Posted: June 6, 2014 in Uncategorized

Exercise 1: Equilibrium

Problem 1:

The following reaction has reached equilibrium in a closed container:

C(s)+H2O(g)  CO(g)+H2(g) ΔH> 0

The pressure of the system is then decreased. How will the concentration of the H2 (g) and the value of Kc be affected when the new equilibrium is established? Assume that the temperature of the system remains unchanged.


Table 2

Hydrogen concentration







stays the same


stays the same

stays the same



stays the same

Problem 2:

During a classroom experiment copper metal reacts with concentrated nitric acid to produce NO2 gas.


The NO2 is collected in a gas syringe.


When enough gas has collected in the syringe, the delivery tube is clamped so that no gas can escape. The brown NO2 gas collected reaches an equilibrium with colourless N2O4 gas as represented by the following equation:

2NO2(g)  N2O4(g)Δ< 0

Once this equilibrium has been established, there are 0,01 moles of NO2 gas and 0,03 moles of N2O4 gas present in the syringe.

  1. A learner, noticing that the colour of the gas mixture in the syringe is no longer changing, comments that all chemical reactions in the syringe must have stopped. Is this assumption correct? Explain.

  2. The gas in the syringe is cooled. The volume of the gas is kept constant during the cooling process. Will the gas be lighter or darker at the lower temperature? Explain your answer.

  3. The volume of the syringe is now reduced (at constant temperature) to 75 cm3 by pushing the plunger in and holding it in the new position. There are 0,032 moles of N2O4 gas present once the equilibrium has been re-established at the reduced volume (75 cm3). Calculate the value of the equilibrium constant for this equilibrium.

Problem 3:

Gases X and Y are pumped into a 2 dm3 container. When the container is sealed, 4 moles of gas X and 4 moles of gas Y are present. The following equilibrium is established:

2X(g)+3Y(g)  X2Y3(g)

The graph below shows the number of moles of gas X and gas X2Y3 that are present from the time the container is sealed.

  1. How many moles of gas X2Y3 are formed by the time the reaction reaches equilibrium at 30 seconds?

  2. Calculate the value of the equilibrium constant at t = 50 s.

  3. At 70 s the temperature is increased. Is the forward reaction endothermic or exothermic? Explain in terms of Le Chatelier’s Principle.

  4. How will this increase in temperature affect the value of the equilibrium constant?

Consider the following hypothetical reaction that takes place in a closed 2 dm3 flask at 298 K.

A2(g)+2B2(g)  2AB2(g)

The graph beneath represents the change in the number of moles of each gas in the flask over a period of 20 minutes.

  1. State how long (in minutes) it took for the reaction to reach equilibrium for the first time.

  2. Write down an expression for the equilibrium constant, Kc, for this particular reaction.

  3. Calculate the concentration of each of the reactants and the product using figures from the graph between 5 minutes and 10 minutes and hence calculate the equilibrium constant Kc for this reaction at 298 K

  4. State what a low value of Kc indicates about the yield of product for a reaction.

  5. Why is it not possible to calculate Kc using figures from the graph during the first 5 minutes

  6. State Le Chatelier’s principle.

  7. At 10 minutes the temperature of the flask was increased. Using Le Chatelier’s principle, determine if the production of AB2 is exothermic or endothermic?

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