Buffering Capacity of Amino Acids, Biology tutorial

Ionization of Water:

Water ionizes according to the equilibrium equation below:

H2O(Water)↔ H+ (Hydrogen ion) + OH(Hydroxide ion)

The equilibrium constant for the ionisation of water is given by Keq

Keq= [H+][OH-]/[H2O] ............. Equation-1

At 25oC, the concentration of one litre of pure water is 55.5 M (Grams of H2O in one litre divided by its gram molecular weight: 1000 gL-1/18.015 gmole-1.Hint: molarity = gram concentration/molecular weight).

So, Keq = [H+][OH-]/55.5M

Or Keq(55.5M) = [H+][OH-]............ Equation-2

The product Keq(55.5 M) can be taken as a new constant called as ionic product of water (kw).

i.e. Keq (55.5 M) = kw = [H+][OH-].

By electrical conductivity measurements, Keq of pure water at 25oC is 1.8 x 10-16 M.

Now, substituting this value into equation-2 above gives

kw = (55.5 M) x (1.8 x 10-16 M)

= 1.0 x 10-14M2

This means that in aqueous solution at 25oC, the product [H+][OH-]is always 1.0 x 10-14 M2.

In pure water (neutral pH), the concentration of hydrogen ion equals the concentration of hydroxide ion ([H+][OH-]).

[H+] and [OH-] can be calculated thus:

kw = [H+][OH-]  .........................equation-3

= [H+]2 (if [H+] = [OH-])

And so you have,

√kw = [H+]

Therefore

[H+] = √1x10-14

= 1x10-7

From the equation-3, any of the concentration terms can be calculated if the other is known.

Definition of pH and Calculation of Hydrogen ion Concentration:

Definition of pH:

The pH is simple and suitable designation for concentration of hydrogen ion in aqueous solution (pOH mean hydroxide ion concentration). pH is stated as negative logarithm to base 10 of molar hydrogen ion concentration.

This can be expressed mathematically thus:

pH=-Log10[H+]

Or pH= Log10(1/[H+])

For a neutral solution at 25oC

pH = - Log 1.0 x10-7

= 7

pH-Scale:

pH scale estimates how acidic or basic a substance is. pH scale ranges from 0 to 14. A pH of 7 is neutral. pH less than 7 is acidic. The pH greater than 7 is basic.

A pH scale is logarithmic and consequently, each whole pH value below 7 is 10 times more acidic than next higher value. For instance, pH 4 is 10 times more acidic than pH 5 and 100 times (10 times 10) more acidic than pH 6. Same holds true for pH values above 7, each of which is 10 times more alkaline than next lower whole value. For instance, pH 10 is 10 times more alkaline than pH 9 and 100 times (10 times 10) more alkaline than pH 8.

pH of biological fluids are estimated roughly using indicators or more accurately with the help of pH meters.

Dissociation of Weak Acids and Definition of pKa:

Dissociation of Weak Acids:

The weak acid is one which dissociates partially in water. Acetic acid is a good example of the weak acid; in water it releases only one of its protons.

CH3COOH(Acetic Acid)↔(Keq) CH3COO-(Acetate ion) +H+ (Proton)

Equilibrium constant in above equation is a measure of tendency of acid (CH3COOH) to lose the proton and form its conjugate base (CH3COO-).

Usually, for dissociation of any weak acid (HA) equilibrium constant is written as Ka

HA(weak acid)↔(Ka)A- (conjugate base) + H+(Proton)

The Ka for acetic acid is 1.74 x10-5.

Weaker acids are known to contain smaller dissociation constants as opposed to larger values for stronger acids.

pKa of Weak Acids:

Just as for pH, pKa is stated as negative logarithm to base 10 of Ka and is expressed mathematically as:

pKa = - logKa

or

= Log 1/Ka

Generally, the stronger the tendency to lose proton, the stronger the acid, the larger its Ka and lower its pKa. Therefore, weaker acids have smaller dissociation constants (Ka) and higher pKa values.

Amino acids are weak acids also. On the titration curve, pKa values of Ionization groups exist at point of inflection of their dissociation.

Buffer Solutions:

Buffer solution is a solution which resists minor changes in pH when small amount of acid or base is added to it. The buffer system comprises of weak acid and its conjugate base. The example of the buffer is a mixture of equal concentration of acetic acid and acetate ion (its conjugate base), found at midpoint of titration curve.

i) Roles of Biological Buffers:

Buffer system in living organisms seems to be their first line of defense against any small change in pH. For instance, excessive production of acid in muscular exercise, release of toxic substances during infection caused by micro organisms or even in nutritionally deficient state (as in diabetes) could cause small change in pH. This should be resisted as fast as possible by actions of many buffers found in body.

ii) Buffer Action:

Buffers execute their action by interaction of weak acid and conjugate base with OH- and H+ respectively. The chart below is summary illustration of buffer action:

793_Buffer Action.jpg

From the chart, the significant points can be noted which are given below:

i) Buffer system is reversible;

ii) If base is added, weak acid (HA) interacts with it to give water and conjugate base;

iii) If acid is added, conjugate base interacts with it to produce weak acid and so oppose change in pH.

iii) Handerson-Hasselbalch Equation:

In laboratory, buffers are at times made using Henderson-Hasselbalch equation, which is a description of titration of all weak acids. Equation is stated mathematically as:

pH=pKa+[conjugate base]/[weak acid]

iv) Amino Acids as Biological Buffers:

Among several other biological buffers, ionisable groups of some amino acid residues in proteins could perform as buffers. For instance, histidine has pKa of 6.0. Proteins having histidine residues thus buffer effectively near neutral pH.

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