Creatine Case Study

Case Study

History

Learning Goals /
Concept Map

Creatine and Related Compounds

Structure

Amino Acids

Creatine in the Body

Equilibrium

Creatine-Creatinine Equilibrium

Creatinine Test for Kidney Function

Detection

Regulation and Ethics

Amine & Nitrile Chemistry

Laboratory Synthesis

Chemical Analysis

Creatine-Phosphocreatine Equilibrium

Uses & Side Effects



Equilibrium

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Definition
Equilibrium Constant, K
Pure Solids and Liquids
Le Châtelier’s Principle
Changes in Concentration
Changes in Temperature
Changes in Pressure and Volume
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Definition
Equilibrium is established when there is no net change in the concentrations of the products and reactants during a reversible reaction [3]. Most biological reactions are considered reversible reactions because they occur in closed systems. A closed system is one where the products are not allowed to escape; therefore, both forward and reverse reaction occur [2]. Initially, a reaction may predominantly occur in the forward direction, but as the amount of products increases, gradually the reverse reaction occurs until an equilibrium is established [4]. Once equilibrium is established, the reaction is not static, but rather both these forward and reverse reactions are occurring at the same rate [1].



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Equilibrium Constant, K
When the rates of the forward and reverse reactions are equal and equilibrium is established, one can determine the equilibrium constant, K, for the reaction at a given temperature. This is determined by first balancing the overall reaction equation. Next, the concentration of each product and reactant is raised to its stoichiometric coefficient. Then the concentrations of products are divided by the concentrations of the reactants [3] when the products and reactants are in either the gaseous (g) or aqueous (aq) phase. For example, consider the following one-step reaction:

K is a constant, so it does not change with time. The reason for this is that at equilibrium, the overall concentration of the substances does not change [4]. Hence, if the value of K is close to 1 at equilibrium, then the reaction mixture contains approximately equal amounts of the reactants and products. If the value of K is very large compared to 1, then the reaction mixture contains very little of the reactants compared to the products. And if the value of K is very small compared to 1, then the reaction mixture contains very little of the products compared to the reactants [4].

Rollover K Values for definitions.

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Pure Solids and Liquids
Some reactions include pure solids and liquids. These are known as heterogeneous reactions. The equilibrium constant expression for these reactions does not include the concentrations of the pure solids and liquids because their molar concentrations are constant. The molar concentration of a substance is determined by dividing the number of moles of the substance by the volume it occupies. This is proportional to the density of a substance, which is the mass divided by the volume. The density of a substance is an intensive property (a property independent of the size of the substance), so the molar concentration must also be intensive. Given that the molar concen
tration of pure solids and liquids is constant, neither solids nor liquids affect the equilibrium of a reaction, and therefore solids and liquids do not need to be taken into account when calculating the equilibrium constant. For example, for the following reaction, the equilibrium constant is determined as shown:

Even though the pure substance (either solid or liquid) is not included in the equilibrium constant calculation, it must be present in the system for the equilibrium to exist [2].

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Le Châtelier’s Principle
Once equilibrium is established, the reaction can be shifted in either the forward or reverse direction by applying a stress to it [1]. This concept is known as Le Châtelier’s principle after the French chemist Henry Louis Le Châtelier. He stated that “a system to which a stress is applied tends to change so as to relieve the applied stress” [3]. Therefore, by determining where stress is applied in a reaction, one can determine whether a reaction at equilibrium will shift in the forward or reverse direction. An applied stress may include a change in concentration, temperature, volume, or pressure.

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Changes in Concentration
A reaction at equilibrium wants to stay at equilibrium, so a change in concentration shifts in the direction that will restore the reaction to equilibrium. For example, consider the following reaction:

If the concentration of reactant A and/or B is increased, the reaction shifts toward the products C and D; in other words, more products C and D are formed. This occurs because as the concentration of A and/or B increases, there is more reactant available to undergo a reaction; therefore, more products are produced. As a result, the reverse reaction begins to occur because more of the products are present [4]. This occurs until the reaction’s equilibrium is restored. The opposite occurs when the concentration of C and/or D is increased. Now, the reaction shifts toward A and B for the same reasons as when the concentration of one or more of the reactants was increased. Conversely, when the concentration of A and/or B are decreased, or when either A or B is removed from the reaction, the reaction again shifts in the direction to restore equilibrium. Here, the reaction will shift towards reactants A and B to re-establish both concentrations and restore the reaction to equilibrium. Once again, the opposite occurs when the concentration of C and/or D is decreased, or when either C or D is removed from the reaction.

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Changes in Temperature
Changes in temperature can also affect the equilibrium of a reaction [1]. To determine whether the reaction will shift in the forward or reverse direction to restore equilibrium, one must determine whether the reaction is exothermic or endothermic. In an exothermic reaction, the heat produced is considered a product. For example, in the following reaction:


In an endothermic reaction, the heat absorbed is considered a reactant [2].

Therefore, if the “concentration” of the heat is decreased, the equilibrium shifts towards the side of the reaction containing the heat. On the other hand, when the “concentration” of the heat is increased, the equilibrium shifts away from that side [4]. For example, if the second reaction occurred in an ice bath, the temperature of the reaction would decrease and, as a result, the reaction would shift to the left to replace the lost heat and restore the equilibrium. However, if the second reaction occurred in a boiling water bath, the temperature of the reaction would increase and the reaction would shift to the right. This occurs because an increase in the heat “concentration” is similar to an increase in the reactants concentration and, to restore equilibrium, the reaction should shift in the forward direction [4].

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Changes in Pressure and Volume
Reactions that involve gases may be influenced by changes in either pressure or volume because gases (unlike liquids and solids) are highly compressible [3]. When a reaction occurs in a system with a constant temperature, changes in pressure and volume are inversely related. (Recall Boyle's Law relating pressure and volume.) This means that an increase in pressure causes a decrease in volume and vice versa. For example, consider the following reaction [1]:

The left side of the reaction has 4 moles of gaseous molecules while the right side has only 2 moles of gaseous molecules. Therefore, when the pressure of the system is increased, the reaction will shift toward the right because there are fewer moles of gaseous molecules on the right side and, hence, less volume. However, when the volume of the system is increased, the pressure decreases and the reaction shifts toward the left to restore the reaction to equilibrium [1].

Le Chatelier’s principle is an important concept in laboratory work. By understanding how a reaction proceeds to equilibrium and the factors that affect its equilibrium, one can control whether or not the forward or reverse reaction occurs and, therefore, whether more products or reactants result.


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[1] Eyring, Henry; Eyring, Edward M. Modern Chemical Kinetics; Reinhold Publishing Corporation: New York, 1963; pp 4, 36, 54.
[2] Jones, Loretta; Atkins, Peter. Chemistry: Molecules, Matter, and Change, 4th ed.; W.H. Freeman and Company: New York, 2000; pp 620-622.
[3] Rothstein MD, Rochelle, ed. MCAT Comprehensive Review. Simon & Schuster: New York, 2000; pp 333-336.
[4] Seager, Spencer L.; Slabaugh, Michael R. Chemistry for Today: General, Organic, and Biochemistry, 3rd ed. Brooks/Cole Publishing Company: New York, 1997; pp 241-247.