The Mole and Stoichiometry: Chemistry 0620 (Cambridge O Level / IGCSE)

Barium nitrate, $\text{Ba(NO}_3\text{)}_2$, is the compound added to fireworks to produce a bright green flame colour.

What is the relative formula mass, $M_r$, of barium nitrate? ($A_r$: Ba $= 137$, N $= 14$, O $= 16$)

Zinc sulfide is the pigment used to make glow-in-the-dark plastic toys and safety signs: after being exposed to light, it slowly releases the energy again as a green glow. It is made by directly heating powdered zinc together with powdered sulfur.

A simple model of solid zinc sulfide shows one zinc atom bonded to one sulfur atom, with this 1:1 pattern repeating throughout the solid.

(a) Use this description of the structure to deduce the formula of zinc sulfide. [1]

(b) Construct the word equation for the reaction between zinc and sulfur. [1]

(c) Construct the symbol equation for this reaction, including state symbols. [2]

(d) Relative atomic mass, $A_r$, is found by comparing the mass of atoms of an element with the mass of atoms of one particular reference element. Describe what is meant by the relative atomic mass of zinc. [2]

A museum conservator is investigating the dark silver sulfide tarnish, $\text{Ag}_2\text{S}$, that forms when silverware reacts with sulfur-containing gases in the air.

In a small controlled trial, $5.4\ \text{g}$ of silver reacts completely to form $6.2\ \text{g}$ of silver sulfide.

(a) An heirloom silver tray contains $21.6\ \text{g}$ of silver. Using the same proportions as the trial (do not use the mole concept), calculate the mass of silver sulfide that would form if all of this silver reacted in the same way. [2]

(b) The conservator cleans the tarnished silver using a solution of ammonia. State the chemical formula of ammonia. [1]

(c) State the two units in which the concentration of this ammonia solution could be expressed. [1]

A ceramics workshop produces a blue glaze pigment by strongly heating cobalt(II) carbonate, $\text{CoCO}_3$, until it fully decomposes into cobalt(II) oxide (the blue pigment) and carbon dioxide gas.

$\text{CoCO}_3\text{(s)} \rightarrow \text{CoO(s)} + \text{CO}_2\text{(g)}$

A batch of $23.8\ \text{g}$ of cobalt(II) carbonate is heated until decomposition is complete. ($A_r$: Co $= 59$, C $= 12$, O $= 16$)

(a) Calculate the number of moles of $\text{CoCO}_3$ in $23.8\ \text{g}$. [2]

(b) Use the balanced equation to calculate the number of moles, and then the mass, of cobalt(II) oxide produced. [2]

(c) Calculate the volume of carbon dioxide gas produced, measured at room temperature and pressure (r.t.p.). [2]

A pigment manufacturer heats $10.4\ \text{g}$ of chromium metal in a stream of oxygen gas until it is completely converted into a green oxide used as a ceramic and paint pigment. The mass of the green oxide formed is $15.2\ \text{g}$. ($A_r$: Cr $= 52$, O $= 16$)

(a) Calculate the empirical formula of this chromium oxide. [3]

(b) For this batch, the manufacturer's target mass of oxide was $16.0\ \text{g}$. Calculate the percentage yield actually obtained. [2]

(c) Chromium forms the ion $\text{Cr}^{3+}$ and oxygen forms the ion $\text{O}^{2-}$. Use the charges on these ions to deduce the formula of chromium oxide, and state whether it agrees with your answer to (a). [2]

Aluminium powder burns in oxygen gas to form aluminium oxide, the compound responsible for the thin, protective coating that gives aluminium metal its natural resistance to corrosion.

Which equation correctly represents this reaction, fully balanced?

A blacksmith heats a ball of iron wool in a Bunsen burner flame in the open air. The iron reacts with oxygen gas to form iron(III) oxide, a reddish-brown solid.

(a) Construct the balanced symbol equation, including state symbols, for this reaction. [2]

(b) Calculate the relative formula mass, $M_r$, of iron(III) oxide, $\text{Fe}_2\text{O}_3$. ($A_r$: Fe $= 56$, O $= 16$) [1]

(c) A separate sample contains $5.6\ \text{g}$ of iron. Calculate the number of moles of iron atoms in this sample. ($A_r$: Fe $= 56$) [2]

A laboratory technician needs to prepare $250\ \text{cm}^3$ of a standard solution of sodium carbonate, $\text{Na}_2\text{CO}_3$, with a concentration of $0.200\ \text{mol/dm}^3$, for use in a titration to test the acidity of a sample of lake water. ($A_r$: Na $= 23$, C $= 12$, O $= 16$)

(a) Convert $250\ \text{cm}^3$ into $\text{dm}^3$. [1]

(b) Calculate the number of moles of $\text{Na}_2\text{CO}_3$ needed to make this solution. [2]

(c) Calculate the relative formula mass, $M_r$, of $\text{Na}_2\text{CO}_3$. [1]

(d) Calculate the mass of $\text{Na}_2\text{CO}_3$ that the technician must weigh out. [2]

(e) Calculate the concentration of this solution in $\text{g/dm}^3$. [1]

A technician adds excess dilute sulfuric acid to $5.4\ \text{g}$ of aluminium turnings to generate hydrogen gas for a fuel-cell demonstration.

$2\text{Al(s)} + 3\text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{Al}_2\text{(SO}_4\text{)}_3\text{(aq)} + 3\text{H}_2\text{(g)}$

($A_r$: Al $= 27$)

(a) Calculate the number of moles of aluminium atoms in $5.4\ \text{g}$. [2]

(b) Use the balanced equation to calculate the number of moles of hydrogen gas produced. [2]

(c) Calculate the volume of hydrogen gas produced, measured at room temperature and pressure (r.t.p.). [2]

A hydrocarbon gas used as a starting material in plastics manufacture contains only carbon and hydrogen. Analysis shows it has the following percentage composition by mass: carbon $85.7\%$, hydrogen $14.3\%$. Its relative molecular mass is $42$. ($A_r$: C $= 12$, H $= 1$)

(a) Calculate the empirical formula of this hydrocarbon. [3]

(b) Use the relative molecular mass to determine the molecular formula of the hydrocarbon. [2]