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48 exam-style questions with model answers, plus 64 quick multi-choice questions — every question on the site for this standard, grouped by the 16 pages of notes they come from.
Write a full answer before you reveal the model one — that comparison is where the marks come from. Every block links back to the notes that teach it.
Give the formula of the ionic compound formed between magnesium and chlorine, and state the charge on each ion.
Explain how the ionic bond in sodium oxide (Na2O) forms, referring to electron transfer and the resulting attraction.
A student says: 'In magnesium oxide, each magnesium ion is bonded to one oxide ion, so MgO is a molecule with one ionic bond.' Evaluate this statement and give a fully correct account of the bonding and structure of magnesium oxide.
Define a covalent bond and state how many electrons are in one single bond.
Explain why oxygen atoms form a double bond in O2 but chlorine atoms form only a single bond in Cl2.
Nitrogen gas (N2) is very unreactive and is used to fill food packaging, while oxygen (O2) reacts readily. Explain this difference in terms of the bonding in each molecule, and comment on what this shows about the relationship between bond order and bond strength.
Describe metallic bonding.
Explain why metals conduct electricity in the solid state, whereas ionic compounds do not.
Compare the bonding in magnesium and in magnesium chloride, and use the comparison to explain why magnesium can be hammered into a sheet while magnesium chloride shatters when struck.
State how many bonding pairs and lone pairs are around the oxygen atom in a water molecule.
Draw the Lewis structure of carbon dioxide (CO2) and explain why the molecule contains double bonds rather than single bonds.
A student draws the Lewis structure of methane with carbon in the centre, four C–H bonds and one lone pair on the carbon. Evaluate the structure, identify the error and its cause, and give the correct structure. Then explain how the same checking method would confirm a correct structure for hydrogen cyanide, HCN.
State the shape and bond angle of a methane molecule, CH4.
Explain why ammonia (NH3) is pyramidal with a bond angle of 107°, while methane (CH4) is tetrahedral with 109.5°.
Compare the shapes and bond angles of CH4, NH3 and H2O, explaining the trend. Then explain why CO2, which also has two atoms bonded to a central atom, is linear rather than bent like H2O.
Explain what makes a covalent bond polar.
Carbon dioxide contains polar bonds but is a non-polar molecule. Explain why.
CH4 and CHCl3 are both tetrahedral, yet one is polar and one is not. Explain this difference fully, and use your reasoning to predict whether CH2Cl2 is polar.
State what is overcome when a molecular substance such as iodine melts.
Explain why molecular substances generally have much lower melting points than ionic substances.
Water (M = 18) boils at 100 °C while methane (M = 16) boils at −162 °C, despite having almost the same molar mass. Explain this difference fully, and explain why the boiling of water is a physical change rather than a chemical one.
Name the four structure types and state the particles present in each.
Carbon dioxide is a gas at room temperature but silicon dioxide melts at 1710 °C. Both contain only non-metals and covalent bonds. Explain the difference.
A white solid does not conduct electricity when solid, dissolves in water to give a solution that conducts well, and melts at 770 °C. A second solid is grey, conducts electricity when solid, and melts at 660 °C. Deduce the structure of each, justify your deductions using all the evidence, and explain why a third substance melting at 3550 °C that does not conduct at all must be different from both.
State whether molecular substances have high or low melting points, and say what is overcome when they melt.
Explain why magnesium oxide has a much higher melting point than sodium chloride.
Explain why the boiling points of the group 17 elements increase from chlorine (−34 °C) to bromine (59 °C) to iodine (184 °C), and then explain why water, a much smaller molecule than iodine, still boils at 100 °C.
State the condition a substance must meet to conduct electricity.
Explain why solid potassium bromide does not conduct electricity but molten potassium bromide does.
Compare the electrical conductivity of copper, sodium chloride, diamond and graphite, explaining each in terms of structure and bonding. Account for why graphite behaves differently from diamond even though both are made only of carbon.
Define malleable and ductile, and state which structure type shows both properties.
Explain why diamond is extremely hard.
Sodium chloride and copper both have giant structures held together by strong forces, and both have high melting points. Explain why copper is malleable while sodium chloride is brittle, and evaluate the claim that 'sodium chloride shatters because its bonding is weaker than copper's'.
State the general rule for predicting solubility, and use it to say whether sodium chloride dissolves in hexane.
Explain, in terms of the particles involved, why sodium chloride dissolves in water.
A student is given three white solids: sodium chloride, sugar (a polar molecular substance) and wax (a non-polar molecular substance). Describe how solubility and conductivity tests could distinguish all three, and justify each prediction from the bonding and structure involved.
State the sign of ΔrH for an exothermic reaction and describe what happens to the temperature of the surroundings.
Explain why ΔrH is negative for an exothermic reaction, even though the surroundings get hotter.
The combustion of methane has ΔrH = −890 kJ mol−1. Sketch and describe the enthalpy diagram, state ΔrH for the reverse reaction with justification, and explain why a reaction being exothermic does not guarantee it happens quickly.
The combustion of methane has ΔrH = −890 kJ mol−1. Calculate the energy released when 2.00 mol of methane burns.
Calculate the energy released when 8.00 g of methane (M = 16.0 g mol−1) burns completely. ΔrH = −890 kJ mol−1.
Ethanol burns with ΔrH = −1367 kJ mol−1 (C2H5OH + 3O2 → 2CO2 + 3H2O). A camping stove burns 15.0 g of ethanol to heat water. Calculate the energy released, then determine what mass of propane (M = 44.1 g mol−1, ΔrH = −2220 kJ mol−1) would release the same energy. Comment on which is the better fuel by mass.
State whether melting is exothermic or endothermic, and give the sign of ΔH.
Explain why the temperature of a substance stays constant while it is melting, even though heat is still being supplied.
Explain why a burn from steam at 100 °C is more severe than a burn from liquid water at 100 °C, using enthalpy changes. Then explain why boiling water requires much more energy than melting the same mass of ice, in terms of the particles.
State whether breaking bonds is exothermic or endothermic, and whether making bonds is exothermic or endothermic.
Calculate ΔrH for H2 + Cl2 → 2HCl using bond enthalpies H–H = 436, Cl–Cl = 243, H–Cl = 432 kJ mol−1, and state whether it is exothermic.
Calculate ΔrH for the complete combustion of ethene (C2H4 + 3O2 → 2CO2 + 2H2O) using C–H = 413, C=C = 614, O=O = 498, C=O = 805, O–H = 464 kJ mol−1. The experimentally measured value is −1411 kJ mol−1. Compare your answer with this and explain any difference.