Acids, Bases & Buffers PDF

Acids, Bases & Buffers I. Chemical Equilibria A. Equilibrium exists in a system, of at least 2 states, when the populations of the two states are constant 1. The members of the system may be changes from one state to another...

Acids, Bases & Buffers I. Chemical Equilibria A. Equilibrium exists in a system, of at least 2 states, when the populations of the two states are constant 1. The members of the system may be changes from one state to another 2. The concentration of the members of the system remains constant in each state B. Most chemical reactions are reversible in which starting materials (reactants) combine to give products, which can breakdown back into reactants Reactants Products Double arrows indicate the process is reversible and equilibrating II. Le Chȃtelier’s Principle A. Systems will always strive to reestablish the state or equilibrium B. Changing concentrations of reactants or products will shift the equilibrium 1. Addition of reactants will shift the equilibrium toward the products 2. Addition of products will shift the equilibrium toward the reactants 3. Example: in the alveoli, oxygen tension is high and this favors movement of oxygen onto the hemoglobin molecule; the lower oxygen concentration in the tissues favors the movement of oxygen off of the hemoglobin molecule C. Changing volume and pressure will also alter the state of equilibrium 1. Example: increasing the FiO2 will increase the dissolved amount of oxygen III. Equilibrium Constant A. The equilibrium constant (K) is a numerical description of the balance of reactants and products in molar concentrations B. In the equilibrium established below: aA + bB cC + dD [D[d[C]c K = [A]a{B]b  C. The constant (K) is often appended with a subscript denoting the type of equilibrium 1. Examble Ka is the constant governing the ionization of weak acids D. As K increases, the equilibrium established favors the products 1. If K > 1 the reaction is product-favored (more product than reactant present) IV. Acids and Bases A. An acid is a species that increases the H+ concentration 1. In an aqueous solution, H+ will exist as the hydronium ion (H3O+) B. A base is a species that increases the hydroxide ion (OH-) in an aqueous solution C. Brônsted definition: 1. Acid donates hydrogen ion to a base 2. Base accepts hydrogen ion from an acid D. In the reaction: HCL H+ + Cl- Here the H+ is the acid and Cl- is known as the conjugate base 1. Conjugate base can accept a hydrogen ion a. Conjugate bases do not have to contain a hydroxide anion E. Amphiprotic species can behave as either an acid or a base: HCO3- + H+ H2CO3 HCO3- H + CO32- + Here bicarbonate anion can act either as an acid or a base 1. The above is also an example of a polyprotic acid (H2CO3) a. Polyprotic acids are able to give more than one hydrogen ion i. In polyprotic acids, the first proton is much more easily released than successive protons F. Weak acids 1. Weak acids can donate hydrogen ions to bases, but are less determined to do so 2. When a weak acid dissolves in water it establishes a dynamic equilibrium 3. The equilibrium constant (Ka) is used to express the equilibration a. Large Ka values reflect stronger acids b. Example: [H+] [HCO3-] Ka = [H2CO3] c. In the example above, the Ka in plasma of carbonic acid is 7.94x10-7 d. As with pH, the pKa is –log(Ka) i. In the example above the pKa is 6.1 V. Buffers A. A pH buffer is a solution that resists changes in pH 1. Buffer solution contain a weak acid (HA) and its conjugate base (A-) B. If a strong acid is added to a buffered solution, the weak base in the buffer, A-, will react with the acid to give HA A- + H+ HA 1. This results in converting the strong acid into the weak acid (HA) C. If a strong base is added to a buffered solution, the weak acid, HA, will react with the base to give water and the weak base A- HA + OH- H2O + A- 1. This results in converting the strong base into the weak base A- D. Buffer solutions can exist with a wide range of pH, depending by the strength of the weak acid HA E. The pH of a buffer solution can be determined by the Henderson-Hasselbalch equation: F. Using the Henderson-Hasselbalch equation for the most effective buffer system in our plasma, bicarbonate pH = 6.1 + log [24 mEq/L / 1.2 mEq/L] 1. Here, 6.1 is the pKa of carbonic acid (H2CO3), plasma bicarbonate is 24 mEq/L and dissolved CO2 is 1.2 mEq/L 2. Solving the equation above results in a plasma pH of 7.4