Biochem.physio

Introduction

The earlier definition said that an acid or a base is a substance that can donate proton or hydroxyl group (NH₃ is base but does not have a hydroxyl group) The definition that is used is that given by Bronsted and Lowry: An acid is a substance that can donate protons or H+ into solution while a base is a substance that can accept protons Many biological macromolecules bear numerous functional groups that determine whether they are acids or bases as well as how they will behave in certain media Under this definition in every acid base reaction, an acid HA reacts with a base H₂O to form a conjugate base of an acid and a conjugate acid of the base HA + H₂O -- H₃O+ + A^- this can also be written as HA  H⁺ + A^- CH₃COOH + H₂O - CH₃COO^- + H₃O⁺ NH₃ + H₂O  NH₄⁺ + OH^- Another definition says that an acid is a substance that can donate an electron pair while a base is a substance that can accept an electron pair. This definition is too broad especially for our biological systems and applies to non-aqueous situations.

Acids and Bases Acids and bases can be classified either as strong or weak The number of protons released determines the strength of the acid Strong acids release all their protons while weak acids don’t release all their protons ie do not ionize completely Strength of an acid is specified by its dissociation constant K K is a measure of the relative proton affinities of the HA/A- and H₃O/H₂O conjugate acid base pairs

K = [H3O+][A-]/[HA][H2O]

Since in aqueous solution the concentration of water is essentially constant the H2O is combined with the dissociation constant giving K = [H⁺][A^-]/[HA] Accordingly acids with K less than 1 (K < 1) i.e. K less than that of H₃O⁺ are weak acids Strong acids have dissociation constants larger than that of H₃O⁺, almost completely ionize

pH Ionization of Water Proton hopping in water Sometimes the hydrogen of one water molecule will "jump" to another water molecule: H₂O + H₂O H₃O⁺ + OH^- This proton hopping is called the ionization of water (an ion is a positively or negatively charged atom or molecule). This ionization creates a H3O+ and a OH^- molecule. The H₃O⁺ is often written as simply H⁺. This is because a H₃O⁺ is just a H⁺ that jumps from one water molecule to another: H₂O H⁺ + OH^- So remember, H₃O⁺ = H

pH How much H⁺ and OH^- exist in water? Very, very little! The ratio of either H+ or OH- to H2O in neutral water is 1:1,000,000,000! Since this is such a small amount of either H+ or OH-, they rarely meet and neutralize each other. The equilibrium constant, Keq describes the ionization equilibrium of water: Keq = [H⁺][OH^-] Because of this relationship it is important to note that if the [H⁺] goes up then the [OH^-] must go down, and vice-versa, for the value for the Keq of water must remain constant. For neutral water, the Keq is 1 x 10^-14 M and the concentrations of [H⁺] and [OH^-] are each 1 x 10^-7 M . Let's look at that last number without the exponent: 0.0000001 M This is obviously a very small number. A more manageable way to discuss small numbers such as this is to take the negative logarithm. For the concentration of [H⁺], this is called the pH. In this case: -log(0.0000001 M) = 7

pH The pH of a solution is simply the negative logarithm of [H⁺]. The pH of a solution describes the acidity of a solution. Acidic solutions are those with a pH of less than 7 and basic solutions have a pH greater than 7. A solution, like H₂O, with a pH = 7 is neutral. Similarly, the pOH could be used to describe a solution in terms of its OH- concentration. pOH is the negative logarithm of the OH- concentration. One useful thing to remember is: pH + pOH = 14. In the body, the pH of blood is 7.4. This corresponds to a [H⁺] of about 40 nM. This value can only vary from 37 nM to 43 nM without serious metabolic consequences.

pKa In living systems, much of the chemistry involves interactions between acids and bases. Acids are H⁺ donors and bases are H⁺ acceptors. HA H⁺ + A^- Acid and base reactions are made up of conjugate acid-base pairs. Strong acids are those that readily give up a H⁺. Strong bases readily accept a H⁺. Weak acids do not readily give up a H⁺, but will under the right conditions. A weak base is one that does not easily take up a H⁺. The conjugate base for a strong acid is a weak base. In contrast, a strong base has a weak acid as its conjugate. However, what about strong base-weak acid conjugate pairs? A^- + H⁺ HA

pH and pKa Think of it in similar terms: a person in dire need of a wig will grab just about anything. This is our strong base. Now on the right side of the equation we have someone with a wig who really doesn't want to give it up, just as a weak acid does not want to part with its H+.

• We can think of a strong acid or base as having a "flip side", which is its weak base or weak acid conjugate, respectively.

How readily an acid gives up its H+ is expressed by the acid dissociation constant, or Ka: Ka = [H+][A-] /[HA] The pKa is the negative logarithm of the Ka. Strong acids have small pKas. Looking at the above equation, it can be seen that when [A-] = [HA], then Ka = [H⁺]. Then the pKa = pH. This is the basis for the Henderson-Hasselbalch equation: pH =pKa + log[H⁺][A^-] /[HA]

 The Henderson-Hasselbalch�      Equation

Titration Curves Titration curves of both a strong and weak acid with a strong base When acids or base are added to a solution, the pH changes. The controlled addition of an acid or base to a solution is called a titration. A titration curve can be made by plotting the pH changes against the volume of acid or base added to a solution. Many titration curves are made by plotting pH on the y axis and the volume of base added to the solution on the x axis. For example, pH changes to acetic acid, CH₃COOH, and hydrochloric acid, HCl, with 0.1M sodium hydroxide, NaOH, is shown to the left. The CH₃COOH curve starts out at a higher pH than the HCl curve. This is because CH₃COOH is a weaker acid than HCl. As NaOH is added to the solution of CH₃COOH, the pH rises. However, the addition of NaOH to the HCl solution does not significantly change the pH. It takes less NaOH to change the pH of a weak acid solution than it does to change the pH of a strong acid solution. Eventually all of the acid in both solutions will have reacted with base, such that adding more base quickly increases the pH. This is called the equivalence point. This is the point where all the acid has reacted with the base, such that further additions of base quickly raise the pH. The equivalence point for a strong acid-strong base titration is pH = 7. For weak acids-strong base titrations, the equivalence point is pH >7. Buffers The bicarbonate blood buffering system As shown above, the pH of a solution is dependent on the concentration of H⁺ ions. Addition or removal of H⁺ ions, then, can greatly affect the pH of a solution. In the body, the pH of cells and extracellular fluids can vary from pH 8 in pancreatic fluid to pH 1 in stomach acids. The average pH of blood is 7.4, and of cells is 7 - 7.3. Although there is great variation in pH between the fluids in the body, there is little variation within each system. For example, blood pH only varies between 7.35 - 7.45 in a healthy individual. Large changes in pH can be life threatening. How does the body maintain a constant blood pH? The body uses a buffer system to withstand changes in pH. Buffers are made up of a mixture of a weak acid with its conjugate base or a weak base with its conjugate acid. Remember that an acid donates a H⁺. A weak acid does not donate its H⁺ as easily. Similarly, a weak base will not accept a H⁺ as well as a strong base. Buffers Buffers maintain pH by binding H⁺ or OH^- ions. This stabilizes changes in pH. The bicarbonate buffer system maintains blood pH near pH 7.4. The carbonic acid, H₂CO₃, in the blood is in equilibrium with the carbon dioxide (CO₂), in the air. Buffers are most effective in a pH range near its pKa. This is where the titration curve is most shallow, and where the pH is least affected by added acid or base. For example, in the titration curve of phosphoric acid (another blood buffering system): pKa1 pKa2 pKa3 H₃PO₄

H₂PO₄^- + H⁺

HPO₄-²⁺ H⁺

PO4^-₃ + H⁺ Buffers The titration of a polyprotic weak acid The region in the rectangle is the buffering region of biological importance. Note that the slope of the curve is shallow within one pH unit of the pKa value. This is where the buffering range is most effective, meaning that the buffer is able to resist changes in pH.

Buffers A buffer - solution that resists changes in pH when acid or base are added to it. Consists of a solution of a weak acid and the corresponding conjugate base. One can also state it consists of a solution of a weak base and its conjugate acid. May be prepared from either a weak acid or a weak base. For example, a 0.05 M acetate solution at pH 5.0 is a buffer. It is prepared from acetic acid by converting some of the acetic acid to acetate ion with strong base. The conjugate weak acid/weak base pair is acetic acid/acetate ion. The word "acetate" in acetate buffer refers to both the conjugate acid (in this example, acetic acid) and the conjugate base (acetate ion). These two species together make up the acetate buffer. Each species is present in the buffer at a specific concentration and the concentration given for the acetate buffer is the sum of the concentrations of these two species, in this example 0.05 M. The ratio of the concentrations of the two species in the buffer is obtained from the Henderson-Hasselbalch equation.