Lecture 4: Water Ionization and Buffers

Ionization of Water

Expressed by an Equilibrium constant

Keq= products / reactants = [H-] [OH-]/[H2O]

In pure water at 25 C the concentration is 55.5 M

Ion Product of Water

By substituting 55.5 M in the eq constant you get

Keq= [H-][OH-]/[55.5]

Kw= (55.5)(Keq)= [H][OH]

In pure water Keq= 1.8×10^-16

Neutral pH

Exactly equal concentrations of H and OH-

Ex: Pure Water

pH Scale

Designates the H and OH- concentrations

pH is based on the ion product of water

pH= log [1/H] or -log[H]

pH greater than 7

Alkaline or basic

Concentration of OH is greater than the concentration of H

pH less than 7

Acidic

Concentration of H is greater than that of OH-

Acidosis

pH of blood plasma below the normal value of 7.4

Common in sever unregulated diabetes

Life threatening

Alkalosis

pH of blood plasma is above the normal value of 7.4

This is found in alcohol abusers

Life threatening

Dissociation Constants

Weak acids and bases have characteristic acid dissociation constants

Conjugate acid and base pair are proton donor and its corresponding proton acceptor

The stronger the acid, the greater its tendency to lose its proton

Ionization Constants

Tendency for any acid HA to lose a proton and form its conjugate base A- is defined as the Keq or Ka

Stronger Acids have larger ionization constants

pKa

Analogous to pH and is defined as the log1/Ka or -logKa

The stronger the tendency to dissociate a proton (strong acid) the lower the pKa is
This can be determined experimentally

Titration Curves

Reveal the pKa of weak acids

A plot pf pH against the amount of NaOH added

Acetic Acid Equilibrium Constants

There are 2 equilibrium present one of water and its ions and one of acetic acid and its conjugate base

Titration Curve of Acetic Acid

At the midpoint, the pH of the equimolar solution is at the pKa of acetic acid

The amount of acetic acid and its conjugate base in solution are the same, this is also the same as the pH of acetic acid

The buffering region is where the acid and its conjugate base acts as a buffer

Buffers and Buffering Systems

Mixture of weak acid and their conjugate base

Aqueous systems that tend to resist change in pH when small amounts of acid or base are added

Buffering systems consist of a weak acid and its conjugate base

Henderson Hasselbalch Equation

Relates pH, pKa and buffer concentration

This describes the shape of the titration curve of a very weak acid

pH= pKa + log [A]/[HA]

Bicarbonate Buffer System

Acts in the blood plasma

H2CO3 → H+ and HCO3-

K= [H+][HCO3-]/[H2CO3]

H2CO3 acts as a proton donor and HCO3- acts as a proton acceptor

pH of a Bicarbonate Buffer System exposed to a gas phase

Depends on the concentration of HCO3-

The partial pressure of CO2 (pCO2) = the concentration of CO2 in the gas phase

Buffer system is effective near pH of 7.4

Bicarbonate Buffer System and Three Reversible Equilibria

The rate of respiration (controlled by the brain stem) can quickly adjust these equilibria to keep the blood pH nearly constant

Hyperventilation raises the pH of blood

Carbonic Acid

CO2 +H2O → H2CO3

K=[H2CO3]/ [CO2][H2O]

CO2 dissolved in an aq solution is in equilibrium with CO2 gas and has an K as well

Equilibrium Constant for the Hydration of CO2

The rapid equilibrium of aqueous CO2 dissolved in blood forms additional H2CO3 for buffering

CO2 + H2O → H2CO3

K = [H2CO3]/[[CO2]

Phosphate Buffer System

The phosphate buffer system acts in the cytoplasm of all cells

H2PO4- → H+ and HPO4 2-

H2PO4- acts as a proton donor and HPO4 2- acts as a proton acceptor

Buffer System is maximally effective at a pH close to its pKa of 6.86 and works over the range of pH of 5.9 and 7.9