Download Buffer - Biochemistry - Lecture Slides and more Slides Biochemistry in PDF only on Docsity!
Buffer Know-how
Buffers are important in biochemical processes. Whether they occur
naturally in plasma or in the cytosol of cells, buffers assure biological
reactions occur under conditions of optimal pH. They do this by controlling the
hydrogen ion concentration of solutions. The word “buffer” is so common in
biochemistry its replaces the word “water” in experimental protocols. For
example, typical of the statement seen in publications is “the pellet was
dissolved in pH 7.5 buffer”. All this should alert you to the importance of
thoroughly understanding buffers and buffering agents. Words such as pH,
pKa, conjugate acid, conjugate base, Henderson-Hasselbalch equation are
used frequently in biochemical language and every publication that describes
an experiment performed “in vitro” ( Lat ., in glass), must include a clear
description of the buffer that was used. In this tutorial we will revisit buffers
and attempt to understand their make up and mechanism of action. We will
also give some insights into how to solve buffer problems.
Rules Governing Buffer Reactions
To help you understand buffer action, consider the equation that describes a buffer reaction ( click 1 ). The HA and A-^ represent the two components of any buffer: the conjugate acid and the conjugate base, respectively. Note that the right side and left side of the equation are the same. Rule1 gives the meaning of the reaction ( click 1 ).
Rules
HA + A-^ + H 2 O HA + A-^ + H 2 O
- Buffer reactions never go to completion.
- H+^ can only react with A-, OH-^ can only react with HA.
- The conjugate acid and conjugate base must change proportionately and in opposite directions.
- The sum of the concentrations of base and acid components stays constant, i.e, HA + A- before = HA + A- after.
Adding NaOH or HCl to this buffer would shift the reaction toward the right, but with different results. This is because of Rule 2 ( click 1 ). Both components in the buffer must change any time acid or base is added to the buffer. This is because of Rule 3 ( click 1 ). Finally, we take into account the HA and A- components, expecting to see a decline in the overall buffer. Such is not found because of Rule 4 ( click 1 ). Hence, Rule 5 summarizes an important principle that we must know (click 1).
- Only the ratio, never the total, of HA to A-^ changes as a result of buffer action.
Principle Behind Buffer Action
Buffers are composed of weak acids and their salts. A salt is the acid minus its proton. Weak acids and their salts have two properties that are important for buffering action. First, weak acids are a reserve of the protons that neutralize OH-^ and prevent the solution from becoming alkaline. Salts of weak acids are strong bases and prevent the solution from becoming acidic. Both components are needed and both are interchangeable through the loss (or gain) of a single proton.
A buffer’s power lies in its reserves ( click 1 ). A buffer is at optimal strength when there is an equal amount of HA and A- in solution as shown. This will only occur when the pH of the solution equals the pKa of the acid’s group. Adding OH- causes the buffer to respond by calling on the reserve pool of HA. A-^ is formed at the expense of HA ( click 1 ). This continues until all the excess OH- is neutralized. At the end the salt pool has increased (and the acid pool has decreased) by the same number of moles of base that were added. Click to go on.
HA
HA
HA HA HA (^) HA HA HA HA
HA
HA
Reserve acid 11 moles
Reserve salt 11 moles
A-^
A- A- A- A- A-
A-
A- A-
A-
A-
OH-
A-
A-
A-
A- A-
6 moles 16 moles
OH-
Neutralized
Focus on the Ratio of [A-]/[HA]
In the previous illustration you saw the importance of knowing the ratio of HA and A-. Now you will see that it is the ratio that determines the pH of the solution, and vice versa, the pH allows you to determine the ratio. It all begins with an equilibrium expression ( click 1 ). If we take the log of all components we derive a logarimic expression of the same equation ( click 1 ),
[H+] = Keq [HA]
[A-]
Log [H+] = Log Keq + Log
[HA]
[A-]
Mutiplying components on both sides of the equation by -1 gives ( click 1 )
- Log [H+] = – Log Keq – Log
[HA]
[A-]
Substituting pH and pK for the appropriate terms in the equation and making the log of HA/A- positive by reversing numerator and denominator gives ( click 1 )
pH = pK + Log
[A-]
[HA]
Note, in the equation, pK is a constant and A/HA is the only variable ( click 1 ). This is the Henderson- Hasselbalch equation. Click to go on.
constant variable
Test and Enhance your Understanding
Q: Why must buffers always have two components?
A: Several reasons. First, Rule 2 say that a base will only react with an acid and an acid with a base. The two components assure the buffer protects against both. Second, it is the ratio of conjugate base to conjugate acid that determines pH. Neither one alone would suffice.
Q: How does one select a buffer in a particular biochemical experiment? A: The decision is based on the desired pH that must be maintained. An acetate buffer, for example, is useless at pH 8.0 because the buffer (pK = 4.8) at this pH is almost all conjugate base. In contrast a buffer with a pK around 8.0, Tris buffer, for example, would be more suitable.
Q: Is there a way of telling when a buffer will be most effective at a given pH? A: Yes. The pK of the buffer (sometimes referred to as pKa for an acid) immediately tells you the pH that will result in equal amounts of conjugate acid and base in solution, which is the optimal condition for any buffer.
Q: Suppose you added 2 moles of acid to 10 moles of buffer in a solution where pH and the pK have the same value. What would be the result?
A: Since the pH = pK, you know that there are equal moles of conjugate acid and conjugate base, 5 moles of each. The 2 moles of acid would convert 2 moles of the base component to the acid changing the ratio of A-/HA from 1.0 to 0.43, which would lower the pH of the solution by 0.37 units.
