Chemical changes

Magnesium burns in oxygen to form $\mathrm{MgO}$. Name the type of reaction and explain your choice in terms of oxygen.

A student heats $6.4\,\mathrm{g}$ of copper in oxygen and obtains $8.0\,\mathrm{g}$ of copper oxide. Calculate the mass of oxygen gained and state what has happened to the copper.

Hydrogen is passed over hot copper oxide: $\mathrm{CuO}+\mathrm{H_2}\rightarrow\mathrm{Cu}+\mathrm{H_2O}$. Identify what is oxidised and what is reduced, explaining each answer using oxygen transfer.

Place magnesium, copper and zinc in decreasing order of reactivity.

Metal P displaces Q from a solution of Q ions. Metal Q displaces R from a solution of R ions. Metal R does not displace P from a solution of P ions. Deduce the reactivity order of P, Q and R and justify it.

Four metals W, X, Y and Z are tested at room temperature. W reacts rapidly with cold water. X reacts slowly with dilute acid but not with cold water. Y does not react with cold water, reacts quickly with dilute acid and displaces X from an X salt solution. Z shows no reaction with dilute acid. Deduce their order from most to least reactive and explain the position of Z relative to hydrogen.

Copper is less reactive than carbon. Explain why carbon can be used to obtain copper from copper oxide.

Aluminium is above carbon, zinc is below carbon and gold is very unreactive. For each metal, choose the most suitable description: found as the metal itself, extracted from its oxide using carbon, or requires a method other than carbon reduction. Explain each choice.

Consider $2\mathrm{CuO}+\mathrm{C}\rightarrow2\mathrm{Cu}+\mathrm{CO_2}$. Identify the substance oxidised and the substance reduced, explaining each identification using oxygen transfer. A proposed alternative uses less fuel but produces only half as much copper per batch. State two pieces of additional information needed to judge which route is preferable.

The half equation is $\mathrm{Zn}\rightarrow\mathrm{Zn^{2+}}+2\mathrm{e^-}$. State whether zinc is oxidised or reduced and explain why.

Complete and balance the half equation $\mathrm{Cl_2}+\square\,\mathrm{e^-}\rightarrow\square\,\mathrm{Cl^-}$ and identify the process.

Magnesium is added to copper(II) ion solution. Write the oxidation half equation, the reduction half equation and the overall ionic equation. Identify the species oxidised and the species reduced.

Name the two products when magnesium reacts with dilute hydrochloric acid.

Write a balanced symbol equation for iron reacting with dilute sulfuric acid to form iron(II) sulfate and hydrogen.

Higher Tier: magnesium reacts with hydrochloric acid according to $\mathrm{Mg}+2\mathrm{H^+}\rightarrow\mathrm{Mg^{2+}}+\mathrm{H_2}$. Write the oxidation and reduction half equations, then identify the species oxidised and the species reduced.

Potassium hydroxide is neutralised by nitric acid. Name the salt formed and the other product.

Complete the word equation and explain the salt name: copper oxide + sulfuric acid $\rightarrow$ ______ + ______.

Aluminium hydroxide reacts with sulfuric acid. Deduce the formula of aluminium sulfate and balance the symbol equation.

Why is an insoluble solid added to an acid until some remains unreacted when making a soluble salt?

Put these stages for making copper sulfate crystals from copper oxide and dilute sulfuric acid into a safe, logical order: crystallise, filter, warm the acid, add excess copper oxide, evaporate some water, dry the crystals.

Describe how to prepare a pure, dry sample of zinc chloride crystals using zinc carbonate and dilute hydrochloric acid. Include the purpose of each separation or heating stage.

Classify solutions with pH $3$, pH $7$ and pH $11$ as acidic, neutral or alkaline.

A colourless solution may have pH $5$, $7$ or $9$. Describe how to determine its approximate pH and state one method that would give a more precise numerical value.

An alkali is added in small portions to an acidic solution until its pH changes from $2$ to $7$. Explain the pH change in terms of ions, write the ionic equation for neutralisation, and explain why adding equal volumes of acid and alkali would not always give pH $7$.

Name the apparatus used to transfer exactly $25.0\,\mathrm{cm^3}$ of alkali and the apparatus used to add a measured, variable volume of acid.

Describe how to obtain an accurate mean titre when titrating a strong alkali with a strong acid.

Higher Tier: $25.0\,\mathrm{cm^3}$ of sodium hydroxide is neutralised by $18.60\,\mathrm{cm^3}$ of $0.150\,\mathrm{mol\,dm^{-3}}$ hydrochloric acid. The equation has a $1:1$ ratio. Calculate the sodium hydroxide concentration in $\mathrm{mol\,dm^{-3}}$ and $\mathrm{g\,dm^{-3}}$. Use $M_r(\mathrm{NaOH})=40.0$.

State the difference between a strong acid and a weak acid in aqueous solution.

Equal-concentration solutions of hydrochloric acid and ethanoic acid are compared. Predict which has the lower pH and explain why. State whether the comparison tells you that either solution is more concentrated.

Solution A has pH $2$ and solution B has pH $5$. Calculate how many times greater the hydrogen ion concentration is in A than in B. Explain why saying 'A is $3$ times more acidic' is incorrect.

State which electrode attracts positive ions during electrolysis and give the charge on that electrode.

Solid sodium chloride does not conduct electricity, but molten sodium chloride does. Explain this difference in terms of ions.

A molten ionic compound contains only $\mathrm{M^{2+}}$ and $\mathrm{X^-}$ ions. Describe the movement of both ions when a direct current is applied, name the type of substance formed at each electrode, and state what carries charge through the melt.

Predict the products at the cathode and anode when molten zinc chloride is electrolysed using inert electrodes.

Molten lead bromide is electrolysed with inert electrodes. Name each product, identify its electrode, and explain why lead bromide must be molten.

A binary compound has formula $\mathrm{CaCl_2}$. Predict both electrolysis products in the molten state and explain how the formula supports the relative numbers of calcium ions and chloride ions discharged.

Explain why aluminium is extracted using electrolysis rather than by reducing aluminium oxide with carbon.

Cryolite is mixed with aluminium oxide before electrolysis. Explain how this reduces the energy cost of extraction.

An aluminium plant considers two electrolytes. Mixture A melts at $950\,^{\circ}\mathrm{C}$; mixture B melts at $1210\,^{\circ}\mathrm{C}$. Both contain the same amount of aluminium oxide and give the same aluminium output. The cells use carbon anodes. Choose the likely lower-energy mixture and explain both your choice and why the anodes need regular replacement.

Predict the products at inert electrodes during electrolysis of aqueous sodium chloride.

Aqueous copper(II) sulfate is electrolysed using inert electrodes. Predict both products and justify each using the discharge rules.

Predict and compare the electrode products for aqueous magnesium chloride and aqueous silver nitrate, both with inert electrodes. Explain every product using relative reactivity or the halide rule.

Complete the cathode half equation $\mathrm{Al^{3+}}+\square\,\mathrm{e^-}\rightarrow\mathrm{Al}$ and name the process.

Write the balanced half equation for bromide ions forming bromine at the anode and explain why it is oxidation.

During electrolysis of aqueous sodium sulfate with inert electrodes, hydrogen forms at the cathode and oxygen forms at the anode. Write a balanced half equation for each electrode and label oxidation and reduction.