Question

Fundamentals of Biochemistry Tutorial: Buffers

Answer 1

new year, new tutorials. unfortunately the two worst trolls on QC are back but I'll try to post tutorials when they're offline so I don't get harassed/spammed/plagiarized.

Answer 2

\({\bf{The~Basics:}}\) aqueous systems that resist pH changes upon addition of strong acid or base
note: for now we will be focusing on the Brønsted-Lowry definition of acids/bases

- generally the cytosol maintains a pH of about 7 with some notable exceptions (will get into this during later tutorials)
- buffer zone: range where adding acid/base does not change the pH significantly. on a plot of (volume added) vs. pH this will appear as a mostly horizontal region
- midpoint of buffer region: conc. of proton donor = conc. of proton acceptor; region where the buffer is the strongest



added [H+] will react with the conjugate base A-, forming a weaker acid
added [OH-] will react with the acid, forming water
equilibrium constant of the weak acid/weak base system is maintained

Answer 3

The Henderson-Hasselbach equation can be used to calculate the pH of the buffer system after a certain amount of acid/base is added. Use stoich the calculate how much of the acid and base are left, taking into account the change in volume of the whole system.

- most buffering is based on weak acid/weak base systems but nucleotides, metabolites, and special compounds like ammonia can also contribute to buffering

- one very important system to remember is the bicarbonate buffers system, which maintains pH in blood

three equations to remember:
H2CO3 (Carbonic acid) <--> H+ + HCO3- (bicarbonate ion)
dissolved CO2 + H2O <--> H2CO3
gaseous CO2 <--> dissolved CO2

two conditions to remember:
1. when [H+] increases, more H2CO3 is produced (le chatelier's principle), which in turn increases the conc. of dissolved CO2, increasing the conc. of gaseous CO2, with the end result being exhalation
2. when [H+] decreases, follow equilibrium in the opposite direction, resulting in more gaseous CO2 dissolving in to the plasma

hyperventilation: results in too much CO2 being lost, raising blood pH as the equilibrium shifts towards loss of dissolved CO2, shifting the equilibrium away from the production of [H+] via carbonic acid
- solution is generally to breathe in a paper bag because this will allow the body to recover lost CO2

Answer 4

cont. we can combine the equilibrium expressions for the bicarbonate buffer equations
using dissolved CO2 + H2O <--> H2CO3 we can make

Kh = [H2CO3]/[dissolved CO2]
or, solving for [H2CO3] we get Kh*[dissolved CO2]
(the conc. of water is basically constant so it's not included in the equation for simplicity's sake)

using H2CO3 <--> H+ + HCO3-
we can make
Keq = [H+][HCO3-]/[H2CO3] for acid dissociation
making the substitution for [H2CO3] we get
Keq = [H+][HCO3-]/ ( Kh*[dissolved CO2] )

if we combine the equilibrium constants (thus, combining the equations to get the overall reaction) we simply multiply Keq * Kh to get

K comb. = [H+][HCO3-]/ ([dissolved CO2])
using literature values this gives us a pKa of about 6.1

using Henderson Hasselbach, with literature values for [HCO3-]/[H2CO3] gives us a blood pH of 7.4

Answer 5

\({\bf{Real-Life~Applications:}}\)
- in diabetic individuals, insulin resistance leads to the inability of cells to reuptake glucose. their cells will instead rely on fatty-acid breakdown for metabolism, which has acidic end products. these end products are present in urine, so this is one possible test for diabetes.
- fasting/starvation also has similar effects
- extreme physical exertion, kidney failure, lung disease can also interfere with the bicarbonate buffer system

Answer 6

Source material is section 2.3 of Principles of Biochemistry 7th edition by Nelson, David L., and Cox, Michael M.