+ Acids and Bases – K b 2.2K: Define Ka, Kb and use these to determine pH, pOH, [H 3 O + ] and [OH – ] of acidic and basic solutions 2.3K: Calculate equilibrium constants and concentrations for Bronsted-Lowry acid base systems when: concentrations at equilibrium are known initial concentrations and one equilibrium concentration are known the equilibrium constant and one equilibrium concentration are known.

+ Base Strength & Ionization Constant, K b Ionic hydroxides, such as NaOH (aq) or Ca(OH) 2(aq), are assumed to dissociate completely upon dissolving. Finding the hydroxide ion concentration does not involve any reaction with water Example: Find the hydroxide ion concentration of a mol/L solution of barium hydroxide

+ Base Strength & Ionization Constant, K b We can communicate the strength of weak bases, with the equilibrium constant for its reaction with water. The equilibrium constant is called the base ionization constant, K b Two common K b calculations: Calculating K b from empirical data (Example #1) Using Kb to predict the concentration of hydroxide ions when the initial concentration is known (Example #2)

+ K b Calculations Example #1: A student measures the pH of a mol/L solution of aqueous ammonia and finds it to Calculate the Kb for ammonia We will use the same method as Ka calculations, but there is usually one extra step because pH values need to be converted to find hydroxide ion concentrations 14 = pH + pOH pOH = = = OH - (aq) K b for ammonia is 1.8 x Remember K b has only 2 sig digs

+ Practice Pg. 746 #10-13 – Check answers in back of book for 11-13

+ Calculating OH - from K b First problem, the data booklet has K a values not K b values. What do we do? K w = K a K b So K b = K w /K a remember K w = 1.00 x Example #1: Solid sodium benzoate forms a basic solution. Determine the Kb for the weak base present.

+ Calculating OH - from K b Example #2: Find the hydroxide ion amount concentration, pOH, pH and the percent reaction (ionization) of a 1.20 mol/L solution of baking soda. Baking soda = NaHCO 3(s)  Na + (aq) + HCO 3 - (aq) For HCO 3 - (aq), the conjugate acid is H 2 CO 3(aq) whose K a is = 4.5 x Approximation Rule: 1.20 = > x So (1.20-x) = x = x 2. x = 1.6 x = OH - (aq) 1.20

+ Calculating OH - from K b Example #2: Find the hydroxide ion amount concentration, pOH, pH and the percent reaction (ionization) of a 1.20 mol/L solution of baking soda. 2.2 x = x 2. x = 1.6 x = OH - (aq) 1.20

+ Amphoteric Species If an entity can react as either a Bronsted-Lowry acid or base, how do you know which will be the predominant reaction? Find the Ka value, calculate the Kb value, which ever is larger wins! Example: Which reaction predominates when NaHSO 3(s) is dissolved in water to produce HSO 3 - (aq) solution? Will the solution be acidic or basic? Ka = 6.3 x Kb = Kw= 1.0 x = 7.1 x Ka 1.4 x The K a value far exceeds the K b value, so an aqueous solution of this substance will be acidic because the hydrogen sulfite ion will react predominately as a Bronsted-Lowry acid.

+ Practice Pg. 750 #1-6