out green-tinged instead of blue. Suggest one impurity and how it got there.[2]

Sodium carbonate solution reacts with dilute hydrochloric acid, fizzing vigorously.

1. (a)Write the full balanced symbol equation, with state symbols.[2]
2. (b)Write the ionic equation for the reaction, and explain why the sodium and chloride ions do not appear in it.[4]

Four solutions are measured with a calibrated pH probe: A reads pH 1, B reads pH 3, C reads pH 5, D reads pH 7.

1. (a)Calculate how many times more concentrated the hydrogen ions in A are than in C.[2]
2. (b)State which solution could be pure water at room temperature.[1]
3. (c)Solution B is diluted 100-fold with water. State and explain the new pH.[3]

Equal volumes of two acids are titrated with the same sodium hydroxide solution. Acid X (pH 1) needs 25.0 cm³ of alkali; acid Y (pH 3) needs 25.0 cm³ too.

1. (a)Compare the hydrogen ion concentrations of X and Y.[2]
2. (b)Explain why the equal titres do not mean the acids have equal concentrations in moles per dm³.[3]
3. (c)Suggest which acid is more likely to be weak, with a reason.[2]

A factory discharges dilute hydrochloric acid waste. The effluent must be neutralised before release into a river.

1. (a)Suggest a suitable alkali for the neutralisation and name the salt formed.[2]
2. (b)Explain why the effluent must not be over-neutralised into strongly alkaline water.[2]
3. (c)The plant measures pH continuously with a probe. Explain why a probe beats universal indicator paper for this job.[2]
4. (d)The waste also contains suspended metal oxide solids. State how they could be removed.[1]

A student follows the salt-making practical with zinc oxide and sulfuric acid, but the final crystals are contaminated with acid.

1. (a)Explain the most likely mistake in the method that caused acid contamination.[2]
2. (b)Describe how the student could test a small sample of the crystals to confirm excess acid is present.[3]
3. (c)Describe how to salvage the batch and obtain pure zinc sulfate crystals.[3]

A student must prepare pure dry potassium sulfate crystals. The only reagents supplied are potassium hydroxide solution and dilute sulfuric acid.

1. (a)Explain why the excess-solid-and-filter method used for copper sulfate cannot work here.[2]
2. (b)Describe how the student should carry out the preparation, naming the technique that finds the right amounts.[4]
3. (c)Write the balanced symbol equation, with state symbols.[2]

A class adds sodium hydroxide solution to 25.0 cm³ of hydrochloric acid, 5 cm³ at a time, measuring the pH and the temperature after each addition.

1. (a)State the volume of sodium hydroxide that exactly neutralised the acid, using evidence from the table.[2]
2. (b)Explain the pattern shown by the temperature column.[2]
3. (c)Describe and explain the pH change between 20 cm³ and 30 cm³ added.[2]
4. (d)Suggest why recording both pH and temperature gives stronger evidence than either measurement alone.[2]