pH: Acid-Base Concentration
As the concentration of hydrogen ions in a solution increase, the more acidic the solution becomes.
As the level of hydroxide ions increases the more basic, or alkaline, the solution becomes.
Clinically, the concentration of hydrogen ions in the body is measured in units called pH units, which is a way to quantify the amount of [H+] in the form of a positive number.
pH scale runs from 0 to 14 and is logarithmic
Each successive unit change in pH represents a tenfold change in hydrogen ion concentration
pH of a solution, therefore, is defined as the negative logarithm, to the base 10, of the hydrogen ion concentration [H+] in moles per liter
When the pH is 7 ([H+] = 10-7 mol/liter):
Number of hydrogen ions precisely equals the number of hydroxide ions (OH–)
And the solution is neutral - Neither acidic or basic
Pure water has a neutral pH of 7, or 10-7 mol/liter (0.0000001 mol/liter) of hydrogen ions.
A solution with a pH below 7, is acidic
There are more hydrogen ions than hydroxide ions
For example, a solution with a pH of 6 has 10 times more hydrogen ions than a solution with a pH of 7
A solution with a pH greater than 7, is alkaline
Hydroxide ions outnumber the hydrogen ions
For example, a solution with a pH of 8 has 10 times more hydroxide ions than a solution with a pH of 7
Thus, as the hydrogen ion concentration increases hydroxide ion concentration falls, and vice versa.
pH Values of Representative Substances
Figure 7-7 The pH values of representative substances. The pH scale represents the number of hydrogen ions in a substance. The concentration of hydrogen ions (H+) and the corresponding hydroxyl concentration (OH–) for each representative substance is also provided. Note that when the pH is 7.0, the amount of H+ and OH– are equal and the solution is neutral.
Acid-base balance is what keeps [H+] in normal range.
Tissue metabolism produces massive amounts of CO2, which is hydrolyzed into the volatile acid H2CO3.
CO2 + H2O → H2CO3 → H+ + HCO3–
Reaction is catalyzed in RBCs by carbonic anhydrase. As the lungs eliminate CO2, the falling CO2 reverses the reaction.
CO2 + H2O ← H2CO3 ← HCO3– + H+
Buffer solution characteristics
A solution that resists changes in pH when an acid or a base is added, in other words they minimize the effects of [H+] or [OH-] changes
Composed of a weak acid and its conjugate base
i.e., carbonic acid/bicarbonate, which in blood exists in reversible combination as NaHCO3 and H2CO3
Add strong acid HCl + NaHCO3 → NaCl + H2CO3, and it is buffered with only a small acidic pH change.
Add base NaOH + H2CO3 → NaHCO3 + H2O, and it is buffered with only a slight alkaline pH change.
Can be open or closed buffer systems.
Bicarbonate and nonbicarbonate buffer systems
Bicarbonate: composed of HCO3– and H2CO3
Open system as H2CO3 is hydrolyzed to CO2. Ventilation continuously removes CO2 preventing equilibration, driving reaction to the right
HCO3– + H+ → H2CO3 → H2O + CO2
Removes vast amounts of acid from body per day
Can buffer both strong acid and base and thus prevent major changes in pH of the solution
Able to buffer nonvolatile fixed acids, however cannot buffer volatile acids as it is in equilibrium with the volatile acid.
Has the greatest buffering capacity
Nonbicarbonate: composed of phosphate and proteins (including hemoglobin)
Consists of a weak acid and its base or salt:
HHB (acid hemoglobin) KHB(potassium hemoglobin)
KH2PO4 (potassium acid phosphate) K2HPO4 ( potassium alkaline phosphate)
H Protein (acid porteinate) NA protein (sodium proteinate)
NaH2PO4 (sodium acid phosphate) NaHPO4 (sodium alkaline phosphate)
Closed system as no gas is available to remove acid by ventilation. All components remain in the system - when equilibrium is reached no further buffering can occur
Hbuf/buf– represents acid and conjugate base.
H+ + buf– ↔ Hbuf reaches equilibrium, buffering stops
Buffers volatile acids.
Both systems are important so that the buffering of both fixed and volatile acids occurs.