Year 11 Chemistry Module 4: Drivers of Reactions Practice Questions

Question 6 ⭐

What is the bond enthalpy of hydrogen gas? (2 marks)

(L2.1: explain the enthalpy changes in a reaction in terms of breaking and reforming bonds, and relate this to the law of conservation of energy)

Question 7 ⭐⭐

With reference to the law of conservation of energy, explain why the dissociation of ionic substances in aqueous solutions are always endothermic on a molecular level. (3 marks)

(L2.1: explain the enthalpy changes in a reaction in terms of breaking and reforming bonds, and relate this to the law of conservation of energy)

Question 8 ⭐⭐⭐

Making reference to the law of conservation of energy, explain the enthalpy changes for the following reactions. (3 marks)

(L2.1: explain the enthalpy changes in a reaction in terms of breaking and reforming bonds, and relate this to the law of conservation of energy)

Question 9 ⭐⭐

 Calculate the reaction enthalpy for the formation of anhydrous aluminium chloride. (3 marks)

 

from the following data:

Question sourced from ChemTeam

(L2.2:investigate Hess’s Law in quantifying the enthalpy change for a stepped reaction using standard enthalpy change data and bond energy data, for example: carbon reacting with oxygen to form carbon dioxide via carbon monoxide)

Question 10 ⭐⭐⭐

Bombardier beetles can fire a hot, toxic mixture of chemicals at their attacker. This mixture contains quinone, C₆H₄O₂, a compound formed by the reaction of hydroquinone C₆H₄(OH)₂ with hydrogen peroxide H₂O₂. (4 marks)

The equation for the overall reaction is:

\text{C}_6\text{H}_4(\text{OH})_{2(aq)} + \text{H}_2\text{O}_{2(aq)} \rightarrow \text{C}_6\text{H}_4\text{O}_{2(aq)} + 2\text{H}_2\text{O}_{(l)}

Using the following data, calculate the enthalpy change for the above reaction.

(L2.2:investigate Hess’s Law in quantifying the enthalpy change for a stepped reaction using standard enthalpy change data and bond energy data, for example: carbon reacting with oxygen to form carbon dioxide via carbon monoxide)

Question 11 ⭐⭐

Using the following information: 

a) Calculate the total enthalpy change of photosynthesis. (3 marks)

b) State what the total enthalpy change of a respiration reaction is and justify your answer. (1 mark)

(L2.3: apply Hess’s Law to simple energy cycles and solve problems to quantify enthalpy changes within reactions, including but not limited to enthalpy changes involved in photosynthesis)

Entropy and Gibbs Free Energy 

Question 12 ⭐⭐⭐

A student conducts an experiment in which solid ammonium chloride (NH₄Cl) is dissolved in water, and the temperature of the solution decreases. 

a) Describe the changes in enthalpy and entropy during the dissolution of ammonium chloride. (2 marks)

b) Based on these changes, explain why the dissolution occurs spontaneously despite the temperature decreasing. (2 marks)

c) Discuss whether entropy or enthalpy is the dominant driving force for this process and justify your answer using chemical reasoning. (2 marks)

(L3.1: analyse the differences between entropy and enthalpy)

Question 13 ⭐

A student carried out the following experiment: 

1. Fill a beaker up with 200mL of water 
2. Increase the temperature of the water such that it is boiling 
3. Record observations 

Justify whether this is an appropriate example of a model illustrating increasing entropy. (2 marks)

(L3.2: use modelling to illustrate entropy changes in reactions)

Question 14 ⭐⭐

For the following reactions, identify whether it is increasing or decreasing in entropy and explain why. (4 marks)

(L3.3: predict entropy changes from balanced chemical reactions to classify as increasing or decreasing entropy)

Question 15 ⭐⭐⭐

Using the Gibbs free energy formula, deduce whether the reaction would be spontaneous in situation 1,2, 3 and 4 in the table below. (8 marks)

(L3.4: explain reaction spontaneity using terminology, including: Gibbs free energy, enthalpy, entropy)

Question 16 ⭐⭐

Given the reaction of diamond converting to graphite is 2C_{(s) diamond}→2C_{(s) graphite} determine ∆G at 298 K and determine if this reaction is spontaneous or not. What does ∆G say about the rate of this reaction? (3 marks)

∆H°f(C_(s)diamond) =1.9kJ/mol 

S°(C_(s)diamond =2.38J/(molK) 

S°(C_(s)graphite)=5.74J/(molK)

Question sourced from Chemistry LibreTexts

(L3.5: solve problems using standard references and the Gibbs free energy formula to classify reactions as spontaneous or nonspontaneous)

Question 17 ⭐

With reference to the Gibbs Free energy, explain why some reactions that go forward are not considered spontaneous reactions. (2 marks) 

(L3.6: predict the effect of temperature changes on spontaneity)

Question 18 ⭐⭐⭐🔥

A student conducts a reaction where 2.0g of magnesium reacts with excess hydrochloric acid in 100.0g of water. The temperature increases from 25.0 C to 32.5 C. Assume the heat capacity of the solution is 4.18 J/gC and no heat is lost to the surroundings.

a) Calculate the enthalpy change for the reaction per mole of magnesium. (3 marks)

b) If the reaction was conducted in an open system where heat was lost to the surroundings, explain how the calculated value of ΔH would differ and why. (2 marks)

c) Explain how the law of conservation of energy applies in this reaction. (2 marks)

(Across L2- L3)

Question 19 ⭐⭐⭐🔥

Ammonium nitrate (NH₄NO₃) dissolves in water, resulting in a temperature drop. 

a) If this scenario occurs in a closed system, explain how the entropy of the surroundings would change and discuss its impact on the overall system’s spontaneity. (3 marks)

b) Compare the entropy change in this process to that of the evaporation of water (at 100°C). (2 marks)

(L3.6: predict the effect of temperature changes on spontaneity)

Question 20 ⭐⭐⭐🔥

A reaction between sodium thiosulfate and hydrochloric acid produces a precipitate of sulfur as shown below:

Na₂S₂O₃(aq) + 2HCl(aq) -> 2NaCl(aq) + H₂O(l) +SO₂(g) +S(s)

a) Given that ΔH is slightly negative and the reaction occurs at room temperature; predict the likely sign of ΔG and justify your reasoning, referring explicitly to (ΔH, T and ΔS). (4 marks)

b) If the reaction were conducted in a sealed container, predict how the pressure of the system would change over time and analyse its effect on the reaction’s spontaneity. Include a discussion of Le Chatelier’s Principle and its relevance to the equilibrium position of the reaction. (4 marks)

(L3.6: predict the effect of temperature changes on spontaneity)

Part A:

Most exothermic reactions involve some kind of heat, and many chemicals used in experiments are toxic. Risk management strategies can include PPE such as safety glasses and gloves, and to avoid burns long hair should be tied back and tongs can be used when handling hot containers.

Part B:
Aim: To determine if dissociation of ionic substances are exothermic or endothermic reactions.

Hypothesis (given): Dissociation of ionic substances are always exothermic.
Ionic substances are often heat resistant due to their strong ionic bonds. The strength of ionic bonds between elements is significantly affected by electronegativity. Smaller ions and more polar ions have incredibly strong bonds, making dissociation (theoretically) endothermic as more energy is required to break the bonds. Whereas, larger ions with weaker ionic bonds are more likely to be exothermic.

Method: By testing the dissociation of a range of ionic substances based on different strengths of ionic bonds. Test each substance at room temperature, in an ice bath, and when heated. This establishes the control and will highlight which type of ionic substances are exothermic or endothermic reactions.

Possible substances that can be tested:

• high strength of ionic bonds – MgO, CaO, Al₂O₃, LiF, NaF, KF
• low strength of ionic bonds – NaCl, KBr, CaCO₂, NH₄Cl

Question 2

(HINT: Start by looking at your formula sheet to find the equations we need to use when finding temperature change)

Assumptions made: The assumptions made throughout these calculations is that the reaction goes to completion and that no heat is lost to the surroundings.

Question 3

(HINT: Use q = mcΔT, to start by finding the energy produced or required for this reaction)

Question 4

(HINT: consider differences in the starting point of enthalpy vs end point of enthalpy? What about differences in activation energy?)

The energy profile diagram to the left displays an endothermic reaction whilst the diagram to the right shows an exothermic reaction. This can be seen first by comparing the activation energy.

Endothermic reactions require significantly more energy to proceed, whilst exothermic reactions are more likely to be spontaneous or require only a little activation energy.

Lastly the overall change in enthalpy (energy) demonstrates that the graph on the left absorbs more energy, meaning there is an overall increase in enthalpy (aligning with an endothermic reaction). Whilst the right diagram aligns with an exothermic reaction which releases energy into its environment.

Question 5

As seen in the diagram, using a catalyst lowers the activation energy, allowing for a more efficient reaction pathway for the system to take. This speeds up the rate of reaction as it is “easier” for the reaction to occur (as the activation energy is lowered).

Question 6

Question 7

(HINT: consider the energy required/released when breaking and forming bonds)

The law of conservation of energy states that energy cannot be created or destroyed. When a reaction occurs, energy is used to break bonds and then released when bonds are formed. Typically the energy required to break a bond is greater than the energy released when bonds are reformed, meaning that the overall change in enthalpy is endothermic.

Question 8

a) This equation relies on the enthalpy of formation. As all the reactants are in their elemental state, the value is zero. Because energy cannot be created or destroyed as stated by the law of conservation of energy, the energy as a result of the formation of calcium carbonate is released into the surroundings.

b) Even though it requires a lot of energy to break the bonds between methane, water and carbon dioxide have significantly stronger bonds. The law of conservation of energy states that energy cannot be created or destroyed which means that the excess energy from the formation of water and carbon dioxide is released into the environment, hence the overall enthalpy change is negative and this is an exothermic reaction.

c) This is the reverse reaction of the Haber Process, and to break the bonds in ammonia requires a lot of energy. This energy is absorbed from the surroundings as energy cannot be created or destroyed. The overall change in enthalpy is positive and therefore this is an endothermic reaction.

Question 9

Question 10

(HINT: Use the given equations to establish the overall equation to find the enthalpy value)

Question 11

a)

b) As the respiration reaction is the reverse reaction of photosynthesis. The enthalpy change for the respiration reaction would be -2804.8 kJ.

Question 12

Dissolution of ammonium chloride

a) As the temperature is decreasing this is an endothermic reaction, meaning that ΔH is positive. However during dissolution, ions dissolve which increases the disorder in the system so ΔS is also positive.

b) For the solution to occur spontaneously the TΔS term must dominate in the Gibbs free energy equation. This can occur either with a high temperature, or a ΔS value which is larger than the ΔH value. Which would result in a negative ΔG value, the requirement for spontaneity.

c) As this reaction occurs without the addition of heat, entropy is the dominant driving force. This can be seen due to the spontaneity of the reaction even though it is an endothermic process.

Question 13

An increase in entropy is classified as a system that has more chaos or movement. By considering the movement alone in the particles of tap water compared to boiling water, boiling water has significantly more movement or chaos. This shows that the entropy has evidently increased making it a good model.

Question 14

Increase in entropy

a) The system shows an increase in entropy as the number of molecules is increasing from one (reactants) to three (products).

b) Entropy is increasing as the state is changing from liquid to gaseous as well as the number of molecules is increasing from nine to thirteen.

c) The equation is decreasing in entropy as the molecules decrease from three to two.

d) As more molecules are being produced (seven reactants compared to twelve products) entropy is increasing.

Question 15

(HINT: Reactions are spontaneous when our ΔG value is less than zero; what values of ΔH, T and ΔS will give us a negative ΔG value?)

1. As the temperature is a small number, the value taken away from the exothermic enthalpy value is most likely not going to exceed it. Therefore the reaction will be spontaneous.
2. This reaction will not be spontaneous as you will end up with a positive ΔG value.
3. This reaction is unlikely to be spontaneous. However, if the ΔH is very small compared to the ΔS then it may be spontaneous.
4. This reaction will be spontaneous as the exothermic value and increase in enthalpy will give a negative ΔG.

Question 16

(HINT: At what value of ΔG is a reaction spontaneous?)

 

Therefore this reaction is spontaneous and likely to move in the forward direction as ΔG is less than zero.

Question 17

Gibbs free energy shows us that for a reaction to be spontaneous, the value of ΔG must be less than zero. This does not mean that it is impossible for reactions with a ΔG greater than zero to occur, they may just need a boost through extra energy to favour them.

Therefore, if forward reaction is not spontaneous, it is true that the reverse reaction will be favoured, but it does not make it impossible for the forward reaction to occur.

Question 18

a)

b) If this reaction occurs in an open system, the temperature increase would be smaller as heat is lost to the surroundings. A smaller ΔT would result in a smaller amount of heat being released which would give a less negative ΔH (smaller in magnitude). Therefore the reaction would appear less exothermic due to the heat lost to the surroundings.

c) As energy cannot be created or destroyed (by the law of conservation of energy), the chemical energy stored in the bonds is transferred into thermal energy. This thermal energy increases the temperature in a closed system, or in an open system some heat would be lost to the surroundings.

Question 19

Dissolution of Ammonium chloride

a) As heat is being absorbed from the surroundings for this reaction to occur (it’s an endothermic process), the entropy of the surroundings would also decrease. Furthermore, as this process is spontaneous, the key driver in this dissolution is the entropy, resulting in a large amount of energy absorbed from the surroundings for this reaction to occur.

b) Whilst the dissolution of ammonium nitrate is resulting in an increase in entropy (as solid particles are becoming ionised in solution), evaporation of water is creating significantly more disorder due to the state change from liquid to gaseous.

Question 20

Sodium thiosulfate and hydrochloric acid

a) As ΔH is slightly negative (exothermic) the system already favours spontaneity. The amount of particles is increasing as well (going from three to five) which demonstrates that ΔS is increasing as well. We are told that the reaction is occurring at room temperature, which contributes to the overall negativity of ΔG. Therefore the reaction is spontaneous with a negative ΔG value.

b) As the reaction produces a gas (SO₂) which is being contained, the pressure will build up over time. However, according to LCP this change will be countered by the system in order to keep the equilibrium balanced. This results in the reaction shifting to the left to counter the production of gas. A shift to the left means the reverse reaction is being favoured, so the spontaneity of the forward reaction decreases (ΔG approaches zero). Therefore, whilst the reaction is spontaneous at its starting conditions, the increase in pressure shifts equilibrium shifts to the left, thereby decreasing spontaneity.