ACID-BASE EQUILIBRIA “Acid-Base Equilibria” is very important not only in Chemistry but also in related sciences such as biology, medicine, agriculture and geology. Topics such as production of chlorophyll in plants, formation of limestone in caves, deterioration of a marble statue in polluted air, and the activity of a lifesustaining enzyme can only be discussed intelligently if we understand acidbase behaviour. Recalling the definitions of Acids and Bases Early & far from ideal definitions: Acids – substances that taste sour; turn litmus red Bases – substances that taste bitter; turn litmus blue. Note that not all compounds exhibiting these properties are chemically similar. Arrhenius Theory of Ionization (1880s) Acids – substances that dissociate into hydrogen ions and anions when dissolved in water. Bases – substances that dissociate into hydroxide ions and cations when dissolved in water. Bases react with acids to form a salt and water. The water produced in an acidbase reaction was considered to be the result of hydrogen ions combining with hydroxide ions. Bronsted-Lowry Concept (1923; England) Acid - any substance capable of donating a proton in a reaction. Base – any substance capable of accepting a proton in a reaction. The B-L concept of acid-base (A-B) behavior has proven especially convenient when using water and water-like (protophilic) solvents. Role of the Solvent One of the most significant features of the B-L concept of the A-B behavior is the incorporation of the solvent as a vital part of the A-B system. For example, when formic acid dissolves in water, ionization occurs because the solvent acts as a base or proton acceptor. HCOOH + H2O <==> H3O+ + HCOO- acid1 base2 acid2 Base1 (Reaction 1) The H3O+ is merely a solvated proton called the hydronium ion. When ammonia is dissolved in water, the solvent act as an acid or proton donor. NH3 base + 1 H2O <==> NH4+ + OH- acid acid base 2 1 (Reaction 2) 2 Amphiprotic solvents – solvents like water that can exhibit both acidic and basic character, depending on the solute. Low MW alcohols and acetic acid are other common amphiprotic solvents. Exercise. Do this in five minutes in a one-half sheet of paper, take a photo & send in my messenger account. 1) HCOOH + C2H5OH <==> 2) NH3 + C2H5OH <==> Conjugate acids and bases When an acid donates a proton in a reaction, it becomes a substance that is capable of accepting a proton to re-form the original acid. Such a substance is A B-L base. When a base accepts a proton in a reaction, it becomes a substance capable of donating a proton. Bronsted & Lowry referred to such substances as conjugate pairs. Every B-L acid has a conjugate base and every B-L base has a conjugate acid. IONIZATION OF WATER H2O + H2O <==> H3O+ + OH- acid1 base2 base1 acid2 (Reaction 3) acid1 & base1 are conjugates; acid2 & base2 are conjugates. All amphiprotic solvents undergo self-ionization or autoprotolysis reactions. The extent to which the reaction occurs is represented by the expression Keq = [(aH3O+)(aOH-)] / (aH2O) = [(aH3O+)(aOH-)] / 1 Keq = Kw = [H3O+][ OH-] (Equation 4) or, taking the negative logarithms of both side pKw = pH + pOH (Equation 5) autoprotolysis constant for water, Kw = 1.00 x 10-14 at 24oC (about room temperature) According to Reaction 3, the concentrations of hydronium and hydroxide ions in pure water must be equal and Equation 4 is used to calculate hydronium ion concentration of pure water: Kw = [H3O+]2 [H3O+]2 = 1.00 x 10-14 [H3O+] = 1.00 x 10-7 M pH = -log (1.00 x 10-7) = 7.000 & pOH of pure water = 7.000 According to Equation 4, an increase in the hydronium-ion concentration resulting from the addition of an acid to water must be accompanied by an equal decrease in the hydroxideion concentration. Conversely, an increase in the hydroxide-ion concentration must be accompanied by a decrease in the hydronium ion concentration. If the concentration of either H3O+ or OH- is known, the other can be calculated using Eq. 4. Exercise (Acid-Base Equilibria) 1. What is the concentration of H3O+ in a 1.00 x 10-2 M NaOH solution? 2. Calculate the pOH of a 0.020 M HCl solution.