Acid-base buffer systems Henderson-Hasselbalch equation

Buffers stabilize a solution at a certain pH. This depends on the nature of the buffer and its concentration. For example, the carbonic acid-bicarbonate system has a pH of 6.37 when the two ingredients are at equimolar concentration. A change in the concentration of the carbonic acid relative to its conjugate base can shift the pH of the buffer. The Henderson-Hasselbalch equation below gives the relationship between pH and concentration. 

There are two ways of dealing with the bicarbonate buffer system. The first uses the Henderson-Hasselbalch equation and an effective pKa of 6.1. If there is more base (HCO 3) than acid (C02), the pH will always be bigger than the pKa. This is usually the case physiologically (pH = 7.4 pKa = 6.1) so that on a molar basis there is always more than 10-fold more HCO 3 than C02. 

The common-ion effect is an application of Le Chatelicr s principle to equilibrium systems of slightly soluble salts. A buffer is a solution that resists a change in pH if we add an acid or base. We can calculate the pH of a buffer using the Henderson-Hasselbalch equation. We use titrations to determine the concentration of an acid or base solution. We can represent solubility equilibria by the solubility product constant expression, Ksp. We can use the concepts associated with weak acids and bases to calculate the pH at any point during a titration. 

The most effective buffering system contains equal concentrations of the acid, HA, and its conjugate base, A-. According to the Henderson-Hasselbalch equation (2.6), when [A-] is equal to [HA], pH equals pKa. Therefore, the pKa of a weak acid-base system represents the center of the buffering region. The effective range of a buffer system is generally two pH units, centered at the pKa value (Equation 2.9). 

HA] is the concentration of the acid and [A-] is the concentration of the conjugate base. The pKa of the carbonic acid-bicarbonate system is 6.37. When equimolar conditions exist, then [HA] = [A ]. In this case, the second term in the Henderson-Hasselbalch equation is zero. This is so because [A ]/[HA] = 1, and the log 1 = 0. Thus at equimolar concentration of the acid-conjugate base, the pH of the buffer equals the pKa in the carbonic acid-bicarbonate system this is 6.37. If, however, we have ten times more bicarbonate than carbonic acid, [A ]/[HA] = 10, then log 10 = 1 and the pH of the buffer will be... 

Buffers are defined as substances that resist changes in the pH of a system. All weak acids or bases, in the presence of their salts, form buffer systems. The action of buffers and their role in maintaining the pH of a solution can best be explained with the aid of the Henderson-Hasselbalch equation, which may be derived as follows. 

The buffer equation is also known as the Henderson-Hasselbalch equation. This equation is in principle just another version of the expression for Ka but it may nevertheless other be easier to apply. Using the buffer equation one must remember that HA and A denotes the corresponding acid-base pair and that pKa is the acid exponent of the acid (HA). In the following example the buffer effect is illustrated in a buffer system consisting of equal amounts of acetic acid and acetate into which strong base is added. 

Henderson-Hasselbalch equation is important for imderstanding buffer action and acid-base balance in the blood and tissues of the mammalian system. The equation is derived in the following way. Let us denote a weak acid by the general formula HA, and its salt by the general formula BA (B being the metal ion and A being the conjugate base). The sedt dissociates completely, while the weak acid dissociates only partly. We can write the equilibrium reactions for the dissociation of HA and BA in the buffer solution as follows ... 

Henderson-Hasselbalch Equation n A formula relating the pH value of a solution to the pK value of the acid in the solution and the ratio of the acid and the conjugate base concentrations pH = pIQ + log( [A—] / [HA]), where [A—] is the concentration of the conjugate base and [HA] is the concentration of the protonated acid. For the bicarbonate buffer system in blood,... 

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