Chemical Buffers

FIELDS OF STUDY: Inorganic Chemistry; Geochemistry; Metallurgy

ABSTRACT

The nature and function of buffer solutions are presented, and their importance in biological systems is discussed. Buffer solutions resist changes in pH as a result of the equilibrium between the conjugate pair specific to a particular solution.

The Nature of Chemical Buffers

Chemical buffers are solutions that resist changes in pH. A buffer solution will generally maintain a constant pH when other materials are added, as long as the quantities of other materials do not overload the ability of the buffer solution to adjust to the change in the system components. The ability of a buffer solution to compensate for the addition of acids or bases, and thus maintain a fairly constant concentration of hydronium ions (H₃O⁺), derives from the equilibrium reaction between the particular weak acid or base and the corresponding salt used to prepare the buffer. The key feature of the system is a weak acid, one that does not dissociate well into its component ions when dissolved in water. Acetic acid is a good example. Despite its acrid odor and acidic character, acetic acid is covalently bonded (as are essentially all other carboxylic acids). A weak acid need not necessarily be a carboxylic acid but absolutely must be one that does not dissociate completely into ions when dissolved in water. When this condition is met, a dissociation equilibrium is established, according to the general equation

H₂O + HA ⇌ H₃O⁺ + A⁻

The equilibrium reaction has two essential characteristics: constant amounts of all reactants and products in a specific characteristic and a constant ratio, defined as the "equilibrium constant." For weak acids, this is the acid dissociation constant, K_a. By adjusting the relative amounts of each component in the system to correspond to the equilibrium’s constant value, an equilibrium system automatically adjusts to any effect that perturbs the state of the system. In a buffer system, the introduction of some additional H⁺ requires it to combine with an appropriate amount of A⁻ both to regenerate HA and to maintain the equilibrium constant value. By the same token, the introduction of additional alkalinity, or basicity, means that some of the HA will undergo neutralization. Regardless of the resiliency of the buffer system, however, the introduction of too much acid or base (or of a much stronger acid or base) will overpower the ability of the system to reestablish its fundamental equilibrium.

For a buffer (or any equilibrium) system to behave in this way, the thermodynamic state of the system must be at an energy minimum when it is at equilibrium. The tendency to regain the minimum-energy state is the driving force for the autoadjustment of an equilibrium system, such as a buffer solution.

Preparation of Chemical Buffer Solutions

Simply stated, a buffer solution is prepared by combining a weak acid and a salt of the weak acid in the same solution. An acetate buffer, for example, would be prepared using acetic acid and sodium acetate (or a similar acetate salt). An alkaline buffer might be similarly prepared from a solution of ammonia in water and a corresponding salt, such as ammonium chloride. The obvious characteristic of the buffer solution is that it contains either a weak acid and its conjugate base or a weak base and its conjugate acid, forming a conjugate pair.

This simple description belies the care with which a buffer solution must be prepared. To achieve a specific pH value for the solution, careful calculations are required to determine the correct amounts of a particular conjugate pair to use. Since buffer systems are used most commonly in biochemical studies, the conjugate pair must be selected according to the nature of the bioreaction system being studied. In biochemical systems, such as the cell cytosol, pH is strictly controlled by materials that are present in the system. With respect to the validity of the results, it would be inappropriate to carry out a biochemical reaction process in a buffer system that is not applicable to cell biochemistry and then apply the observed results to a biochemical system.

Another consideration when discussing chemical buffers is the mode of reaction by which the conjugate pair undergoes neutralization and dissociation reactions. A monoprotic weak acid, such as acetic acid, donates a single proton (H⁺) from each molecule in solution and functions in an entirely different manner than does a diprotic or polyprotic acid, such as citric acid, oxalic acid, or phosphoric acid. The mode of reaction also determines both the magnitude and the range of pH in which the buffer is effective. A simple comparison of the relevant equilibrium reactions make this easy to see. Oxalic acid, a diprotic acid, has two neutralization equilibriums that occur simultaneously, as

HOOC−COOH ⇌ HOOC−COO⁻ + H⁺ HOOC−COO⁻ ⇌ ⁻OOC−COO⁻

Phosphoric acid, a common component of biochemical systems and a triprotic acid, has three neutralization equilibriums to consider:

H₃PO₄ ⇌ H⁺ + H₂PO₄⁻

H₂PO₄⁻ ⇌ H⁺ + HPO₄²⁻

HPO₄²⁻ ⇌ H⁺ + PO₄³⁻

Each separate neutralization equilibrium occurs at a different pH, but all are simultaneously involved in the overall equilibrium condition of the system. The calculations required to achieve a desired pH using such conjugate pairs must take this into account.

The Henderson-Hasselbalch Equation

The [H₃O⁺] of a conjugate pair buffer solution can be calculated proportionately as

Mathematically, this is equivalent to the logarithmic form

This is known as the Henderson-Hasselbalch equation. For most applications, it works quite well for the calculation of pH at each individual step in an overall equilibrium system. Successive calculations for systems with more than one equilibrium to consider are used to calculate the overall pH for such systems. Knowing the concentrations of the acid and conjugate-base materials allows for calculation of the required masses of the two components, and when properly mixed, the system can come to equilibrium at the desired pH.

Buffers in Biological Systems

Buffer systems are important in both living systems and the laboratory study of processes that affect living systems. Enzyme-mediated processes are especially sensitive to changes in pH. The fluid components of a cell structure such as the cytosol are very strictly controlled, and complex buffer systems involve the interaction of several kinds of buffers. The various functional groups of amino acids are typically rendered nonfunctional outside of a narrow pH range. The process taking place under the control of the enzyme may result in the release of protons into the reaction medium. The effect that these would have is nullified by the surrounding buffer solution. Structural features such as disulfide bridges, necessary for the enzyme to maintain its proper structure, are sensitive to hydrolysis as well.

PRINCIPAL TERMS

• conjugate acid: the material formed from a base when it accepts a proton (H⁺) from an acid, thus gaining a unit of positive charge.
• conjugate base: the material formed from an acid when it donates a proton (H⁺) to a base, thus losing a unit of positive charge.
• dissociation constant: a characteristic value representing the extent to which a compound dissociates into component ions in a certain solvent and under specific conditions.
• neutralization: a chemical reaction between an acid and a base that results in the formation of a salt, usually accompanied by water.
• pH: a numerical value that represents the acidity or basicity of a solution, with 0 being the most acidic, 14 being the most basic, and 7 being neutral.

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