Chem Sem 2 Exam Style Cards

[5 marks] Ethanol (CH3CH2OH) is soluble in water. With reference to intermolecular forces, explain why.

1) Both ethanol and water are polar covalent molecular substances. 2) Solute-solute: ethanol has hydrogen bonds, dipole-dipole and dispersion forces. 3) Solvent-solvent: water has hydrogen bonds, dipole-dipole and dispersion forces. 4) Solute-solvent: ethanol and water form hydrogen bonds, dipole-dipole and dispersion forces with each other. 5) Dissolves because the energy released forming the (similar strength) solute-solvent bonds is sufficient to compensate for the energy needed to break the solute-solute and solvent-solvent bonds.

[4 marks] Explain why carbon dioxide has low solubility in water.

1) CO2 is non-polar so has dispersion forces only between molecules. 2) Water has hydrogen bonds, dipole-dipole and dispersion forces. 3) Only weak dispersion forces can form between CO2 and water, so the energy released forming solute-solvent bonds is not enough to compensate for the energy needed to break water's strong hydrogen bonds. 4) Therefore CO2 does not dissolve to a significant extent.

[4 marks] Explain why iodine (I2) dissolves in hexane but not in water.

1) I2 and hexane are both non-polar, with only dispersion forces between particles. 2) Dispersion forces formed between I2 and hexane are similar in strength to those broken, so energy released compensates for energy absorbed and I2 dissolves. 3) Water has strong hydrogen bonds. 4) I2-water dispersion forces are not strong enough to break water's hydrogen bonds, so I2 does not dissolve in water.

[5 marks] Sodium chloride dissolves in water but not in hexane. Explain using intermolecular forces.

1) NaCl is an ionic lattice held by strong ionic bonds. 2) Water is polar and its molecules form strong ion-dipole attractions with Na+ and Cl-. 3) The energy released forming ion-dipole attractions is sufficient to break the ionic lattice and water's hydrogen bonds, so NaCl dissolves. 4) Hexane is non-polar and can only form weak dispersion forces with the ions. 5) These are not strong enough to overcome the ionic bonds, so NaCl is insoluble in hexane.

[3 marks] Oil and water do not mix. Explain.

1) Oil is non-polar with only dispersion forces between its molecules. 2) Water has strong hydrogen bonds between molecules. 3) Water-oil attractions (dispersion only) are too weak to compensate for breaking water's hydrogen bonds, so the two do not mix.

[8 marks] The boiling points of the Group 16 hydrides are H2O 100 °C, H2S -60 °C, H2Se -41 °C, H2Te -2 °C. Describe and explain the trend, referring to IMFs and the anomaly.

1) All are polar covalent molecular substances. 2) H2S, H2Se and H2Te have dipole-dipole and dispersion forces. 3) Boiling point increases H2S < H2Se < H2Te. 4) Molecules have an increasing number of electrons. 5) Dispersion forces get stronger, so the sum of IMFs increases and more energy is needed to overcome them. 6) H2O is an anomaly: its boiling point is far higher than expected. 7) H2O has hydrogen bonds (as well as dipole-dipole and dispersion). 8) Hydrogen bonding is much stronger than dipole-dipole for similar sized molecules, so the sum of IMFs is greatest and the boiling point anomalously high.

[6 marks] Explain why NH3 has a higher boiling point than PH3, but AsH3 is higher than PH3 as well.

1) NH3 has hydrogen bonds (N-H with lone pair on N), dipole-dipole and dispersion forces. 2) PH3 has dipole-dipole and dispersion forces only (P is not sufficiently electronegative and large). 3) Hydrogen bonds are much stronger than dipole-dipole for similar sized molecules, so NH3 has the greatest sum of IMFs despite having the fewest electrons. 4) From PH3 to AsH3 the number of electrons increases. 5) Dispersion forces therefore get stronger. 6) More energy is needed to overcome the larger sum of IMFs, so boiling point rises.

[6 marks] Butane and methylpropane (both C4H10) have boiling points of -0.5 °C and -11.7 °C. Match and explain.

1) Butane -0.5 °C; methylpropane -11.7 °C. 2) Both are non-polar so dispersion forces are the only IMF. 3) Both have the same number of electrons. 4) Butane is straight chain; methylpropane is branched. 5) Straight chains allow greater surface contact between molecules, so stronger dispersion forces (molecules can move closer, stronger induced dipoles). 6) More energy is needed to separate butane molecules so it has the higher boiling point.

[5 marks] Pentane boils at 36 °C and 2,2-dimethylpropane at 9.5 °C. Explain the difference.

1) Both are non-polar alkanes (C5H12) with only dispersion forces. 2) They have the same number of electrons. 3) Pentane is a long straight chain, 2,2-dimethylpropane is compact and nearly spherical. 4) Greater surface contact in pentane gives stronger dispersion forces. 5) More energy is needed to separate pentane molecules, so it has the higher boiling point.

[6 marks] Ethanol (C2H6O) boils at 78 °C and methoxymethane (CH3OCH3) at -24 °C. Explain.

1) Both have the same molecular formula and number of electrons, so similar dispersion forces. 2) Ethanol has an O-H bond so forms hydrogen bonds between molecules (plus dipole-dipole and dispersion). 3) Methoxymethane has no H bonded to O, so no hydrogen bonds; only dipole-dipole and dispersion forces. 4) Hydrogen bonding is stronger than dipole-dipole. 5) Ethanol has the greater sum of IMFs. 6) More energy is needed to separate ethanol molecules, so higher boiling point.

[4 marks] Sketch and describe a diagram showing a hydrogen bond between two water molecules.

1) Two bent H2O molecules, each with two lone pairs on O. 2) O is delta negative and each H is delta positive on both molecules. 3) Hydrogen bond drawn as a dashed line from the delta positive H of one molecule to a lone pair on the O (delta negative) of the other. 4) The hydrogen bond is labelled, with the O-H…O arrangement roughly in a line.

[3 marks] Why can HF form hydrogen bonds but HCl cannot?

1) Hydrogen bonding requires H bonded to N, O or F and a lone pair on a neighbouring N, O or F. 2) F is small and highly electronegative, making the H-F bond very polar and giving H a high charge density so it can approach the lone pair closely. 3) Cl is larger and less electronegative, so HCl only has dipole-dipole and dispersion forces.

[3 marks] State the IMFs present in methane, ethanol and carbon dioxide.

1) Methane (non-polar): dispersion forces only. 2) Ethanol (polar, O-H): hydrogen bonds, dipole-dipole and dispersion. 3) Carbon dioxide (non-polar, symmetrical): dispersion forces only.

[4 marks] State the strongest IMF between molecules of NH3, HCl, C2H6 and CH3Cl and justify each briefly.

1) NH3: hydrogen bonds (H on N, lone pair on N). 2) HCl: dipole-dipole (polar, no H-N/O/F). 3) C2H6: dispersion (non-polar). 4) CH3Cl: dipole-dipole (polar since the C-Cl bond differs from C-H so asymmetric).

[4 marks] Explain why water has a higher boiling point than H2S.

1) Water has hydrogen bonds, dipole-dipole and dispersion forces. 2) H2S has dipole-dipole and dispersion forces only. 3) The hydrogen bonds are much stronger than dipole-dipole forces. 4) The sum of IMFs is greater in water, so more energy is needed to overcome them and its boiling point is higher.

[4 marks] Explain why ice is less dense than liquid water.

1) In ice each water molecule forms hydrogen bonds in a fixed open hexagonal lattice. 2) This arrangement holds molecules further apart than in the liquid. 3) When ice melts some hydrogen bonds break and molecules move closer together. 4) Liquid water therefore has more molecules per unit volume (greater density), so ice floats.

[3 marks] Explain why water has a high surface tension.

1) Water molecules form hydrogen bonds with each other. 2) Molecules at the surface have no neighbours above so are pulled inward and sideways by hydrogen bonding. 3) This net inward attraction makes the surface behave like a skin that is hard to break.

[5 marks] Explain how you could tell, from temperature-drop data, which of two liquids has weaker IMFs, and why.

1) Evaporation is endothermic. 2) Particles need enough energy to overcome the IMFs holding them in the liquid. 3) The highest-energy particles escape, so the average kinetic energy (temperature) of the remaining liquid falls. 4) The liquid with the faster and larger temperature drop evaporates more readily. 5) It therefore has weaker IMFs.

[3 marks] Define polar bond and polar molecule.

1) Polar bond: a covalent bond between atoms of different electronegativity where the shared electrons are drawn closer to the more electronegative atom, giving partial charges (delta+ and delta-). 2) Polar molecule: a molecule with polar bonds that are arranged asymmetrically so the dipoles do not cancel. 3) It has an overall positive end and negative end.

[4 marks] CCl4 is non-polar but CHCl3 is polar. Explain.

1) Both contain polar C-Cl bonds. 2) CCl4 is tetrahedral with four identical bonds, so it is symmetrical and the bond dipoles cancel. 3) In CHCl3 the C-H bond is different from the three C-Cl bonds, so the molecule is asymmetrical. 4) The dipoles do not cancel, leaving a net dipole, so CHCl3 is polar.

[4 marks] CO2 is non-polar but SO2 is polar. Explain with shapes.

1) CO2 is linear (two electron areas, no lone pairs). 2) The two C=O dipoles are equal and opposite and cancel, so it is non-polar. 3) SO2 is bent because S has a lone pair (three electron areas). 4) The S=O dipoles do not cancel in a bent shape, so SO2 is polar.

[3 marks] State the shape and polarity of NH2Br and justify.

1) Pyramidal: N has three bonding pairs and one lone pair (four electron areas). 2) The N-H and N-Br bonds are polar and not identical. 3) The molecule is asymmetrical so the dipoles do not cancel; it is polar.

[3 marks] State the shape and polarity of CH2O (methanal) and justify.

1) Trigonal planar: C has three bonding areas (two C-H single bonds and one C=O double bond) and no lone pairs. 2) The C=O bond is polar and the C-H bonds are not equivalent to it. 3) Asymmetrical so dipoles do not cancel; polar.

[3 marks] State the shape of SO4 2- and PCl3 and justify.

1) SO4 2-: tetrahedral; S has four bonding areas and no lone pairs. 2) PCl3: pyramidal; P has three bonding pairs and one lone pair. 3) Lone pair repulsion compresses the PCl3 bond angle to slightly less than 109.5 degrees.

[3 marks] Why does a lone pair reduce bond angles, and why is any diatomic molecule linear?

1) A lone pair is held closer to the central atom and repels more strongly than a bonding pair, pushing bonding pairs closer together (angle below the ideal). 2) A diatomic molecule has only two nuclei, so the only possible arrangement is a straight line regardless of lone pairs.

[3 marks] State and explain the shape and polarity of SO3 and of H2S.

1) SO3: trigonal planar (three bonding areas, no lone pair), symmetrical so dipoles cancel; non-polar. 2) H2S: bent (two bonding pairs and two lone pairs), asymmetrical so dipoles do not cancel; polar.

[6 marks] Explain why SiO2 has a much higher melting point than CO2.

1) SiO2 is a covalent network substance: a 3D lattice of Si and O atoms joined by strong covalent bonds. 2) CO2 is covalent molecular: discrete molecules with covalent bonds inside and weak dispersion forces between molecules. 3) Melting SiO2 requires breaking strong covalent bonds throughout the lattice. 4) Melting CO2 only requires overcoming weak dispersion forces. 5) Far more energy is needed for SiO2. 6) Hence SiO2 has the much higher melting point.

[6 marks] Compare diamond and graphite: hardness, conductivity and why.

1) Diamond: each C bonded to four others in a 3D tetrahedral network. 2) Strong covalent bonds in all directions make it extremely hard with a very high melting point. 3) All valence electrons are in bonds so no mobile charge carriers; non-conductor. 4) Graphite: each C bonded to three others in hexagonal layers. 5) The fourth electron per carbon is delocalised and mobile, so graphite conducts. 6) Layers are held by weak dispersion forces so slide easily; graphite is soft and slippery (but has a high melting point because strong covalent bonds within layers must break).

[4 marks] Why is the melting point of C60 much lower than diamond?

1) C60 is a covalent molecular substance made of discrete spherical molecules. 2) Only weak dispersion forces between C60 molecules need to be overcome when melting. 3) Diamond is a covalent network, so strong covalent bonds throughout must be broken. 4) Therefore much less energy is needed for C60.

[4 marks] Define allotrope. How do allotropes differ from isotopes?

1) Allotropes are different structural forms of the same element where atoms are bonded or arranged differently (diamond, graphite). 2) They have significantly different physical properties but the same chemical composition. 3) Isotopes are atoms of the same element with the same number of protons but different neutrons. 4) Allotropes differ in structure; isotopes differ in mass number.

[4 marks] Explain why iron(III) oxide is brittle while iron is malleable.

1) Iron(III) oxide is ionic: a lattice of alternating cations and anions. 2) Iron is metallic: a lattice of cations in a sea of delocalised electrons. 3) When force distorts the ionic lattice, ions of the same charge are brought together and repel, so the crystal shatters. 4) When force distorts the metal lattice, the delocalised electrons move to maintain the attraction, so layers slide without breaking.

[3 marks] Cu conducts as a solid but Cu2O does not. Explain.

1) Cu is metallic with delocalised valence electrons that are free to move and carry charge. 2) Cu2O is an ionic lattice where ions are fixed in position. 3) Cu has mobile charge carriers in the solid state while Cu2O has none (comparative statement).

[4 marks] Explain why ionic compounds conduct when molten or dissolved but not as solids.

1) In the solid, ions are held in fixed lattice positions and cannot move. 2) Conduction requires mobile charged particles. 3) When molten, the lattice breaks down and ions are free to move. 4) When dissolved, the lattice separates into ions that are free to move in solution and carry charge.

[5 marks] Explain why aluminium has a higher melting point than sodium.

1) Both are metallic substances: lattice of cations in a sea of delocalised electrons. 2) Al contributes three delocalised electrons per atom, Na contributes one. 3) Al3+ has a greater charge (and the delocalised electron density is higher). 4) The electrostatic attraction between cations and delocalised electrons is therefore stronger in Al. 5) More energy is needed to overcome the metallic bonding, so Al has the higher melting point.

[8 marks] Describe and explain the trend in melting point across Period 3 (Na to Ar).

1) Na, Mg, Al are metallic; melting point increases Na < Mg < Al. 2) More delocalised electrons and higher cation charge give stronger electrostatic attraction. 3) More energy is needed to break the metallic bonding. 4) Si has the highest melting point; it is a covalent network. 5) Strong covalent bonds throughout the lattice must be broken. 6) P4, S8, Cl2 and Ar are covalent molecular (Ar is monatomic) with very low melting points. 7) Only weak dispersion forces between molecules are overcome. 8) Among these, S8 > P4 > Cl2 > Ar because of decreasing numbers of electrons per molecule, giving weaker dispersion forces.

[4 marks] Explain why melting point decreases down Group 1.

1) Atomic radius increases down the group (extra electron shells). 2) The delocalised valence electrons are further from the cation nuclei (more shielding). 3) The electrostatic attraction (metallic bonding) weakens. 4) Less energy is required to break the lattice, so melting point falls.

[3 marks] Why does mercury melt at -39 °C when most metals are solid?

1) Mercury's filled d-orbitals (electron configuration) make its metallic bonding unusually weak. 2) Less energy is required to overcome the attractions, so it has an unusually low melting point and is liquid at room temperature.

[4 marks] A solid has a high melting point, is hard and brittle, does not conduct as a solid but conducts when molten. Identify the bonding and justify.

1) Ionic. 2) High melting point and hardness come from strong electrostatic attractions in all directions in a 3D lattice. 3) Brittle because layers shifting align like charges which repel. 4) Does not conduct as a solid because ions are fixed, but conducts when molten as ions become mobile.

[4 marks] A solid melts at 119 °C and never conducts electricity. Identify the bonding type and explain.

1) Covalent molecular (such as S8). 2) Low melting point because only weak dispersion forces between molecules are overcome. 3) Does not conduct because molecules are neutral and all electrons are held in bonds or lone pairs. 4) No mobile charged particles in any state.

[4 marks] Explain why the melting point of metals can be explained using the metallic bonding model.

1) Metals are cations in a lattice surrounded by delocalised valence electrons. 2) There is strong electrostatic attraction between cations and electrons throughout the lattice. 3) A large amount of energy is needed to overcome this. 4) So metals generally have high melting points.

[4 marks] Define electronegativity and explain the trend across Period 2.

1) Electronegativity is the tendency of an atom to attract a bonding pair of electrons towards itself. 2) It increases from Li to F. 3) Nuclear (core) charge increases across the period while bonding electrons remain in the same shell. 4) Bonding electrons are attracted more strongly.

[3 marks] Explain why electronegativity decreases down a group.

1) Atomic radius increases (more shells), so the bonding pair is further from the nucleus. 2) Inner shells shield the bonding electrons from the nuclear charge. 3) Attraction for the bonding pair is weaker, so electronegativity decreases.

[3 marks] Predict the electronegativity of neon and justify.

1) A value of zero (or lower than lithium). 2) Neon is monatomic with a full outer shell and does not generally form covalent bonds, so it does not attract a bonding pair.

[4 marks] Explain the trend in atomic radius across Period 3 and down Group 2.

1) Across: radius decreases. 2) Same number of shells but increasing nuclear charge; greater attraction pulls valence electrons closer. 3) Down: radius increases. 4) An extra shell is added each time, placing valence electrons further away with more shielding by inner shells.

[3 marks] Explain why the first ionisation energy of rubidium is lower than that of sodium.

1) Rb has a larger atomic radius, so the outer electron is further from the nucleus. 2) Core charge is the same (protons and shielding both increase). 3) Attraction is weaker, so less energy is needed to remove the outer electron.

[3 marks] Explain why argon has a higher first ionisation energy than chlorine, and why potassium has a lower one than argon.

1) Ar has one more proton than Cl and valence electrons in the same shell, so stronger attraction and higher energy needed. 2) K's outer electron is in a new, higher shell much further from the nucleus. 3) Attraction is weaker despite the extra proton, so IE is lower.

[3 marks] Successive ionisation energies of an element show a large jump between the 4th and 5th. What does it show and why?

1) The element has 4 valence electrons (such as carbon, Group 14). 2) After four are removed the fifth must come from an inner shell. 3) This shell is much closer to the nucleus with a much stronger attraction, so energy needed jumps.

[2 marks] Why is the second ionisation energy always larger than the first?

1) After one electron is removed the ion is positive. 2) Removing an electron from a positive ion needs more energy because of the greater electrostatic attraction (fewer electrons, same protons).

[3 marks] Why do isotopes of an element have the same chemical properties but different physical properties?

1) Isotopes have the same number of protons and electrons, so the same electron configuration and valence electrons. 2) Chemical properties depend on valence electrons. 3) They differ in neutrons and mass, so physical properties such as density differ.

[4 marks] Describe how a mass spectrometer works and explain why only positive ions are detected.

1) Ionisation: sample bombarded by high-energy electrons forming positive ions. 2) Acceleration: ions accelerated towards a negatively charged plate. 3) Deflection and detection: magnetic field deflects ions according to mass/charge; lighter or more highly charged ions deflect more. 4) Negative ions are repelled by the negative plate and never reach the detector.

[5 marks] Describe how Rutherford's gold foil experiment led to the nuclear model.

1) Alpha particles were fired at thin gold foil. 2) Most passed straight through, so atoms are mostly empty space. 3) A few were deflected, so there is a concentrated positive charge. 4) About 1 in 8000 bounced straight back, so the positive charge and most mass are in a tiny dense nucleus. 5) This disproved Thomson's plum pudding model (spread-out positive charge).

[4 marks] Explain how Bohr's model accounts for line emission spectra, and state one limitation.

1) Electrons occupy fixed energy levels (shells). 2) Electrons absorb specific energy to jump to a higher shell and release it as a photon when falling back. 3) Energy gaps are fixed, so only specific wavelengths (lines) are seen. 4) Limitation: only accurate for hydrogen, it cannot explain multi-electron atoms.

[3 marks] How did Chadwick's discovery improve the atomic model?

1) He bombarded beryllium with alpha particles and detected neutral radiation of significant mass (neutrons). 2) This showed the nucleus contains neutrons as well as protons. 3) It explained why atomic mass exceeds proton mass alone (e.g. helium).

[3 marks] Why can flame tests and AAS identify elements?

1) Electrons absorb energy and move to higher levels. 2) They emit/absorb energy at specific wavelengths when changing levels. 3) The line spectrum is unique for each element.

[3 marks] Why do nanomaterials have properties different from the bulk material, and why are they regulated?

1) Particles of 1-100 nm have a very high surface area to volume ratio. 2) Properties change with size. 3) Different properties can pose new risks to health, safety and the environment, so regulation is required.

[5 marks] Explain the difference between an alkene reacting with bromine and an alkane reacting with bromine.

1) Alkene: addition; bromine adds across the C=C. 2) No UV needed; orange bromine decolourises rapidly. 3) Alkane: substitution; one H is replaced by Br and HBr forms. 4) Requires UV light; without UV there is no reaction (orange persists). 5) Alkenes are reactive because the double bond can open; alkanes have only strong single bonds.

[4 marks] Describe how to distinguish hex-1-ene from cyclohexane, both C6H12.

1) Add bromine water to each in the dark. 2) Hex-1-ene: orange solution rapidly decolourises (colourless), addition at C=C. 3) Cyclohexane: orange colour remains, no reaction without UV. 4) Hex-1-ene is unsaturated; cyclohexane is saturated.

[4 marks] Explain why benzene does not decolourise bromine water although it is unsaturated-looking.

1) Benzene has a planar ring with six delocalised electrons. 2) Delocalisation makes the ring unusually stable. 3) Addition would destroy this stability so benzene resists it. 4) It reacts only by substitution with a halogen carrier catalyst (e.g. AlBr3).

[4 marks] Explain why cis/trans isomers exist for but-2-ene but not but-1-ene.

1) Rotation about a C=C is restricted. 2) Cis/trans isomers need each carbon of the double bond to have two different groups. 3) In but-2-ene each double bond carbon has an H and a CH3. 4) In but-1-ene one double bond carbon has two H, so swapping gives the same molecule.

[3 marks] Explain structural isomers using C5H12.

1) Structural isomers have the same molecular formula but different arrangements of atoms. 2) C5H12 has pentane, 2-methylbutane and 2,2-dimethylpropane. 3) They differ in branching so physical properties (boiling point) differ.

[5 marks] Explain why fractional distillation separates crude oil.

1) Crude oil is heated until it vaporises. 2) Vapours rise up a column with a temperature gradient (hot bottom, cooler top). 3) Longer chains have more electrons and stronger dispersion forces, so higher boiling points. 4) They condense lower in the column. 5) Short chains with weaker dispersion forces condense higher up and each fraction collects where the temperature matches its boiling range.

[4 marks] Distinguish complete from incomplete combustion and give one danger of the latter.

1) Complete combustion (excess O2) forms CO2 and H2O only. 2) Incomplete (limited O2) forms CO and/or C (soot) and H2O. 3) CO is toxic. 4) It binds to haemoglobin preventing oxygen transport.

[4 marks] Compare fossil fuels and biofuels.

1) Fossil fuels formed over millions of years so are non-renewable; biofuels come from recent plants and are renewable. 2) Fossil fuels have higher energy content per gram (octane 47.8 kJ/g vs ethanol 29.7). 3) Biofuels have lower net CO2 because CO2 emitted was recently absorbed by plants. 4) Biofuels need farmland and processing energy and have lower energy density.

[3 marks] How do catalytic converters help the environment?

1) They contain catalysts (e.g. metal nanoparticles). 2) They convert CO, unburnt hydrocarbons and nitrogen oxides into CO2, H2O and N2. 3) This lowers emissions of toxic pollutants.

[4 marks] Explain the difference between a strong acid and a weak acid, and between concentrated and dilute.

1) Strong acid ionises completely in water; weak acid only partially ionises (equilibrium, mostly molecules). 2) Concentrated means a large amount of solute per litre; dilute means a small amount. 3) Strength is about degree of ionisation; concentration is about amount per volume. 4) They are independent: a dilute solution can be a strong acid, a concentrated solution a weak acid.

[5 marks] Explain why 0.10 M ethanoic acid has a higher pH than 0.10 M HCl, with equations.

1) HCl + H2O → H3O+ + Cl- (complete). 2) CH3COOH + H2O ⇌ CH3COO- + H3O+ (partial). 3) Ethanoic acid only partially ionises so only a small proportion forms H3O+. 4) [H3O+] is lower than in HCl. 5) pH = -log[H+], so lower [H+] means higher pH.

[3 marks] Why does 0.01 M H2SO4 have a lower pH than 0.01 M HCl?

1) H2SO4 is a strong diprotic acid that ionises completely for the first H+. 2) Each mole can release up to two H+ (HSO4- ionises further). 3) [H+] is about twice that of HCl, so pH is lower.

[4 marks] Can a concentrated weak acid conduct better than a dilute strong acid? Explain.

1) Yes. 2) Conductivity depends on the concentration of ions. 3) A concentrated weak acid still ionises partially, giving a significant number of ions. 4) A very dilute strong acid has fully ionised but few ions overall.

[4 marks] Equal volumes of 1.0 M HCl and 1.0 M ethanoic acid react with excess magnesium. Compare the rate and the total volume of hydrogen.

1) HCl reacts faster. 2) HCl has a higher [H+], so more frequent collisions with Mg. 3) The total volume of H2 is the same. 4) Both contain the same moles of acid; as H+ is consumed more ethanoic acid ionises, so all of it eventually reacts.

[4 marks] Explain the Bronsted-Lowry model using NH3 + H2O ⇌ NH4+ + OH-.

1) A B-L acid is a proton donor; a B-L base is a proton acceptor. 2) NH3 accepts H+ from water so NH3 is the base. 3) H2O donates H+ so it is the acid. 4) NH4+ is the conjugate acid of NH3 and OH- is the conjugate base of H2O.

[3 marks] Why is the Arrhenius model limited? Give an example.

1) It only applies to aqueous solutions. 2) It requires bases to contain OH-. 3) Ammonia acts as a base (accepts H+) but has no OH-, which Bronsted-Lowry explains.

[4 marks] Explain why HCO3- is amphiprotic, with equations.

1) An amphiprotic species can both donate and accept a proton. 2) As an acid: HCO3- + H2O ⇌ CO3 2- + H3O+. 3) As a base: HCO3- + H2O ⇌ H2CO3 + OH-. 4) Therefore it acts as both.

[3 marks] Why is pure water a weak electrolyte and why is neutral not always pH 7?

1) Water self-ionises: H2O + H2O ⇌ H3O+ + OH-, but only to a tiny extent. 2) Very few ions are present so conductivity is very low. 3) Kw changes with temperature, so the pH of neutral water changes (pH 7 only at 25 °C).

[4 marks] Describe the observations when copper(II) oxide is added to excess nitric acid and the reaction type.

1) A black solid is added to a colourless solution. 2) The solid is consumed forming a blue solution. 3) Acid + metal oxide → salt + water. 4) Ionic equation: CuO + 2H+ → Cu2+ + H2O.

[4 marks] Describe the observations when solid ammonium carbonate is added to excess hydrochloric acid.

1) A white solid is added to a colourless solution. 2) Bubbles form (CO2). 3) The solid dissolves/is consumed to give a colourless solution. 4) The reaction is carbonate + acid → salt + water + carbon dioxide.

[3 marks] Why is a solution of sodium carbonate basic?

1) Carbonate ions are a weak base. 2) CO3 2- + H2O ⇌ HCO3- + OH-. 3) The OH- produced makes [OH-] greater than [H+], so pH is above 7.

[4 marks] Explain why sulfur dioxide and sodium oxide give different universal indicator colours in water.

1) SO2 is a non-metal oxide: SO2 + H2O → H2SO3 (acidic), so UI is red/pink. 2) Na2O is a metal oxide: Na2O + H2O → 2NaOH (strongly basic), so UI is blue/violet. 3) Non-metal oxides are acidic, metal oxides basic.

[4 marks] Describe a test to distinguish hydrochloric acid and nitric acid solutions of equal concentration.

1) Add silver nitrate solution to each. 2) HCl gives a white precipitate: Ag+ + Cl- → AgCl. 3) HNO3 gives no precipitate (no reaction). 4) Ag+ only precipitates with the chloride ion.

[4 marks] Describe how to distinguish solid zinc hydroxide from solid barium sulfate.

1) Add dilute acid (HCl) to each. 2) Zn(OH)2 dissolves forming a colourless solution: Zn(OH)2 + 2H+ → Zn2+ + 2H2O. 3) BaSO4 does not react or dissolve. 4) Insoluble solids that dissolve in acid have reacted with it.

[3 marks] Why are weak acids written in molecular form in ionic equations?

1) Weak acids only partially ionise so exist mainly as molecules in solution. 2) Only strong electrolytes (strong acids and soluble ionic compounds) are written as ions. 3) Spectator ions are removed.

[3 marks] Define limiting reagent and explain why it determines the amount of product.

1) The limiting reagent is the reactant that is completely consumed first. 2) Once it is used up the reaction stops. 3) Other reactants are in excess, so the amount of product is set by the limiting reagent.

[4 marks] Explain why doubling the amount of one reactant does not always double the product.

1) Product amount is limited by whichever reactant runs out first. 2) If the reactant that was doubled is already in excess, the limiting reagent is unchanged. 3) Product will not increase until the limiting reagent increases. 4) Excess reactant is left over.

[3 marks] Describe how you could observe that magnesium was the limiting reagent in a reaction with acid.

1) All the magnesium solid is consumed. 2) Bubbling stops even though acid remains. 3) Adding more Mg restarts the bubbling (acid in excess).

[4 marks] Why is one reactant often used in excess in industry?

1) To ensure the more expensive or valuable reactant is used up completely. 2) Increases the yield based on that reactant. 3) Helps the reaction go to completion at a faster rate. 4) The excess can be recovered or recycled.

[4 marks] Give reasons why the actual yield is lower than the theoretical yield.

1) Reaction does not go to completion (reversible). 2) Side reactions form other products. 3) Product lost in transfers, filtration or left in solution. 4) Impure reactants or inefficient steps.

[3 marks] Why can a percentage yield greater than 100% occur?

1) The product is impure (contains reactant or by-product). 2) The product is wet and not fully dried. 3) Measurement or mass error.

[4 marks] Explain why the overall yield of a multi-step process is lower than each individual step efficiency.

1) The product of step 1 is the reactant for step 2. 2) Each step loses some material. 3) Losses compound since efficiencies multiply. 4) Overall yield is the product of step efficiencies, lower than any single step.

[4 marks] In a cyclic process some excess reactant remains after the first cycle. Explain why it matters for the second cycle.

1) Unreacted excess reactant carries over. 2) It adds to the amount regenerated from the cycle. 3) The relative amounts change, so a different reactant can become limiting. 4) Therefore the limiting reagent must be re-evaluated and the yield recalculated.

[4 marks] After neutralising an acid with a base, the final pH is not 7. Explain what determines it.

1) Whichever reagent is in excess determines pH. 2) Excess H+ gives pH below 7. 3) Excess OH- gives pH above 7. 4) Exactly equal moles (strong acid and strong base) gives pH 7.

[3 marks] A student assumed a precipitate was pure silver chloride but it also contained lead chloride. State the type of error and effect.

1) Systematic error. 2) Pb2+ also precipitates with chloride, adding mass. 3) The calculated %Ag is always too high.

[6 marks] Explain, using collision theory and a Maxwell-Boltzmann distribution, why a reaction is faster at higher temperature.

1) Higher temperature increases average kinetic energy. 2) Particles move faster so the frequency of collisions increases. 3) The proportion of particles with energy at least Ea increases. 4) Therefore the frequency of successful collisions increases and rate increases. 5) Diagram: higher temperature curve is lower and shifted right with the same area. 6) Ea is marked and the area beyond Ea is larger.

[4 marks] Explain, using collision theory, why increasing concentration and surface area increase rate.

1) Concentration: more particles per unit volume. 2) More frequent collisions and so more successful collisions per second. 3) Surface area: more reactant particles exposed. 4) More frequent collisions with the other reactant so rate increases.

[4 marks] Explain how a catalyst increases reaction rate without changing the enthalpy change.

1) It provides an alternative pathway with lower activation energy. 2) A greater proportion of collisions have energy at least Ea. 3) More frequent successful collisions. 4) Reactants and products have the same energy so deltaH is unchanged; the catalyst is not consumed.

[3 marks] Why is the initial rate fastest, and why does a gas volume graph level off?

1) At the start reactant concentration is highest so collisions are most frequent. 2) As reactants are consumed concentration falls and collisions become less frequent. 3) When the limiting reagent is used up the reaction stops and the graph is horizontal.

[4 marks] Why does a small rise in temperature have a large effect on rate?

1) Particles have a spread of energies (Maxwell-Boltzmann). 2) Ea lies on the high-energy tail of the curve. 3) A small shift greatly increases the proportion of particles with energy at least Ea. 4) So successful collisions increase disproportionately.

[4 marks] Thermite mixture is stable until ignited with a magnesium coil. Explain.

1) The reaction has a high activation energy. 2) At room temperature few particles have energy of at least Ea. 3) Heating from Mg combustion raises particle energy. 4) Once started the exothermic reaction releases enough heat to sustain it.

[5 marks] Describe and justify an energy profile diagram for an exothermic catalysed reaction.

1) Axes: enthalpy vs reaction coordinate. 2) Reactants higher than products. 3) deltaH shown as a downward arrow (negative). 4) Ea from reactants to peak of the activated complex. 5) Catalysed pathway has a lower peak but the same start and end levels.

[4 marks] Explain why boiling water is endothermic and compare potential energy of liquid and gas.

1) Heat is absorbed to break intermolecular (hydrogen) bonds between molecules. 2) No covalent bonds are broken. 3) Gaseous molecules have higher potential energy. 4) Energy was absorbed from the surroundings.

[3 marks] Explain bond breaking and bond forming in terms of energy.

1) Breaking bonds absorbs energy (endothermic). 2) Forming bonds releases energy (exothermic). 3) If more energy is released forming bonds than absorbed breaking bonds, the reaction is exothermic (deltaH negative).

[5 marks] Write a hypothesis, identify the variables and explain why controlled variables matter for Mg reacting with acid at different temperatures.

1) If temperature increases then the rate of reaction increases because collisions are more frequent and energetic. 2) Independent: temperature. 3) Dependent: rate (e.g. gas volume per time). 4) Controlled: acid concentration and volume, mass of Mg. 5) Controlling variables ensures only temperature affects rate, improving validity.