NC State University

1 Practical aspects of buffers Chemistry 201 NC State University 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 acid and conjugate base to the solution in a defined proportion. Method 1 Method 2 Add a strong acid 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. 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.

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. 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. 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.

3 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. 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 acid and Its conjugate base to a solution or by titrating in strong base to acid (or vice versa).

4 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. 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. 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.

5 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 acid-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. The equivalence point is reached when the added is equal to the original acid 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.

6 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). 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 acid crystals under polarized light Citric acid is found in abundance in citrus fruits.

7 It is also part of the citric acid cycle in biochemistry. It is a triprotic acid, with three carboxylic acid 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 acid 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. Amino acid side chain pKa Proteins are made of up of a number of titratable amino acids. At pH 7 the terminal carboxyl, aspartate and glutamate have a negative charge. Terminal amino, lysine and arginine are positively charged. Others are neutral, but can be charged due to interactions within the protein.

8 The pKa of any of these groups may be altered by the protein. The isoelectric point of a protein pI Proteins have many titratable groups on their surface. It is not possible to define a single pH since all of the groups have different pKa values. However, we can define the point at which protein is neutral in charge: the isoelectric point. At the isoelectric point the protein has 0 net charge, which means that there as many positive as negative groups on the surface. The isoelectric point concept applies to polymers, nanoparticles etc. Any macromolecule can be described in terms of its overall charge. IMPORTANT: When pH = pI a macromolecule has a neutral surface. This is the minimum stability point. Macromolecules tend to precipitate at this poin. Hemoglobin is the most abundant protein inside of red blood cells and accounts for one-third of the mass of the cell. During the conversion of CO2 into HCO3-, H+ liberated in the reaction are buffered by hemoglobin, which is reduced by the dissociation of oxygen.

9 This buffering helps maintain normal pH. The process is reversed in the pulmonary capillaries to re-form CO2, which then can diffuse into the air sacs to be exhaled into the atmosphere. Hemoglobin as a buffer As with the phosphate buffer , a weak acid or weak base captures the free ions, and a significant change in pH is prevented. Bicarbonate ions and carbonic acid are present in the blood in a 20:1 ratio if the blood pH is within the normal range. With 20 times more bicarbonate than carbonic acid, this capture system is most efficient at buffering changes that would make the blood more acidic. This is useful because most of the body s metabolic wastes, such as lactic acid and ketones, are acids. Carbonic acid levels in the blood are controlled by the expiration of CO2 through the lungs. Bicarbonate (hydrogen carbonate): an important regulator in the body The role of carbonic anhydrase The enzyme does not change the equilibrium, but it accelerates the Rate of reaching the equilibrium on each side of a membrane.

10 Acid/Base Mixtures : Reactions How do you calculate pH after an acid/base reaction occurs? Text : Section What is the pH when: a) 25 mL of M HCl are added to 35 mL of M NaOH? b) 15 mL of M HClO4 are added to 25 mL of M NaOH? Examples: Strong acids and bases What is the pH when 25 mL of M HCl are added to 35 mL of M NaOH? Examples: Strong acids and bases What is the pH when 25 mL of M HCl are added to 35 mL of M NaOH? Examples: Strong acids and bases Step 1. Calculate dilutions. First add the volumes Total volume = 25 mL + 35 mL = 60 mL Calculate concentrations in the solution What is the pH when 25 mL of M HCl are added to 35 mL of M NaOH? Examples: Strong acids and bases Step 2. Write a balanced chemical reaction for the limiting reaction and the excess reaction. Limiting reaction Species HCl NaOH Na+ Cl- Initial Difference -x -x x x Final x x What is the pH when 25 mL of M HCl are added to 35 mL of M NaOH?