You have probably heard the word buffer - it is one of those words that has different meanings in “real” life” and in chemistry. In everyday usage buffer means a safeguard or a barrier - something that provides a cushion or shield between you and something harmful. But in chemistry (and biology) the meaning of buffer is different and quite specific - a buffer is a solution that resists changes in pH. As we will see this ability is critically important to living systems. Many reactions are affected by changes in pH. For example, strong acid and base solutions are harmful to living tissue because they cause rapid hydrolysis of the bonds that hold you together.

9.1 Reaction systems
9.2 Buffered systems
9.3 Activation energy

That is, acidic or basic solutions can speed up the reactions in which the bonds are broken (dead bodies are often disposed of in murder mysteries - and sometimes in real life - by dissolving them in strong acid or base).

Even relatively small fluctuations in pH can cause devastating changes in the rates (and types) of reactions that are continuously occurring in living systems - the major reason being that proteins have many weak acidic and basic groups. As pH changes, the charges on these groups also changes, leading to changes in protein structure and function.

Question to answer:

• What factors might make reaction being sensitive to pH?
• Why is protein structure and activity sensitive to changes in pH?
• Which parts of proteins do you think might be affected by changes in pH - what kind of chemical properties would they have to have?
• What might these “bits” of proteins look like? (what groups of atoms would they contain?)
  

Questions to ponder:

• Would you expect nucleic acids to be more or less sensitive to pH changes than proteins?

The aqueous solution chemistry of living systems is terrifically complicated, however, we can simplify the systems that we investigate to understand how buffers work by looking at simple chemical buffer systems. Let us start by considering what would happen if we took 0.10 moles of hydrogen chloride gas and dissolved it in enough water to make one liter of solution. The resulting 0.10 M solution of hydrochloric acid HCl (aq) would have a pH of 1 (pH = – log (0.10), = – log (1.0 x 10^–1). That is the pH of the solution would change from 7 (for pure water) to 1, a change of 6 orders of magnitude in [H+]. Now if we do the same experiment adding 0.10 mol HCl(g) to an appropriately buffered solution, we would find the pH of the resulting solution would not change much.

To understand how this happens we have to review some acid-base chemistry, and look at what happens when acids and base react, what the products are, and how those products behave. As we just calculated the pH of a 0.10 M strong acid is 1.0. It would not matter which strong acid we chose - as long as it only had one proton to donate, so the pH of solutions HCl, HBr, and HClO₄ would all be the same - because they are all almost completely ionized in aqueous solution. However if we take a weak acids, like acetic acid (CH₃COOH), hydrogen fluoride (HF), phosphoric acid (H₃PO₄) or carbonic acid (H₂CO₃), the pH of each will differ, and none of them will be as low as the strong acids, because they do not ionize completely in solution. So for example, the pH of 0.10 M acetic acid is ~ 2.9, since the concentration of H+ is lower than in 0.10 M HCl. While that might not seem much different from a pH of 1, remember the [H+] is 10^–1 or 0.1 M for a pH of 1 and 10-2.9 or 0.0012 M for a pH 2.9.

Now if we look at the conjugate bases of weak and strong acids we will see an analogous difference in their behavior to produce solutions with different pHs. The conjugate base of HCl is Cl– (the chloride ion). However since we can’t just get a bottle of chloride (we need a counter ion for charge balance), we will use sodium chloride, NaCl, since we know that sodium ions are not reactive (they are usually spectator ions). If we measure the pH of a solution of NaCl, we will find that, just like water, it is 7. Neither the sodium ion, nor the chloride ion affect the pH. However if we take the corresponding conjugate base from acetic acid, for example sodium acetate (CH₃COONa), we find that a 0.1 M solution has a pH of about 9. Now this is quite surprising at first glance.

Sodium acetate belongs to the class of compounds that we generically label as “salts”. In everyday life we use salt to mean sodium chloride, but in chemistry the term salt refers to compound that contains the conjugate base of an acid, and a cation. While it might be tempting to think of all salts being rather innocuous and unreactive (like sodium chloride) it turns out that both components of the salt (the conjugate base anion, and the cation) can affect the properties of the salt - even in a simple reaction like dissolving in water. In fact, the pH of any conjugate base of a weak acid tends to be basic.

Let us investigate a bit further. This observation implies that the acetate ion (CH₃COO^–) must be reacting with water to produce hydroxide (since we already know the Na+ does not react with water). This reaction is called a hydrolysis reaction. The name is derived from the Greek words for water (hydro) and to break or separate (lysis) - it refers to reactions in which water is one of the reactants.
CH₃COO^– (aq) + H₂O(l) ↔ CH₃COOH (aq) + –OH(aq)

The production of hydroxide increases [–OH], which in turn, affects [H+] since the two are related by the equilibrium expression [H+] [–OH] = 1 x 10^–14 = K_w.

That is: when the salt of a weak acid (that is its conjugate base) is dissolved in water, a weak base is produced - and that weak base has all the properties of any base - that is, it can react with an acid - as we will see.

It is possible to calculate the pH of solutions of weak bases, just as it is to calculate pH of weak acids, if you know the acid equilibrium constant. However, what is more interesting is what happens when a solution contains significant amounts of both a weak acid and its conjugate base. For example if we take a solution that is 0.10 M in both acetic acid and sodium acetate, we can calculate what the pH would be by setting up the equilibrium table

                                     AcOH + H₂O ↔ H₃O+    + ^–OAc    + Na+ 
Initial concentrations 0.10 M               1 x 10^–7    0.10 M    0.10 M

Note that even though acetate is present in the initial mixture we have put it on the product side because once both acetic acid and acetate are present in the same solution their concentrations are “linked” - they become part of an equilibrium system that can described by the equilibrium constant for acetic acid. If the concentration of one species is changed, the other must respond. For example: recall in Chapter 8 we looked at what happens to the pH of a solution of acetic acid when acetate ion is added. The presence of acetate affects the position of equilibrium for the acetic acid dissociation, and instead of a pH of 2.9 (for 0.10M acetic acid) the concentration of solution that is 0.10 M in acetic acid and sodium acetate is 4.7. The presence of the common ion acetate has suppressed the ionization of acetic acid. We can calculate the pH of the resulting solution by adapting the expression for the acid dissociation equilibrium:

Ka = [H+][AcO^–]/[AcOH]

We are going to ignore any reaction with water from both the acetic acid and the acetate ion because they will not affect the pH significantly, since both are relatively “weak”. Even if we take these reactions into account it will not change the answer we get. Substituting in the equation for Ka we get

Ka = 1.8 x 10^–5 = [H+] (0.10)/(0.10).

Alternatively we can take negative logs of both sides giving us:

pKa = pH – log [AcO–]/[AcOH]      or      pH = pKa + log [AcO–]/[AcOH]

This equation is known as the Henderson-Hasselbalch equation; it is a convenient way to calculate the pH of solutions that contain weak acids and their conjugate bases (or weak bases and their conjugate acids).

Question to answer:

• What would be the pH of a buffer system if the concentration of the acid component is equal to the concentration of its conjugate base?
• Do you think that a particular buffer system can buffer any pH? For example could an acetic acid/acetate system effectively buffer a pH of 9?
• What criteria would you use to pick a buffer system for a particular pH?
  

Recall that a buffer can resist changes in pH - and the question is - how exactly does this happen? Let us take a closer look. Imagine we have a buffer solution that is 1.0 M in both acetic acid and acetate. The pH of this system is – log 1.8 x 10^–5 = 4.74 (since [AcO–] = [AcOH]). Now let us add some acid to this buffer; to make calculations easy we will add 0.01 mol HCl to 1.0 L of buffer solution. What happens? The major species in the buffer solution are acetic acid and acetate and water, (hydronium ion and hydroxide ion are minor components), and we have to decide which one will react with HCl(aq). Just as in any acid-base reaction, it is more likely that the base will react with the acid - that is the acetate part of the buffer will react with the H+.

The resulting reaction will be: H+ + –OAc ↔ AcOH + H₂O

That is the acetate concentration will decrease and the acetic acid will increase. We can write the reaction down and calculate the initial (before reaction) and final (after reaction) concentrations.


The pH of this system can be calculated from the Henderson-Hasselbalch equation:
pH = pKa + log (0.99/1.01) = 4.73.

That is the pH has hardly budged! (recall that the pH of 0.01 M HCl is 2.0). Even if we add more acid, for example 0.1 mol HCl to our liter of buffer, the resulting pH will not change much (it will be pH = 4.74 + log (0.90/1.10) = 4.65). Note that the addition of acid has moved the pH in the direction we would expect. That is: the addition of acid has made the pH of the solution slightly more acidic (with a lower pH) but nowhere near the pH of the solution if we had added the HCl directly to 1L of water.

We can also think about what happens when we add a strong base to the buffer solution. If we add 0.01 mol sodium hydroxide to our liter of buffer, we can calculate what the new pH will be. Now the “active” component of the buffer will be the acid, and the reaction can be written:

HOAc + –OH ↔ AcO– + H₂O

that is the strong base will react with the weak acid. The acid concentration will fall, and its conjugate base concentration will rise.

the new pH of the solution is pH = 4.74 + log (1.01/0.99) = 4.75 - a slight increase - but hardly detectable. Note that the pH of a 0.01 M solution of NaOH would be 12.

So, buffers can keep the pH of a solution remarkably constant - which as we will see if very important for biological systems, but the question arises - just how much acid could we add to the system before the pH does change appreciably, or rather, enough to influence the behavior of the system? In biological systems, this tolerance is fairly low - changes in pH can cause a cascade of reactions that may prove catastrophic for the organism. Almost all biomolecules contain groups that are sensitive to pH. That is: they contain weakly acidic (carboxylic acid groups) or basic (amino) groups, groups that can donate or accept a proton, respectively. Because they are weak acids and basis, their behavior is highly pH dependent..

The amount of acid or base that a buffer solution can absorb is called its buffering capacity. It depends on the original concentrations of conjugate acid and base in the buffer and their ratio after reaction: that is [conjugate acid]/[conjugate base]. If you start with a buffer that has equal amounts of acid and base this ratio is equal to 1.0. As the ratio moves further away from 1.0 the pH will be affected more and more, until the pH changes out of the desired range.

Another important property of buffers is the range of pH that they can act over. As we have seen from the Henderson-Hasselbalch equation, when the concentration of acid = concentration of base, the pH of the solution = pKa of the acid, so the acetic acid/acetate buffer has a pH = 4.74. Generally, the effective buffering range is +1 or – 1 pH unit around the pKa, meaning that the acetic acid/acetate acts as an effective buffer in the range of pH 3.7 - 5.7, well within the acidic pH region. While there are biological compartments (the stomach, lysosomes, and endosomes) that are acidic, the major biological fluids (cytoplasm, blood/plasma) have pHs around 7.2-7.4. In these systems buffering is done by phosphate or carbonate buffer systems. For example, the phosphate buffer system is composed mainly of H₂PO₄^– as the proton donor (acid), and HPO₄^2– as the proton acceptor (base)

H₂PO₄^– + H₂O ↔ HPO₄^2– H₃O+

It is important to see that what counts as an acid or a base depends entirely on the reaction system you are studying. Both H₂PO₄– and HPO₄^2– can act as either an acid or base (try writing out the reactions), depending upon the pH. But at physiological pH (7.2-7.4) the predominant forms are H₂PO₄– and HPO₄^2–sup>. The pKa of the conjugate acid is 6.86 - so it makes sense that this buffer system is active in cellular fluids.

Question to answer:

• How much acid would you have to add to change the pH of a buffer that is 1.0 M in acid and conjugate base by 1 full pH unit?
• If the buffer were 0.1M in acid and conjugate base would you have to add the same amount of acid? Why or why not?
• What buffer systems would you use to buffer a pH of 4, 6, 8, 10? What factors would you take into account?
• Carbonic acid H₂CO₃ has two acidic protons. Draw out the structure of carbonic acid, and show how each proton can take part in an acid base reaction with a strong base such as sodium hydroxide.
  

Amino acids, proteins and pH

Another way to look at this idea of mixtures of acids and their conjugate bases is to think about the effect of pH on a particular acid or base (rather than thinking about how adding strong acid or base affects a buffer solution). This is particularly important in biological systems where there are many weakly acidic or basic groups that can be affected by the pH. For example, proteins contain both weakly acidic –COOH and weakly basic –NH₂ groups. At physiological pH (~7). For a simple carboxylic acid like acetic acid (and it turns out that most carboxylic acids behave in a similar way), a 1.0 M solution of has a pH of ~ 3.8. If we manipulate the pH (for example by adding a strong base), the acid reacts with the base to form an acetate ion. Based on the Henderson-Hasselbalch equation, when [acetate] = [acetic acid], the pH equals the acid’s pKa - that is 4.74. As the pH increases further, [acetate] must also increase, until by pH ~7, the [acetic acid] is very small indeed. The ratio of base to acid is about 200/1. That is at physiological pHs groups such as carboxylic acids are deprotonated and exist in the carboxylate (negatively charged) form.
Conversely, if we look at the amino group (–NH₂) of a protein, it is actually the base part of a conjugate acid base pair in which the acid is the protonated form –NH₃⁺. The pKa of an –NH₃⁺ group is typically ~9. Meaning that at a pH of 9 there are equal amounts of the protonated (– NH₂) and unprotonated (–NH₃⁺) forms. If we lower the pH,by adding strong acid for example, the concentration of –NH₃⁺ form increases as the the base form –NH₂ is protonated. At pH ~ 7 there is little of the –NH₂ form remaining.

Interestingly this means that an amino acid - as pictured here, can never exist in the predominant form where both the amino (–NH2) group and the carboxylic acid (–CO2H) exist at the same time.

The “neutral” species is in fact the one in which –NH₃⁺/–CO₂^–are present at the same time. This form is called a zwitterion, and is the predominant form at physiological pH.

A protein is composed mainly (and sometimes solely) of polymers of amino acids, known as polypeptides. In a polypeptide, all but one amino (–NH2) and carboxylic acid (–CO2H) groups are bonded together form a peptide bond, in which they are transformed into an amide group, which is neither acidic or basic.

That said, many of the amino acids found in proteins have acidic (aspartic acid or glutamic acid) or basic (lysine, arginine, or histidine) side chains. The practical implication of all interactions between these charged side chains are influenced by the pH of the environment in which a protein finds itself, and changes from the “normal” environment can lead to changes in protein structure, and in turn its biological activity. In some cases protein activity is regulated by environmental pH, in other cases, changes in pH can lead to protein misfolding (denaturation). For example: if these groups are protonated or deprotonated, the electronic environment in that region of the protein can change drastically, which may mean that the protein will not only change how it interacts with other species, but its shape may change so as to minimize repulsive interactions or produce new attractive interactions. Small changes protein shape can have profound effects on how the protein interacts with other molecules and, if it is a catalyst, its efficiencyand specificity. In fact, there are cases where environmental pH is used to regulate protein activity.

Another important buffer system is the carbonic acid (H2CO3) bicarbonate (HCO3–) buffer, which is a major buffering component of blood plasma. However this system is more complex than the phosphate buffer, because carbonic acid is formed by the reversible reaction of carbon dioxide in water:
H₂O + CO₂ ↔ H₂CO₃ and H₂CO₃ + H₂O ↔ HCO₃– + H₃O+

In this instance we have two reactions linked (or coupled) by common intermediate, which leads us into a discussion of how some systems can exist under non-equilibrium conditions, and how some reactions can occur despite the fact that they have a positive free energy change (that is, they appear to contravene the Second Law of Thermodynamics).

9.1 Reaction systems
9.2 Buffered systems
9.3 Activation energy

Question to answer:

• What would be the ratio of –NH₃⁺/–NH₂ in a solution of a protein at pH 5, pH 7, and pH 9?
• What kinds of interactions would each form participate in?
• What is the predominant form of a carboxylic acid group at pH 5? pH 7, pH 9
• What kinds of interactions would each form participate in?