HSC Chemistry Module 5: Equilibrium and Acid Reactions Practice Questions

Factors that Affect Equilibrium 

Question 6 ⭐⭐

Account for the effect of increased temperature and increased volume on the reaction between nitrogen dioxide and dinitrogen tetraoxide using the Le Chatelier’s principle. (4 marks)

(L2.1: Investigate the effects of temperature, concentration, volume and/or pressure on a system at equilibrium and explain how Le Chatelier’s principle can be used to predict such effects)

Question 7 ⭐⭐

Using the Le Chatelier’s principle, explain how the addition of NaCl can be used to produce more Fe³⁺ and SCN^– ions in the following reaction. (4 marks)

(L2.1: Investigate the effects of temperature, concentration, volume and/or pressure on a system at equilibrium and explain how Le Chatelier’s principle can be used to predict such effects)

Question 8 ⭐⭐⭐

Using the Le Chatelier’s principle, deduce how we can increase the yield of Co(H₂O)₆²⁺ by temperature, concentration and pressure. The chemical reaction is as follows: 

(L2.1: Investigate the effects of temperature, concentration, volume and/or pressure on a system at equilibrium and explain how Le Chatelier’s principle can be used to predict such effects)

Question 9 ⭐⭐⭐🔥

Consider the reaction in question 8. Explain whether the addition of NaCl increases the production of its products using the collision theory. (4 marks)

(L2.2: Explain the overall observations about equilibrium in terms of the collision theory) 

Question 10 ⭐⭐

Predict the most likely position of the equilibrium in the following conditions. (4 marks)

a) Low activation energy in forward reaction but high activation energy in reverse reaction 

b) High activation energy in forward reaction but low activation energy in reverse reaction 

c) Negative heat of reaction within high temperatures 

d) Negative heat of reaction within low temperatures 

(L2.3: Examine how the activation energy and heat of reaction affect the position of equilibrium) 

Question 11 ⭐

Write the equilibrium expression for each of the reactions below (3 marks): 

(L3.1: Deduce the equilibrium expression (in terms of Keq) for homogeneous reactions occurring in solution) 

Question 12 ⭐⭐⭐

A mixture of 5.0 mol H_{2 (g)} and 10.0 mol I_{2 (g)} are placed in a 5.0 L container at 450°C and allowed to come to equilibrium. At equilibrium the concentration of HI_(g) is 1.87 mol/L. Calculate the value for the equilibrium constant, K, for this reaction under these conditions. (4 marks) 

(L3.2: Perform calculations to find the value of Keq and concentrations of substances within an equilibrium system and use these values to make predictions on the direction in which a reaction may proceed) 

Question 13 ⭐⭐⭐

Explain how an increase in temperature can affect the K_eq of an endothermic and exothermic reaction. (4 marks)

(L3.3: Qualitatively analyse the effect of temperature on the value of Keq)

Question 14 ⭐

During the practical that investigates the K_eq of iron (III) thiocyanate equilibrium, students obtained a different K_eq value to that of the theoretical value. Suggest a source of error for this. (3 marks)

(L3.4: Conduct an investigation to determine Keq of a chemical equilibrium system, for example: Keq of the iron (III) thiocyanate equilibrium) 

Question 15 ⭐⭐⭐

Deduce what would be observed in the following conditions. (6 marks)

(L3.5: Explore the use of Keq for different types of chemical reactions, including but not limited to: Dissociation of ionic solutions, Dissociation of acids and bases)

Solution Equilibria 

Question 16 ⭐

Explain how the dissolution of cations differ to the dissolution of anions in water. Use diagrams if appropriate. (4 marks) 

(L4.1: Describe and analyse the processes involved in the dissolution of ionic compounds in water)

Question 17 ⭐⭐

How do Aboriginal and Torres Strait Islander People utilise the solubility equilibria to remove toxins in cycad fruits? (4 marks) 

(L4.2: Investigate the use of solubility equilibria by Aboriginal and Torres Strait Islander Peoples when removing toxicity from foods, for example: Toxins in cycad fruit)

Question 18 ⭐⭐⭐

Complete the following table. (9 marks)

(L4.3: Conduct an investigation to determine solubility rules, and predict and analyse the composition of substances when two ionic solutions are mixed, for example: Potassium chloride and silver nitrate, Potassium iodide and lead nitrate, Sodium sulfate and barium nitrate)

Question 19 ⭐⭐⭐🔥

Calculate the the following (6 marks): 

a) What is the solubility of Sr²⁺when the K_sp of strontium sulfate is 3.44×10^-7? 

b) What is the solubility of CO₃^2-when the K_sp of calcium carbonate is 3.36×10^-9? 

c) What is the solubility of CrO₄^2- when the K_sp of silver chromate is 1.12×10^-12

(L4.4: Derive equilibrium expressions for saturated solutions in terms of Ksp and calculate the solubility of an ionic substance from its Ksp value) 

Question 20 ⭐⭐⭐🔥

25.0 mL of 0.0020 M potassium chromate are mixed with 75.0 mL of 0.000125 M lead(II) nitrate. Will a precipitate of lead(II) chromate form. K_sp of lead(II) chromate is 1.8 x 10^-14 (3 marks)

(L4.5: Predict the formation of a precipitate when given the standard reference values for Ksp) 

Question 2

a) The burning of magnesium is an irreversible combustion reaction. As it reaches a point of completion and will not change from the formed product MgO the system has reached static equilibrium.

b) This reaction is in dynamic equilibrium as it is reversible and any change can be countered to re-establish equilibrium.

c) The Haber process is a great example of dynamic equilibrium as it is reversible and follows Le Chatelier’s principle when changes are made to the system.

Question 3

Dynamic equilibrium can proceed in open systems, however this is only possible when the rates of input and output are balanced. An open system must be able to self-regulate in order to maintain dynamic equilibrium. For example, body temperature regulation is an example of dynamic equilibrium in an open system, our bodies sweat or shiver to lose or generate heat (respectively), maintaining the desired temperature.

Question 4

In a combustion reaction, energy is released resulting in a negative enthalpy (exothermic reaction). Entropy is positive as the disorder increases due to the production of a larger amount of molecules.
On the other hand, photosynthesis has positive enthalpy as heat is absorbed in order for this reaction to proceed. Entropy is negative as the products are more organised (and there’s less molecules).
These factors make combustion a spontaneous reaction and photosynthesis a non-spontaneous reaction.

Question 5

Collision theory states that particles must collide with the required activation energy and correct orientation for a reaction to occur. Systems with more energy will have more successful collisions as the required activation energy is met. Furthermore, having a high concentration of particles will allow for more possible collisions as the desired orientation is more likely to be achieved and there are overall more particles for collision.

Question 6

The forward reaction of nitrogen dioxide and dinitrogen tetroxide is exothermic, meaning it releases heat. If the temperature of this system is increased, Le Chatelier’s principle (LCP) states that this stress must be countered. Adding heat will favour the endothermic (reverse) reaction, resulting in the production of NO₂. Meanwhile, increasing the volume results in more space for gaseous particles to occupy. According to LCP, equilibrium will counteract this change by shifting to the side with more gas moles, therefore favouring the reverse reaction as two molecules of NO₂ are produced for every one molecule of N₂​O₄.

Question 7 

The addition of NaCl to the reaction between Fe³⁺ and SCN^– results in an equilibrium shift to the left. Le Chatelier’s principle states that every stress or change placed on a system will be countered in order to re-establish equilibrium. By adding NaCl, there is another substance that can react with the Fe³⁺ ions, creating FeCl^4- complexes. This means there are less Fe³⁺ ions to react with the SCN^–, shifting the equilibrium to the left so that more Fe³⁺ ions become available.

Question 8

In order to increase the yield of Co(H₂O)₆²⁺ the conditions must favour the reverse reaction. The reaction displayed shows that heat must be added for the forward reaction to occur, meaning the forward reaction is endothermic. To favour the reverse reaction, the system must be cooled, forcing equilibrium to shift to the left by LCP. By increasing the concentration of the products the reverse reaction can be achieved as well (for example by adding water, i.e. diluting the solution). Lastly, pressure can be used to favour the reverse reaction. By increasing the pressure of the system, it will naturally favour the side with fewer particles, again favouring an equilibrium shift to the left. Hence, by LCP the reverse reaction can be favoured by decreasing the temperature, diluting the system with water and increasing the pressure.

Question 9

Adding NaCl to the system above will provide further Cl^– for the reaction. A greater abundance of Cl^– ions increases the availability of particles for collision as stated by collision theory. As the Cl^– ions are primarily on the reactant side, there will be more successful collisions with the reactants, resulting in the forward reaction proceeding. According to collision theory, these greater successful collisions between reactants will increase the products.

Question 10

a) Equilibrium will most likely favour the forward reaction, being positioned to the right.

b) Equilibrium will most likely favour the reverse reaction, positioned to the left.

c) Adding heat to an exothermic reaction will favour the reverse reaction, equilibrium will be positioned to the left.

d) Adding heat to an endothermic reaction favours the forward reaction, equilibrium will be positioned to the right.

Question 11

Question 12

Question 13

The K_eq of any reaction is essentially a ratio of concentrations of products to reactants. An endothermic reaction will continue in the forward direction when the temperature of its system is increased. This means that the concentration of the products becomes stronger and the concentration of the reactants weaker, resulting in a larger K_eq. However exothermic reactions will favour their reverse reaction when heat is added. This means that the concentration of the reactants will increase and the products decrease. The K_eq will become smaller in this case.

Question 14

In this practical investigation, a possible source of error is failing to maintain a constant temperature. For equilibrium to be established, the reaction must be allowed to occur in constant conditions. If the conditions are constantly changing the system will be adjusting to match the new conditions and equilibrium will not be reached.

Question 15

Question 16

Water molecules have positive (hydrogen) and negative (oxygen) dipoles and they are attracted to the anions and cations respectively. For example, when NaCl is dissolved in water, the Cl^– anions will be attracted to the positive dipole of the water molecule, resulting in a ‘shell’ around the anion. Meanwhile, the Na⁺ cation is attracted to the negative dipole of the water molecule.

Question 17

Solubility equilibria was used by the Aboriginal and Torres Strait Islander Peoples to make cycads edible. This was achieved through a process called leaching, where the cycad fruits were placed in a flowing creek, allowing the toxins to dissolve and be washed away with the river. Prior to leaching the cycad fruits, Aboriginal and Torres Strait Islander Peoples would increase the surface area through pounding, slicing or grinding the fruit, making the leaching process more effective.

Question 18

Question 19

Question 20