Temperature Effect on Chemical Equilibrium (Nancy Nasr)

This demonstration effectively shows how the input or removal of heat to/from a system can shift the chemical equilibrium according to Le Chatelier’s principle. When household ammonia (a weak base) is diluted with water, the reaction produces hydroxide ions as evidenced by the following chemical reaction:

(6.1) NH_{3 (aq)} + H₂O _(l) à NH₄⁺ _(aq) + OH^- _(aq)

ß

According to Le Chatelier’s principle, depending on whether a reaction is endothermic or exothermic, the addition or removal of heat to a chemical reaction will shift the equilibrium in a direction opposite to that stress. In the above demonstration example, if heat is added to the chemical reaction (i.e. when we put the diluted ammonia solution into a hot water bath) the equilibrium shifts to the left to increase the concentration of ammonia. This is evidenced by the phenolphthalein indicator turning the solution a very light shade of pink. Contrary to this, when heat is removed from the chemical reaction (i.e. when the diluted ammonia solution is put into an ice water bath), the equilibrium shifts to the right (i.e. towards the side that produces more heat), thereby increasing the concentration of hydroxide ions. This is evidenced by the phenolphthalein indicator turning the solution more pink.

Questions and Answers:

Q: Why is phenolphthalein the indicator of choice in this demonstration?

A: Different indicators have different pH intervals in which they change color. For example, Methyl Orange changes color over the pH range from about 3 – 4. If a solution is about pH 3 the Methyl Orange indicator will be red, and when the solution is about pH 4 Methyl Orange Indicator will be yellow. Because the demonstration that we conducted is dealing with bases that have pH values much higher than 4, the Methyl Orange Indicator would not show us any visible color change. Phenolphthalein is used in our demonstration because this indicator changes color over the pH interval of roughly 8 – 10, which is the pH range expected when dealing with dilute ammonia (which is a weak base) and hydroxide ions (which are more strongly basic).

Q: Based on the demonstration, predict whether the reaction in (6.1) is endothermic or exothermic. Explain.

A: The reaction in (6.1) is exothermic. In an exothermic reaction, we consider heat as a product, while in endothermic reactions we consider heat as a reactant. The difference between endothermic and exothermic reactions can be illustrated as follows:

Endothermic: Reactants + heat à Products

Exothermic: Reactants à Products + heat

If we treat heat as if it were another reactant or product, we can see that, according to Le Chatelier’s principle, the addition of heat to endothermic reactions will cause an increase in the concentration of products, while in exothermic reactions the addition of heat will cause an increase in the concentration of reactants. Looking back at our demonstration, and equation (6.1), indeed the removal of heat from this reaction increased the concentration of hydroxide ions since the reaction shifts to the side that produces more heat (as evidenced by the deep pink shade of phenolphthalein observed), thus it can be concluded that it is an exothermic reaction.

Q: The base-dissociation constant for ammonia dissolved in water is 1.8 X 10 ^-5 at 25ºC. What do you expect will happen to the value of the base-dissociation constant when a diluted solution of ammonia is heated?

A: The base-dissociation constant for ammonia dissolved in water will become smaller at higher temperatures. Referring back to the reaction in (6.1), the equilibrium constant expression for the reaction is as follows:

K_b = [NH₄⁺] [OH^-]

------------------

[NH₃]

(**Note: water is omitted from the above expression because it is a solvent)

When we heated the diluted solution of ammonia, the equilibrium shifted to the left (as evidenced by the light pink color of the solution), thereby increasing the concentration of NH₃. If we assume a higher concentration of NH₃, then looking at the above expression we can see that the numerator will become smaller relative to the denominator. At 25ºC we do not know the exact values of each of the concentrations, but we know that if one of those values increases then there will be a relative change to the dissociation constant value. Again, since the value of [NH₃] increases, we can predict that the dissociation constant value will decrease from its value at 25ºC.

Applications to everyday life:

1) The Haber process provides an interesting example of the above concept in an exothermic reaction. The Haber process essentially forms ammonia from hydrogen and nitrogen gases. This reaction can be summarized as follows:

N_{2 (g)} + 3 H_{2 (g)} à 2 NH_{3 (g)} + 92 KJ

ß

The Haber process is important to us because the ammonia produced in this reaction is then used to generate fertilizer. The fertilizer ultimately generated from this process is estimated to be responsible for sustaining 1/3 of the world’s population. In order to maximize the amount of ammonia generated from this reaction, scientists utilize the concepts of Le Chatelier’s principle. Looking at the above reaction, if scientists were to increase the temperature of the reaction, the equilibrium would shift to the left producing more reactants than products. As a result, to use temperature as a means to shift the equilibrium to the right, scientists must actually decrease the temperature of the reaction. **Note: A decrease in temperature alone will not be sufficient at producing economical amounts of ammonia. As a result, the Haber process relies on increased pressure (also in accordance with Le Chatelier’s principle), in addition to low temperature to maximize ammonia yield.

2) Nitric Oxide emissions show the practical importance of temperature and how it affects reaction equilibria. Nitric oxide is formed by way of the following endothermic reaction:

½ N_{2 (g)} + ½ O_{2 (g)} à NO _(g)

ß

Nitrogen and Oxygen gases combine in the atmosphere to form Nitric Oxide. Because this reaction is endothermic, the formation of Nitric Oxide increases dramatically at higher temperatures. In automobile engines, the combustion of fuels is exothermic, thus increasing the temperature of the surroundings. When Nitrogen and Oxygen gases are trapped in the engines, the formation of Nitric Oxide occurs rapidly (i.e. the equilibrium shifts to the right) due to the increase in temperature caused by the combustion of fuels. Though the temperatures of the gases decreases as the formed products make their way out of the exhaust pipe, the temperatures are not low enough to again shift the equilibrium to the left, and convert existing Nitric Oxide back to Nitrogen and Oxygen gases. In effect, the Nitric Oxide remains, and is then released into the atmosphere via the exhaust of the vehicle. Nitric Oxide is an undesirable byproduct of this combustion because it aids in the formation of acid rain and contributes to global warming.

3) Thermoregulation, a form of homeostasis, is used by animals and humans as a way of regulating internal body temperature. An internal body temperature that is too warm or too cold can have adverse effects on the way in which the animal’s biochemical processes function. In humans for example, the rate of many enzyme-dependent reactions increase 2-3 times for every 10ºC increase in body temperature, until temperatures become so high that the enzymes begin to denature. If enzymes denature, many vital biochemical reactions cannot proceed in a timely manner which can result in the depressed production of many important chemical byproducts, necessary for survival. Similarly, if internal temperatures become too cold, the lipid bi-layer present in cell membranes may actually become less fluid. This is adverse because decreased fluidity in membranes can decrease the membrane’s permeability. Cells in the body depend on semi-permeable membranes to get nutrients successfully into cells, and get waste products out. Wastes can accumulate in cells, triggering cell death. In effect, the human body maintains equilibrium at temperatures around 37ºC. Shifting the equilibrium towards either direction becomes adverse for the organism, and is carefully regulated by complex pathways and mechanisms in the brain.

Photographs: