NC State University

1 Practical aspects of buffers Chemistry 201 NC < strong >Statestrong > < strong >Universitystrong > Lecture 15 The everyday pH scale To review what pH means in practice, we consider the pH of everyday substances that we know from experience. Remember that [H+] = 10-pH. pH + pOH = 14 Therefore that [OH-] = 10pH-14. Two ways to make a buffer Add the < strong >acidstrong > and conjugate base to the solution in a defined proportion. Method 1 Method 2 Add a strong < strong >acidstrong > to the weak base (or vice versa) until the desired proportion [A-]/[HA] is obtained. Buffer strength The ratio [A-]/[HA] should be as close as possible 1:1, but the amounts may vary. To make a stronger buffer you simply need to increase the amount of each component.

2 Let s investigate. Suppose we add 1 mL of 1 M HCl to 1 liter of solution. The final concentration of HCl is M. Buffer strength The ratio [A-]/[HA] should be as close as possible 1:1, but the amounts may vary. To make a stronger buffer you simply need to increase the amount of each component. Let s investigate. Suppose we add 1 mL of 1 M HCl to 1 liter of solution. The final concentration of HCl is M. pH = 3 Buffer strength The ratio [A-]/[HA] should be as close as possible 1:1, but the amounts may vary. To make a stronger buffer you simply need to increase the amount of each component. Let s investigate. Suppose we add 1 mL of 1 M HCl to 1 liter of solution.

3 The final concentration of HCl is M. pH = 3 Suppose we add 1 mL of 1 M HCl to 1 liter of 10 mM optimal phosphate buffer solution (pKa = ) solution. The final concentration of HCl is M. pH = pKa + log10([A-]/[HA]) Buffer strength Keeping in mind that the unbuffered solution in this example ([HCl] = M) would be pH = 3 Suppose we add 1 mL of 1 M HCl to 1 liter of 10 mM optimal phosphate buffer solution (pKa = ) solution. The final concentration of HCl is M. pH = pKa + log10([A-]/[HA]) [A-] = = [HA] = + = pH = + log10( ) = If the target pH = ( pH = pKa) then this buffer is too weak.

4 An error of pH units could be significant. Buffer strength Keeping in mind that the unbuffered solution in this example ([HCl] = M) would be pH = 3 Suppose we add 1 mL of 1 M HCl to 1 liter of 100 mM optimal phosphate buffer solution (pKa = ) solution. The final concentration of HCl is M. pH = pKa + log10([A-]/[HA]) [A-] = = [HA] = + = pH = + log10( ) = If the target pH = ( pH = pKa) then this buffer is reasonable. The difference is only Buffer strength Keeping in mind that the unbuffered solution in this example ([HCl] = M) would be pH = 3 Suppose we add 1 mL of 1 M HCl to 1 liter of 300 mM optimal phosphate buffer solution (pKa = ) solution.

5 The final concentration of HCl is M. pH = pKa + log10([A-]/[HA]) [A-] = = [HA] = + = pH = + log10( ) = If the target pH = ( pH = pKa) then we would say that this buffer is definitely strong enough, difference = You can create a buffer either by adding the < strong >acidstrong > and Its conjugate base to a solution or by titrating in strong base to < strong >acidstrong > (or vice versa). Remember, regardless of the method used to prepare it: The buffering strength is maximum when [HA] = [A-] The buffering range is considered to extend from [HA] / [A-] = to [HA] / [A-] = 10. This is subjective. Wertz suggests to 100 is an acceptable range.

6 Titrating to make a buffer Understanding the titration curve Starting point [HA] = [HA]0 Suppose we want to make a buffer by titrating [OH-]. We cannot use the H-H equation initially. We do not know the concentration of [A-]. Instead at this initial point we will use the other form of the equilibrium constant and make an ICE table. Added [OH-] = 0 Understanding the titration curve Starting point [HA] = [HA]0 Suppose we want to make a buffer by titrating [OH-]. We cannot use the H-H equation initially. We do not know the concentration of [A-]. Instead at this initial point we will use the other form of the equilibrium constant and make an ICE table.

7 Added [OH-] = 0 Understanding the titration curve Starting point [HA] = [HA]0 We can calculate x = [H+] and therefore the pH from the equilibrium constant. Added [OH-] = 0 Understanding the titration curve Maximum buffer capacity [HA] = [A-]. When we have a buffer we can use the Hendersen-Hasselbach equation. This is nice since it is the simplest treatment of the < strong >acidstrong > -base equilibrium. In the case shown we have pH = pKa. Added [OH-] = 1/2 [HA]0 Understanding the titration curve Maximum buffer capacity [HA] = [A-] when [OH-] ~ [HA]0 The buffer range is defined as approximately from: pH = pKa 1 R = [A-]/[HA] = [OH-] ~ [HA]0 to pH = pKa + 1 R = [A-]/[HA] = 10 [OH-] ~ [HA]0 Buffer region Understanding the titration curve Once the solution moves outside the buffer range the pH shoots up.

8 The equivalence point is reached when the added is equal to the original < strong >acidstrong > concentration, [OH-] ~ [HA]0 At this point one can no longer use the H-H equation. Instead, we assume that [A-] ~ [HA]0 Then we use the base equil- -ibrium: Note that pKb = 14 - pKa Buffer region Types of buffers There are inorganic buffers, phosphate, but there are many more organic buffers. In fact, the number of buffers is staggering. Organic buffers: Tris, HEPES, MOPS, Biological buffers: citrate, acetate, carbonate, malonate, Proteins themselves are polyelectrolytes and therefore tend to the buffer the solution they are in. This can have important physiological impact ( hemoglobin).

9 Tris buffer Tris(hydroxymethyl)aminomethane), is an organic buffer with the formula (HOCH2)3 CNH2. Tris has a pKa = The buffer range is It is is widely used as a component for solutions of nucleic acids, proteins and for any application in which phosphate is not a good choice. For example, calcium phosphate has a low solubility product, Which means that phosphate buffers are a poor choice in any application where calcium is present. Important point: Tris can react with aldehydes since it is a primary amine. Choice of buffer should done by consulation of the chemical interactions in your application. Other organic buffers HEPES: pKa = MOPS: pKa = Citrate buffer Citric < strong >acidstrong > crystals under polarized light Citric < strong >acidstrong > is found in abundance in citrus fruits.

10 It is also part of the citric < strong >acidstrong > cycle in biochemistry. It is a triprotic < strong >acidstrong > , with three carboxylic < strong >acidstrong > groups as seen by its structure (on the right). pKa1 = pKa2 = pKa3 = Citrate buffer: species in solution The middle H atom has the lowest pKa. This is because the neighbor- ing OH group has an electron with- drawing effect that stabilizes the negative charge created. Amino acids are amphipathic All amino acids contain the carboxylic < strong >acidstrong > and amino group. These have very different pKa values so the amino acids have a doubly charged form (zwitterion) at pH 7. In proteins only the N- and C-terminus have these pKas. Amino acids can be classified in part according to the pKa of their side chains.