pH and Buffering Flashcards
Define pH
pH = -log [H+]
It defines how acidic/alkaline a solution is
Acidity depends on ONLY free H+ ions- NOT on those bound to anions
What is the living range of pH in the blood?
7.0-8.0
however ideally its between 7.35-7.45
Why is regulating the blood pH important?
It is important due to proteins- the [H+] and [OH-] that can affect the bonds that stabilise the protein ie. VdW, ionic, electronic interactions
What is the definition of acidosis ?
excessively acidic condition
What is the definition of alkalosis ?
excessively alkali condition
Where does acid come from in the body?
breakdown of protein
Some acids enter in foods eg. lemon and vinegar
incomplete oxidation fats/glucose
Loading and transport of CO2 in the blood- carbonic acid
What statements can be made about concentrations at neutral pH and room temperature
[H+] = [OH-] = 10^7
How is acid-base equilibria regulated in the body?
The lungs-expelling CO2
Kidneys- regulate blood pH
Systems in the blood known as chemical buffers
What are Chemical Buffers?
Buffers resist abrupt and large swings in the pH of body fluids- by releasing H+ ions when pH is too high and by binding to H+ when pH is too low. Both counteract the deviation.
A buffer is a combination of the acid and the conjugate base
a strong acid-
proton donator that fully dissociates in solution
a strong base -
proton acceptor that fully dissociates in solution
a weak acid-
proton donator that slightly dissociates in solution- it is able to manipulate it so it dissociates fully by increasing pH or not at all by decreasing pH
a weak base-
proton acceptor that slightly dissociates - it is able to manipulate it so it dissociates fully by increasing pH or not at all by decreasing pH
Ionisation of water
the idea that water is a very very weak acid- hence will dissociate slightly.
Pure water is 55.6 M
[H+]{OH] = 10^-14
Why does urine have a greater range of pH than bloody (5-8)
Kidney regulates the blood pH therefore what the urine composition is like reflects what the kidney has done
The pKa of ethanoic acid is 4.75, why isn’t this the ideal to be used as a buffer
As blood pH is 7.4 (approx) hence would mean that using ethanoic acid would result in the blood pH falling below the lower range of pH.
List the dissociations of phosphoric acid
H3PO4—> H2PO4 -
H2PO4- —-> HPO4 2-
HPO4 2- ——> PO4 3-
Which 2 acids would be most commonly used in the body
H2PO4 - and HPO4 2-
What is the definition of pKa
pKa= -log [Ka]
Ka = [H+][OH-] / [HA]
At the half neutralisation point pH= pKa
The Lower the pKa the higher the Ka hence the stronger the acid
What is Henderson- Hasselbalch equation?
pH= pKa + log [A-] / [HA]
What are physiologically important buffers?
H2CO3 —> HCO3-
H2 PO4 2- —-> HPO4 -
Protein —-> Protein -
Protein + —–> Protein
What determines whether a protein would be a good buffer?
The R groups of the individual amino acids?
eg. Histidine
Describe the example of carbonic acid
H2CO3 ---> HCO3- pKa=6.1 [H2CO3] is proportional to the pCO2 the HH equation helps you distinguish the cause eg. with high pH low [A-] means diabetic causes high [acid] means respiratory causes
Why is haemoglobin useful in terms of acid and base
it is a good buffer due to large number histidine residues
However pKa of histidine in Hb is different to that of free His due to the neighbouring residues affecting the pKa
