Acid-Base Equilibria

acid

proton donor

base

proton acceptor

Ionization of H20

H₂O + H₂O → H₃O⁺ + OH^-

A + B → CA + CB

equilibrium water constant 

k_w  = 10^_-14 

[H+] > 10^-7

acid

[H+] < 10^-7

base

pH =

-log[H+]

Henderson-Hasselback Equation

pH = pka + log [A^-]/[HA]

if low pKa

will be deprotonated above its pKa

if high pKa

will be deprotonated below its pKa

amine groups

never have a negative charge

if more base than acid [A^-] > [HA⁺]

species will be deprotonated at neutral

if more acid than base [A^-] < [HA⁺]

species will be protonated at neutral

K_^w = 

K_eq * [H₂O]

-log[H+] - log [OH-] =

+14

k_a increases

if concentration of acid increases

if pH > pKa

base > acid

deprotonation

if pH < pka

acid > base 

so protonation!

pH

acidity of the environment

pKa

point of deprotonation and protonation

buffer range 

pKa +- 1

in proteins, COOH is deprotonated more at higher pH values

so the more basic the solution, the more negative the charge of the protein

the more acidic the solution, the more positive the charge of the protein

what is a buffer

weak acids + salts

weak base + salts

• resist change +- 1

• at +-2 the group is either fully protonated/deprotonated

pH = pKa

50% deprotonated and 50% ionized

pI (isoelectric point)

pH where the average charge on the molecule is zero

• average of the two pKa values surrounding the isoelectric species

pI = pKa + pKb/2 

pKa of H₃N+

9.6

pKa of COO-

2.3

increase the pH of an amino acid

more negatively charged the amino acid will be

decrease the pH of an amino acid

more positively charged the amino acid will be

strong acids disassociate fully

weak acids don’t

determining pH of a weak acid

k_a= [H⁺][A^-]/[HA]

when to ignore x in weak acid disaasociation

1. x ➗initial conc is <5%

2. c ➗ ka > 100

3. eliminate x when k <= to a number x 10^-4