I. pH= -log [H+] on a scale 0-14. Lower pH means more acidic or more hydrogen ions; high pH means more basic (alkaline) or less hydrogen ions or more hydroxide ions; If [H+] has a molar concentration of 10-7 M, the pH is 7.
A. Acids are proton donors that release [H+] in water; pH 0-7
B. Bases are proton acceptors or release [OH–] in water; pH 7-14
II. H2 O <–> H+ + OH–
This equation is important because about two thirds of our bodies are water. At equilibrium, this reaction is pushed to the left and the concentrations of H+ and OH– are very small. Additionally, because of the covalent bond between oxygen and hydrogen, water is usually electrically neutral.
III. Electrical Neutrality (water is electrically neutral)
At equilibrium, the sum of the positive charges always equals the sum of the negative charges
IV. Strong ions– any ion that tends to be free (ie it readily dissociates to become an ion) and doesn’t readily combine with other ions. (Examples of strong ions are Na+, K+, Mg2+, Cl–, lactate–)
A. NaCl is a strong electrolyte
an electrolyte is a salt that ionizes in water and becomes a strong ion.
NaCl –> Na+ + Cl–
B. SID – Strong Ion Difference:The sum of all positive strong ions minus the sum of all negative strong ions.
1. Negative SID is an acid
2. Positive SID is alkaline. Human blood has a pH of 7.4 with a positive SID.
-Adding CO2 to a solution with a positive [SID] has a dramatic effect on [H + ]; Remember the carbonic acid reaction. CO2 can push the equation (CO2+H2O….) to the right and thus H+ are released which can counteract the positive SID.
V. Regulating body [H+]
A. Regulating PCO2
B. Regulating SID –renal system can control [Cl–] and other ions as well
C. Slight disturbances in pH can seriously disrupt physiological functions.
1. Metabolic acidosis and alkalosis
2. Respiratory acidosis and alkalosis
V. Buffers– resist changes in pH by converting strong acids or bases to weak ones
A. Physiologic Buffers –Process by which body systems that control the body’s output of acids, bases, or CO2 and therefore stabilizes pH. These systems resist changes in pH when acid or base is added.
1. Respiratory System
H+ + HCO3– <–> H2CO3 <–> CO2 (expired) + H2 O
a. A drop in pH stimulates pulmonary ventilation, which expels excess CO2 and therefore reduces H+ concentration
b. An increase in pH reduces pulmonary ventilation, which allows metabolic CO2 to accumulate in ECF faster than it is expelled and lowers pH to normal.
2. Urinary/Excretory System –this system can buffer more acid/base than the respiratory system but is slower to work. Hydrogen ions can be directly excreted into the kidney tubules.
B. Buffer Systems = Chemical Buffers –A substance that bind hydrogen ions and removes it from a solution as its concentration begins to rise, or releases hydrogen ions into a solution as its concentration falls. They work quickly.
1. Bicarbonate Buffer System This is an important buffer system because it can be regulated by the respiratory system to adjust carbon dioxide levels and the urinary system to alter hydrogen and bicarbonate levels.
CO2 + H2 O<–> H2CO3 <–> H+ + HCO3–
a. Carbonic acid is a weak acid. Enzyme for this reaction is called carbonic anhydrase and works ideally at a pH of 6.1
b. If reaction goes to the right, carbonic acid releases hydrogen ions and therefore lowers the pH. Additionally, the kidneys can excrete bicarbonate ions to keep this reaction moving to the right and increasing hydrogen ion concentrations. For example, if a base is added to the bloodstream, you have less hydrogen ions and this causes the equation to push to the right. This release more hydrogen ions to neutralize or buffer the base.
c. If reaction goes to the left, bicarbonate acts as a weak base by binding hydrogen ions, removing them from the solution, and raising pH.Additionally, the lungs and kidneys could continuously remove CO2 to keep this reaction going to the left so that more hydrogen ions are neutralized. For example, if you add an acid to the bloodstream, you have more hydrogen ions and this pushes this equation to the left. As a result you will decrease your hydrogen levels and buffer the acidic change.
2. Phosphate Buffer System This system is most important inside of our cells.
H2PO4– <–> HPO42- + H+
a. If reaction goes to the right, more hydrogen is liberated and the pH drops, thus counteracting an alkaline condition.
b. If reaction goes to the left, hydrogen ions are bound to HPO42-, and the pH increases to counteract acid conditions.
c. Enzyme for this reaction works ideally at a pH of 6.8
3. Protein Buffer System –proteins are more concentrated in the body than bicarbonate or phosphates, and therefore accounts for the majority of all chemical buffering in the body; these systems relate to amino acid carboxyl groups and amino groups.
—COOH –> –COO– + H+ (here carboxyl group is acting like an acid, as it is a proton donor, and can therefore be an effective buffer if base is added to the bloodstream)
–NH2 + H+ –> –NH3+ (here amine group is acting like a base, as a proton acceptor, and therefore can be an effective buffer if acid is added to the bloodstream. As NH3+ levels increase, the NH3+ can be excreted by the kidneys)