Table of Contents
Nature of Chemical Equilibrium
Definition of chemical equilibrium
Chemical equilibrium is the state of a chemical system at which a constant concentration of products and reagents is present. Reactions, which take place in homogeneous solutions, seem to have come to rest because no changes in concentrations of the participating substances can be determined. Substance turnover occurs only on the particle level, which is why chemical equilibrium is also referred to as dynamic equilibrium.
For each reaction, the position of equilibrium, under certain surrounding conditions, is determined by a natural constant.
This form of reaction is also known as a reversible reaction, as it occurs in both directions and simultaneously. This condition results in the reaction equation for the reaction type of the equilibrium reaction to contain a double arrow. However, reversible reactions can take place only if none of the reaction partners leave the system.
Examples of equilibrium reactions
- Pure water: H2O dissociates into H+ and OH–. In pure water, there is an equilibrium between H2O and the dissociated ions. The position is very far on the side of H2, and this results in a pH value of 7.
- If glucose is dissolved in water at room temperature, a stable concentration relation results in 63 % β-glucose and 37 % α-glucose.
Requirements of chemical equilibrium
- Closed or concluded system: Reversible reactions can occur only if none of the participating substances can escape.
- Reversible reaction: If a reaction has begun and the first products have formed, an instant and immediate backward reaction take place so that the products are decomposed to their original substances again. The reaction velocities of the reaction partners adjust due to the back and forth swinging of the reactions until a constant relation has developed after a certain period of time.
Features of chemical equilibrium
- Forward and backward reaction happens simultaneously: dynamic equilibrium.
- Identical reaction velocities [vforth = vback]
- Adjustable from both sides
- Original substances and reaction products are present simultaneously and in a constant concentration relation.
- Incomplete substance turnover
- Substance conversion is observable only on the particle level due to constant substance concentration.
- This applies: CRP / COS = constant.
- Catalyzers do not influence the location of the equilibrium.
- Catalyzers shorten the time until equilibrium is reached.
Development of chemical equilibrium
In order to explain the equilibration of a reaction, be conscious of the meaning of the term reaction velocity. Many reactions can go backward, as well as forward. Reaction velocity is the change in substance concentration during a certain period of time. If forward and backward reactions occur simultaneously, which is typical for an equilibrium reaction, the following applies:
A + B ⇔ C + D
Vforward = k1 • cA • cB
Vbackward = k2 • cC • cD
[k= proportionality factor, c = concentration]
Thus, the following applies in chemical equilibrium:
Vforward = vbackward, and k1 • cA • cB = k2 • cC • cD
If a reaction takes place incompletely in a closed system and is also reversible (equilibrium reaction), the reaction initially has a high reaction velocity, as the concentrations of the original substances are high. The reaction velocity of the forward reaction gradually decreases because the substance concentrations of the reagents constantly decrease, while the backward reaction gains velocity because the substance concentrations of the products increase in the course of the reaction.
This process swings back and forth until a state is reached at which the same amount of products and reagents is formed. In this state, the velocities of the forward and backward reactions are equal. This is why the reaction seems to have halted. Macroscopically, no changes can be observed as the chemical transitions occur only on the particle level.
The position of the chemical equilibrium is specific for each reaction and corresponds to a natural constant, which means that it cannot be changed. However, the time for equilibration can be shortened with the help of catalyzers.
Equilibration time is also specific for each reaction, but only at constant temperature conditions. Shortening of the time can be explained by the ability of catalyzers to cause more ‘effective collisions’ in their active state so that the chemical reaction is accelerated.
Law of Mass Action
The law of mass action, or LMA, offers the mathematical instrument to quantitatively describe the position of chemical equilibrium.
If more than 50% of the original substances react to products, the equilibrium is rather on the right side of the overall picture, referred to as equilibrium located on the right side.
Requirements of the law of mass action
- Closed system
- Equilibrated state
Calculation of the law of mass action
k = equilibrium or mass action constant
c = substance concentration
A, B, C, D = reaction partners or their substance concentrations
a, b, c, d = stoichiometric numbers; can be taken from the reaction equation
Explanation of the law of mass action
One main statement of the LMA is that the relation k of the multiplied product concentrations to the multiplied reagent concentrations is constant for certain reactions under determined conditions. Thus, the quotient K of this equation is also referred to as equilibrium or mass action constant.
However, one must consider that the LMA can be applied only to diluted solutions. In more concentrated solutions, there are deviations between the particles due to interactions. For example, OH ions cannot move in strong bases, as there is not enough solvent present for the ions. This makes one conclude that the position of the equilibrium fluctuates, depending on the concentration. Yet, the same concentrations develop in the equilibrium according to the LMA, which is why one always considers diluted solutions as ‘reference’.
Furthermore, location of equilibrium depends on temperature and, possibly, on pressure relations. The section below discussing disturbances of chemical equilibrium will provide further explanation.
The calculated equilibrium constant Kc has great importance for further calculations. On the bases of the value of Kc, one can calculate the transformed amount of reagents or the exploit of reaction products. Kc can be determined from experimental values.
Disturbances or Influencing the Equilibrium
If a chemical equilibrium is disturbed, an acceleration of the reaction occurs, which then eliminates or reverses the disturbance. This rule is also known as the principle of least constraint, or Le Châtelier’s principle. The “constraint” refers to the disturbance of the equilibrium, which leads to the reaction having to be compensated by acceleration.
If one applies a constraint to a system in equilibrium, the system shifts in the direction of a new equilibrium level so that the effect of the constraint becomes minimal. That is the smallest.
A disturbance can be triggered by different factors. As mentioned previously, the position of the equilibrium can be changed by deviation of temperature and pressure conditions. In addition, participating substance amounts also have an influence.
The following section examined other factors to illustrate the way changes can be triggered by different factors.
Energy input (e.g., via heating) results in a reinforced “uphill” reaction. An increase in the formation of reagents occurs, which actually forms products, which store their energy there, under ‘normal’ conditions.
This means that an increase in temperature promotes an endothermic reaction and that the value of the equilibrium constant decreases. Analogously, the opposite happens in the event of a decrease in energy or temperature: The location of the equilibrium shifts in the direction of the products and the exothermic reaction is promoted.
Changes in the amount of substance
For the sake of “rescue,” the following reactions occur with the adding or removal of reaction partners: [reaction: A + B → C + D].
- Adding of original substances A or B → increased formation of the reaction products C and D
- Adding of the products C or D → increased formation of the reagents A and B
- Removal of A or B → increased formation of A or B
- Removal of C or D → increased product formation (C, D)
An increase in the concentration of a substance promotes its consumption and a decrease in concentration promotes its reproduction.
Via adding of acids, bases, or precipitants, the concentration of reaction partners can, however, also be disturbed. In such a case, two coupled equilibrium reactions often occur parallel.
Changes in pressure conditions
If the participating substances in an equilibrium reaction in a closed system are gasses, a change in pressure results in a change in the location of the chemical equilibrium. If the reaction partners have another aggregate phase than gaseous, the equilibrium is not affected or changed. The background of this phenomenon is that changes in volume at reactions with non-gaseous substances are so small that the dependency of the location of the equilibrium on the pressure can be neglected.
If an increase in pressure occurs during a reaction that takes place under a decrease in volume, the chemical equilibrium shifts to the side of the products. An increase in pressure during a reaction that takes place under an increase in volume leads the location of the equilibrium to be shifted to the reagents.
A decrease in pressure promotes the reaction, which occurs under an increase in volume.