Chemical energetics Banker

1. (a) Which statement is correct for the energy profile of an exothermic reaction?A The products are at a higher energy level than the reactants.B The products are at a lower energy level than the reactants.C The reactants and products are at the same energy level.D There is no activation energy. [1]

1. (a) Define activation energy, Ea. [1]
2. (b) Describe the energy profile for an endothermic reaction. Say where the reactants and products are, and show how Ea and ΔH are measured. [3]

1. (a) On an energy profile, the reactants are at 120 kJ/mol, the top of the curve is at 290 kJ/mol and the products are at 40 kJ/mol. What are the activation energy and the enthalpy change, ΔH?A activation energy 170 kJ/mol, ΔH −80 kJ/molB activation energy 250 kJ/mol, ΔH −80 kJ/molC activation energy 170 kJ/mol, ΔH +80 kJ/molD activation energy 290 kJ/mol, ΔH −40 kJ/mol [1]

A reaction has ΔH = +58 kJ/mol and activation energy Ea = 140 kJ/mol.

1. (a) Is the reaction exothermic or endothermic? Give a reason. [1]
2. (b) Calculate the energy of the top of the curve relative to the products. [2]
3. (c) State the activation energy of the reverse reaction. [1]

Many homes in Pakistan cook on piped natural gas (mainly methane).

1. (a) The burning of natural gas is exothermic, yet a gas stove needs a spark or match to light it. Explain why. [2]
2. (b) Explain why the flame keeps burning once it has been lit. [1]

Hydrogen peroxide decomposes slowly to water and oxygen. When a little manganese(IV) oxide is added, the reaction becomes much faster.

1. (a) Describe how the energy profile for the catalysed reaction differs from that for the uncatalysed reaction. [2]
2. (b) Hydrogen peroxide decomposition is exothermic. State where the products are on the energy profile. [1]
3. (c) State what happens to the mass of manganese(IV) oxide during the reaction. [1]

1. (a) A catalyst is added to a reaction. What happens to the activation energy and to ΔH?A activation energy decreases, ΔH decreasesB activation energy decreases, ΔH unchangedC activation energy unchanged, ΔH decreasesD activation energy increases, ΔH unchanged [1]

An energy profile is described by these points (energy in kJ/mol):reactants 50; top of curve 230; products 180.

1. (a) Calculate ΔH for the forward reaction. [1]
2. (b) Calculate the activation energy of the uncatalysed forward reaction. [1]
3. (c) With a catalyst the peak is at 150 kJ/mol. Calculate the new activation energy. [1]
4. (d) Calculate the activation energy of the uncatalysed reverse reaction. [1]
5. (e) Explain why the catalysed reaction is faster. [2]

Data for four reactions:Reaction P: Ea 95 kJ/mol, ΔH −210 kJ/molReaction Q: Ea 40 kJ/mol, ΔH −60 kJ/molReaction R: Ea 180 kJ/mol, ΔH −390 kJ/molReaction S: Ea 120 kJ/mol, ΔH +45 kJ/mol

1. (a) Which reaction would be fastest at room temperature? Explain. [2]
2. (b) Which reaction is endothermic? [1]
3. (c) For reaction P, calculate the activation energy of the reverse reaction. [1]
4. (d) Which reaction gives out the most energy when 1 mole reacts? [1]

1. (a) Which process is endothermic?A burning methaneB neutralising an acid with an alkaliC thermal decomposition of calcium carbonateD respiration [1]

A student mixed 25 cm³ of dilute hydrochloric acid with 25 cm³ of aqueous sodium hydroxide in a polystyrene cup. Initial temperature of both solutions = 22.5 °C; highest temperature reached = 29.0 °C.

1. (a) Calculate the temperature rise. [1]
2. (b) State whether the reaction is exothermic or endothermic. Give a reason. [1]
3. (c) Suggest why a polystyrene cup is used instead of a glass beaker. [1]
4. (d) Suggest one further improvement to make the temperature rise more accurate. [1]

1. (a) Classify each change as exothermic or endothermic.(i) burning petrol in a car engine(ii) electrolysis of molten aluminium oxide(iii) adding magnesium to dilute hydrochloric acid(iv) melting ice [4]
2. (b) Define an exothermic reaction in terms of energy transfer. [1]

1. (a) When ammonium nitrate dissolves in water, the temperature of the solution falls from 21 °C to 13 °C. Which statement explains this?A The process is exothermic; energy is transferred to the surroundings.B The process is endothermic; energy is transferred from the surroundings.C The process is endothermic; energy is transferred to the surroundings.D The process is exothermic; energy is transferred from the surroundings. [1]

A student dissolved 5.0 g of each salt in 50 cm³ of water and recorded the temperatures.Ammonium chloride: start 21.0 °C, end 16.5 °CCalcium chloride (anhydrous): start 21.0 °C, end 32.0 °CPotassium nitrate: start 21.5 °C, end 15.5 °CSodium chloride: start 21.0 °C, end 20.5 °C

1. (a) Calculate the temperature change for each salt. Use + for a rise and − for a fall. [2]
2. (b) Which salt dissolves in the most endothermic process? [1]
3. (c) Which salt would be most suitable for a self-heating food can? Explain. [1]
4. (d) Give two variables that should be kept constant. [2]

1. (a) A physiotherapist at a cricket match uses an instant cold pack. Squeezing it lets ammonium nitrate dissolve in water. Explain how this cools the injury. [2]
2. (b) A hand warmer contains iron powder that reacts with oxygen from the air. State whether this reaction is exothermic or endothermic, and describe the energy transfer. [2]
3. (c) Suggest why the iron in the hand warmer is a fine powder rather than a lump. [1]

Excess zinc powder was added to 50 cm³ of aqueous copper(II) sulfate in an insulated cup at time 0. The temperature was recorded every 30 seconds.Time/min: 0.0, 0.5, 1.0, 1.5, 2.0, 2.5, 3.0, 3.5Temperature/°C: 21.0, 27.5, 32.0, 34.0, 34.5, 34.0, 33.5, 33.0

1. (a) Calculate the maximum temperature rise. [1]
2. (b) Explain why the temperature falls after 2.0 minutes. [1]
3. (c) Write the ionic equation for the reaction. [1]
4. (d) The experiment is repeated with magnesium powder instead of zinc, using the same number of moles of metal. Predict how the maximum temperature rise would compare. Explain. [2]
5. (e) Suggest why zinc was used in excess. [1]

Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O ΔH = −2800 kJ/mol

1. (a) State the value of ΔH for photosynthesis, 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂, and say whether it is exothermic or endothermic. [2]
2. (b) Calculate the energy released per gram of glucose in respiration. [Mr glucose = 180] Give your answer to 3 significant figures. [1]
3. (c) Describe the energy level diagram for respiration. [2]
4. (d) State where the energy for photosynthesis comes from. [1]

At the high temperatures inside a car engine, nitrogen reacts with oxygen: N≡N + O=O → 2 N=OBond energies in kJ/mol: N≡N 945, O=O 498, N=O 630

1. (a) Calculate the energy needed to break the bonds in the reactants. [1]
2. (b) Calculate the energy released when the bonds in the products form. [1]
3. (c) Calculate the energy change, ΔH, for the reaction and state whether it is exothermic or endothermic. [2]

In the Haber process: N≡N + 3 H–H ⇌ 2NH₃ (each NH₃ has three N–H bonds)Bond energies in kJ/mol: N≡N 945, H–H 436, N–H 391

1. (a) Calculate the energy change, ΔH, for the reaction in kJ/mol. [3]
2. (b) State the sign of ΔH for the reverse reaction, the decomposition of ammonia. [1]

1. (a) Which statement about bond energies is correct?A Breaking bonds releases energy and making bonds takes in energy.B Breaking bonds takes in energy and making bonds releases energy.C Both breaking and making bonds release energy.D Both breaking and making bonds take in energy. [1]

For a reaction, the total energy taken in to break bonds is 1850 kJ/mol and the total energy released in making bonds is 2100 kJ/mol.

1. (a) Calculate ΔH for the reaction. [1]
2. (b) State whether the reaction is exothermic or endothermic and explain your answer. [2]
3. (c) State what happens to the temperature of the surroundings during the reaction. [1]

Natural gas (methane) burns in a gas stove:CH₄ + 2O₂ → CO₂ + 2H₂OBond energies in kJ/mol: C–H 413, O=O 498, C=O 805, O–H 464

1. (a) Calculate the total energy needed to break all the bonds in the reactants. [1]
2. (b) Calculate the total energy released when the bonds in the products form. [1]
3. (c) Calculate the energy change for the reaction, in kJ/mol. [1]
4. (d) Explain, in terms of bond breaking and bond making, why this reaction is exothermic. [2]

1. (a) Hydrogen iodide decomposes: 2HI → H₂ + I₂Bond energies in kJ/mol: H–I 299, H–H 436, I–I 151What is the energy change for this reaction?A −11 kJ/molB +11 kJ/molC +148 kJ/molD −288 kJ/mol [1]

A city plans hydrogen-powered buses. Hydrogen burns in oxygen to form water.Bond energies in kJ/mol: H–H 436, O=O 498, O–H 464

1. (a) Calculate ΔH for the reaction 2H₂ + O₂ → 2H₂O in kJ/mol. [3]
2. (b) Calculate the energy released per mole of hydrogen, H₂, burned. [1]
3. (c) State one advantage of hydrogen as a fuel for buses. [1]

Methane reacts with chlorine: CH₄ + Cl₂ → CH₃Cl + HClBond energies in kJ/mol: C–H 413, Cl–Cl 242, C–Cl 327, H–Cl 431

1. (a) Calculate ΔH for this reaction in kJ/mol. [3]
2. (b) State the condition needed for this reaction to happen. [1]

Bond energies in kJ/mol: H–H 436, F–F 158, H–F 568, Cl–Cl 242, H–Cl 431

1. (a) Calculate ΔH for H₂ + F₂ → 2HF in kJ/mol. [2]
2. (b) Calculate ΔH for H₂ + Cl₂ → 2HCl in kJ/mol. [2]
3. (c) Which reaction releases more energy? Use the bond energies to explain why. [2]

Ethene reacts with hydrogen using a nickel catalyst: CH₂=CH₂ + H–H → CH₃–CH₃Bond energies in kJ/mol: C=C 612, C–C 347, C–H 413, H–H 436

1. (a) Calculate ΔH for the reaction, in kJ/mol. [3]
2. (b) Suggest why bond-energy calculations give only an approximate value of ΔH. [1]

H₂ + Br₂ → 2HBr ΔH = −103 kJ/molBond energies in kJ/mol: H–H 436, Br–Br 193

1. (a) Calculate the bond energy of the H–Br bond. [3]
2. (b) Is the H–Br bond stronger or weaker than the H–Cl bond (431 kJ/mol)? Explain. [1]

Ethanol, CH₃CH₂OH, is used as a fuel in some countries:C₂H₅OH + 3O₂ → 2CO₂ + 3H₂OEthanol contains 5 C–H, 1 C–C, 1 C–O and 1 O–H bond.Bond energies in kJ/mol: C–H 413, C–C 347, C–O 358, O–H 464, O=O 498, C=O 805

1. (a) Calculate the energy needed to break all the bonds in one mole of ethanol and three moles of oxygen. [2]
2. (b) Calculate ΔH for the combustion of ethanol in kJ/mol. [2]
3. (c) Calculate the energy released per gram of ethanol. [Mr ethanol = 46] Give your answer to 3 significant figures. [1]

Hydrogen peroxide decomposes: 2H–O–O–H → 2H–O–H + O=OBond energies in kJ/mol: O–H 464, O–O 146, O=O 498

1. (a) Calculate ΔH for this reaction in kJ/mol. [3]
2. (b) The value measured by experiment is −196 kJ/mol. Suggest why the calculated value is different. [1]
3. (c) Manganese(IV) oxide speeds up this reaction. State its effect on the value of ΔH. [1]