3.2 Chemical equilibrium  (Page 2/2)

A closed system is one in which energy can enter or leave, but matter cannot. The second beaker covered by the bell jar is an example of a closed system. The beaker can still be heated or cooled, but water vapour cannot leave the system because the bell jar is a barrier. Condensation changes the vapour to liquid and returns it to the beaker. In other words, there is no loss of matter from the system.

Open and closed systems

An open system is one whose borders allow the movement of energy and matter into and out of the system. A closed system is one in which only energy can be exchanged, but not matter.

Reversible reactions

Some reactions can take place in two directions. In one direction the reactants combine to form the products. This is called the forward reaction . In the other, the products react to form reactants again. This is called the reverse reaction . A special double-headed arrow is used to show this type of reversible reaction :

$XY+Z\rightleftharpoons X+YZ$

So, in the following reversible reaction:

$H_2\left(g\right)+I_2\left(g\right)\rightleftharpoons 2HI\left(g\right)$

The forward reaction is $H_2\left(g\right)+I_2\left(g\right)\to 2HI\left(g\right)$ . The reverse reaction is $2HI\left(g\right)\to H_2\left(g\right)+I_2\left(g\right)$ .

A reversible reaction

A reversible reaction is a chemical reaction that can proceed in both the forward and reverse directions. In other words, the reactant and product of one reaction may reverse roles.

Demonstration : the reversibility of chemical reactions

Apparatus and materials:

Lime water (Ca(OH) ${}_2$ ); calcium carbonate (CaCO ${}_3$ ); hydrochloric acid; 2 test tubes with rubber stoppers; delivery tube; retort stand and clamp; bunsen burner.

Method and observations:

1. Half-fill a test tube with clear lime water (Ca(OH) ${}_2$ ).
2. In another test tube, place a few pieces of calcium carbonate (CaCO ${}_3$ ) and cover the pieces with dilute hydrochloric acid. Seal the test tube with a rubber stopper and delivery tube.
3. Place the other end of the delivery tube into the test tube containing the lime water so that the carbon dioxide that is produced from the reaction between calcium carbonate and hydrochloric acid passes through the lime water. Observe what happens to the appearance of the lime water. The equation for the reaction that takes place is: $Ca{\left(OH\right)}_2+C{O}_2\to CaC{O}_3+H_{2}O$ CaCO ${}_3$ is insoluble and it turns the limewater milky.
4. Allow the reaction to proceed for a while so that carbon dioxide continues to pass through the limewater. What do you notice? The equation for the reaction that takes place is: $CaC{O}_3\left(s\right)+H_{2}O+C{O}_2\to Ca{\left(HC{O}_3\right)}_2$ In this reaction, calcium carbonate becomes one of the reactants to produce hydrogen carbonate (Ca(HCO ${}_3$ ) ${}_2$ ) and so the solution becomes clear again.
5. Heat the solution in the test tube over a bunsen burner. What do you observe? You should see bubbles of carbon dioxide appear and the limewater turns milky again. The reaction that has taken place is: $Ca{\left(HC{O}_3\right)}_2\to CaC{O}_3\left(s\right)+H_{2}O+C{O}_2$

Discussion:

• If you look at the last two equations you will see that the one is the reverse of the other. In other words, this is a reversible reaction and can be written as follows: $CaC{O}_3\left(s\right)+H_{2}O+C{O}_2\rightleftharpoons Ca{\left(HC{O}_3\right)}_2$
• Is the forward reaction endothermic or exothermic? Is the reverse reaction endothermic or exothermic? You should have noticed that the reverse reaction only took place when the solution was heated. Sometimes, changing the temperature of a reaction can change its direction.

Chemical equilibrium

Using the same reversible reaction that we used in an earlier example:

$H_2\left(g\right)+I_2\left(g\right)\rightleftharpoons 2HI\left(g\right)$

The forward reaction is:

The reverse reaction is:

When the rate of the forward reaction and the reverse reaction are equal, the system is said to be in equilbrium . [link] shows this. Initially (time = 0), the rate of the forward reaction is high and the rate of the reverse reaction is low. As the reaction proceeds, the rate of the forward reaction decreases and the rate of the reverse reaction increases, until both occur at the same rate. This is called equilibrium.

Although it is not always possible to observe any macroscopic changes, this does not mean that the reaction has stopped. The forward and reverse reactions continue to take place and so microscopic changes still occur in the system. This state is called dynamic equilibrium . In the liquid-vapour phase equilibrium demonstration, dynamic equilibrium was reached when there was no observable change in the level of the water in the second beaker even though evaporation and condensation continued to take place.

There are, however, a number of factors that can change the chemical equilibrium of a reaction. Changing the concentration , the temperature or the pressure of a reaction can affect equilibrium. These factors will be discussed in more detail later in this chapter.

Chemical equilibrium

Chemical equilibrium is the state of a chemical reaction, where the concentrations of the reactants and products have no net change over time. Usually, this state results when the forward chemical reactions proceed at the same rate as their reverse reactions.