THINK ABOUT IT

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Human blood has a certain pH. When people exercise, they produce a lot of carbon dioxide, which in turn builds up carbonic acid in their bodies. A build-up of carbonic acid can alter the blood’s pH, which can put stress on the body. How does the body maintain a healthy equilibrium?

What happens to a system at equilibrium when conditions change?

To answer this question, you will explore

EXPLORING THE TOPIC

Changing Conditions

When a mixture is at equilibrium, the forward and reverse processes are dynamically balanced, going backward and forward at equal rates. For example, if an indicator is dissolved in water, the indicator, HIn, dissociates in the forward process and the ions, H⁺ and In⁻, attract one another in the reverse process. Although equilibrium is dynamic, once equilibrium is established, the concentrations of HIn, H⁺, and In⁻ in the aqueous solution do not change over time, as long as the conditions stay the same.

However, conditions do not always stay the same for an equilibrium mixture. Something new might be added to the mixture or the mixture might be heated or cooled. For example, HCl might be added to the indicator solution described above. Such changes disturb the equilibrium, which changes in a predictable way by adjusting to new concentrations to counteract the disturbance.

Here are three types of changes that can be made to an equilibrium mixture:

Changing the conditions for an equilibrium mixture will disturb the system. The examples here demonstrate how each of the three types of changes (or disturbances) affects a reversible process at equilibrium.

COBALT CHLORIDE EQUILIBRIA

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The dissociation of tetrachlorocobalt (II) ions, CoCl₄²⁻(aq), to aqueous cobalt (II) ions, Co²⁺(aq) and chloride ions, Cl⁻(aq), is especially useful in demonstrating changes in equilibrium because it involves a striking color change.

Cobalt chloride solutions can be pink, blue, or various shades of purple. This figure shows molecular views of tiny volumes of a pink and blue cobalt solution. Notice that there are more Co²⁺ ions (pink circles) than CoCl₄²⁻ ions (blue circles) in the pink solution, but more CoCl₄^2– ions in the blue solution. Notice also that the blue solution has a larger number of Cl⁻ ions (gray circles).

It is possible to predict how a disturbance, such as the addition of Cl⁻ ions to the pink cobalt solution, will change the equilibrium mixture.

CHANGES IN CONCENTRATION

Imagine that you have a pink sample of a cobalt chloride solution. The pink color indicates that the sample has more Co²⁺ ions (one of the products) than CoCl₄²⁻ ions (starting substance). If you change the conditions by adding more Cl⁻ ions to the mixture, the sample turns blue. So adding more products results in the formation of more starting substance.

If you take this blue sample and add silver ions, AgNO₃(aq), to remove chloride ions, Cl⁻, by precipitating AgCl(s), the mixture changes back to pink. So removing Cl⁻ ions (one of the products) results in the formation of more products (Co²⁺ and Cl⁻ ions). The illustration shows the results of these two changes.

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The color changes indicate that adding or removing one of the products in an equilibrium mixture has the effect of changing the final concentrations of the new equilibrium mixture. Likewise, the equilibrium mixture also changes if the concentration of starting substance is changed. The change depends on which side of the reversible process gets the stress.

CHANGES IN TEMPERATURE

This same reversible process can be used to test the effects of changing temperatures on an equilibrium mixture. If a pink starting sample of cobalt chloride solution is heated, it turns blue. If a blue sample of this mixture is cooled, it turns back to pink or pinkish-purple. In this particular reversible process, heating favors the starting substance.

Heating does not always cause more starting substance to form. The direction of the change depends on whether the process is exothermic or endothermic. For an endothermic process, heating produces more products. For an exothermic process, cooling produces more products. The reversible cobalt process is exothermic in the directions the chemical equation is written, so heating causes the formation of more starting substance.

CHANGES IN PRESSURE

Pressure changes are particularly important when considering equilibrium mixtures that are in the gas phase. Recall from Unit 3: Weather that it is common to use a cylinder and piston as a flexible container for a gas. Pushing down on the piston to decrease the gas volume increases the external pressure on a gaseous equilibrium mixture. A change will occur in the direction that decreases the number of moles of gas and relieves some of the pressure. Recall that the pressure decreases as the number of moles decreases for the same temperature and volume.

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For example, consider the process for the formation of ammonia gas from nitrogen and hydrogen gases.

N₂(g) + 3H₂(g) ⇋ 2NH₃(g)

Decreasing the volume of the container increases the pressure on the gas mixture. The system will change so that more ammonia forms. This is because there are four moles of gas on the starting substance side of the equation and only two moles of gas on the product side.

The creation of more ammonia relieves the stress of the added pressure by reducing the total number of moles of gas in the mixture.

Le Châtelier’s Principle

These very predictable changes to equilibrium mixtures were first noted around 1885 by a French chemist named Henry Le Châtelier. He observed that when the conditions of an equilibrium mixture are modified, the system responds in such a way as to reduce the effect of the change. This idea is known as Le Châtelier’s Principle.

Le Châtelier’s principle allows you to predict qualitatively how the system responds to changes in concentration, pressure, or temperature.

AN EXPLANATION OF LE CHÂTELIER’S PRINCIPLE

If more of one of the products is added to the equilibrium mixture of starting substances and products, this momentarily increases the rate of the reverse process. The concentrations in the mixture adjust until the rates of the forward and reverse processes are equal. This adjustment results in a new equilibrium mixture that has more starting substance compared with that of the initial solution. The value of the equilibrium constant K is the same for both the adjusted mixture and the initial solution.

A second change is to raise or lower the solution’s temperature, which also changes the composition of the equilibrium mixture of starting substances and products. In this case, the changes in composition result in a new value of K depending on if the forward process is exothermic or endothermic. Because there is a new value of K, the concentrations of starting substances and products change to satisfy the new equilibrium constant equation. If the forward process is exothermic (heat is released), the value of K gets smaller as the temperature is raised. Likewise, if the forward process is endothermic (heat is required), the value of K gets larger as the temperature is raised.

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BLOOD pH AND EXERCISE

The human body is a large collection of reversible processes. For processes to proceed at the appropriate rates, and for substances to stay in the appropriate concentrations within our bodies, a balance must be maintained. For example, the pH of the body’s blood is normally around 7.4, slightly basic. When you exercise, more carbon dioxide is produced, and it dissolves in the bloodstream. This is a natural waste product that you normally exhale from the lungs. However, carbon dioxide in the bloodstream also leads to the formation of carbonic acid, H₂CO₃. A large build up of carbonic acid would increase the acidity of the blood and lower the pH.

One of the ways your body deals with a change in the blood’s pH is to exhale more rapidly, removing CO₂ from the bloodstream and shifting the process back toward the right. This is one reason why we breathe more heavily when exercising—to rid the body of excess carbon dioxide and increase the blood’s pH.

LESSON SUMMARY

What happens to a system at equilibrium when conditions change?

When conditions are changed for a system at equilibrium, the system will respond by reducing the effect of the change. This is known as Le Châtelier’s principle. Changing conditions (or stresses) include changes to temperature, concentration, or pressure. A system regains equilibrium by rebalancing the concentrations of starting substances and products. The exact nature of the effects on a system are predictable because they depend on the type of change in condition that the system encounters. A temperature change affects the value of K. A pressure change or a change in concentration stresses the system, but does not change the value of K.