Acid-base Equilibria n K a and K b n % dissociation of weak acid/bases n pH of weak acid/base solutions n pH of salt solutions n Buffers n pH of buffer solutions n Titrations:selecting indicators

How do you calculate pH of weak acids/bases n From % dissociation n From K a or K b n What is % dissociation n Amount dissociated n % Dissoc. = x 100 n Initial amount

How do you calculate % dissociation from K a or K b n 1.00 M solution of HCN; K a = 4.9 x n What is the % dissociation for the acid?

n 1.00 M solution of HCN; K a = 4.9 x n First write the dissociation equilibrium equation: n HCN(aq) + H 2 O(l) H 3 + O(aq) + CN - (aq) n [HCN] [H + ] [CN - ] n Ini. Con M 0.0 M 0.00 M n Cha. Con -x x x n Eq. Con x x x n [H 3 + O ][CN - ] x 2 n K a = = n [HCN] x

n x ~ 1.00 since x is small n x 2 n K a = ; K a =4.9 x = x 2 n 1.0 n x = sqrt 4.9 x = 2.21 x n Amount disso x n x 100 = x 100 n Ini. amount 1.00 n % Diss.=2.21 x x 100 = %

% Dissociation gives x (amount dissociated) need for pH calculation n Amount dissociated n % Dissoc. = x 100 n Initial amount/con. n x n % Dissoc. = x 100 n concentration n

n 1 M HF, 2.7% dissociated n Notice the conversion of % dissociation to a fraction: 2.7/100=0.027) x=0.027 Calculate the pH of a weak acid from % dissociation

n HF(aq) + H 2 O(l) H 3 + O(aq) + F - (aq) n [H + ][F - ] n K a = n [HF] n [HF] [H + ] [F - ] n Ini. Con M 0.0 M0.00 M n Eq.Con pH = -log [H + ] n pH = -log(0.027) pH = 1.57

pH from K a or K b n 1.00 M solution of HCN; K a = 4.9 x n First write the dissociation equilibrium equation:HCN(aq) + H 2 O(l) n H 3 + O(aq) + CN - (aq) n [HCN] [H + ] [CN - ] n Ini. Con M 0.0 M 0.00 M n Chg. Con -x x x n Eq. Con x x x

n [H 3 + O ][CN - ] x 2 n K a = = n [HCN] x n x ~ 1.00 since x is small n x 2 n K a = ; K a =4.9 x = x 2 n 1.0 n x = sqrt 4.9 x = 2.21 x n pH = -log [H + ] n pH = -log(2.21 x ) pH = 4.65

What salt solutions would be acidic, basic and neutral? 1)strong acid + strong base = neutral 2)weak acid + strong base = basic 3)strong acid + weak base = acidic 4)weak acid + weak base = neutral, basic or an acidic solution depending on the relative strengths of the acid and the base.

What pH? Neutral, basic or acidic? a)NaCl neutral b) NaC 2 H 3 O 2 basic c) NaHSO 4 acidic d) NH 4 Cl acidic

How do you calculate pH of a salt solution? n Find out the pH, acidic or basic? n If acidic it should be a salt of weak base n If basic it should be a salt of weak acid n if acidic calculate K a from K a = K w /K b n ifbasic calculate K b from K b = K w /K a n Do a calculation similar to pH of a weak acid or base

What is the pH of 0.5 M NH 4 Cl salt solution? (NH 3 ; K b = 1.8 x ) n Find out the pH, acidic n if acidic calculate K a from K a = K w /K b n K a = K w /K b = 1 x /1.8 x ) n K a = X n Do a calculation similar to pH of a weak acid

NH H 2 O H 3 + O + NH 3 [NH 4 + ] [H 3 + O ] [NH 3 ] Ini. Con. 0.5 M 0.0 M0.00 M Eq. Con x x x [H 3 + O ] [NH 3 ] K a (NH 4 + ) = = [NH 4 + ] x ; appro.:0.5 - x. 0.5 (0.5 - x)

x 2 K a (NH 4 + ) = = 5.56 x x 2 = 5.56 x x 0.5 = 2.78 x x= sqrt 2.78 x = 1.66 x [H + ] = x = 1.66 x M pH = -log [H + ] = - log 1.66 x pH = 4.77 pH of 0.5 M NH 4 Cl solution is 4.77 (acidic )

Selection of an indicator for a titration n a) strong acid/strong base n b) weak acid/strong base n c) strong acid/weak base n d) weak acid/weak base n Calculate the pH of the solution at he equivalence point or end point

What is the pH? n 50 mL of 0.1 M HCl and 50 mL 0.1M NaOH solution n 25 mL 0.5 M HC 2 H 3 O 2 and 25 mL 0.5 M NaOH n 50 mL of 1.0 M HCl and 50 mL of 1.0 NH 4 OH (NH 3 ) solution.

What is an Indicator? Indicator is an weak acid with different K a, colors to the acid and its conjugate base. E.g. phenolphthalein n HIn H + + In - n colorless pink n Acidic colorless n Basic pink

pH range of Indicators litmus ( ) bromothymole blue ( ) methyl red ( ) thymol blue ( ) phenolphthalein ( ) thymolphthalein ( )

Common Ion Effect n Weak acid and salt solutions n E.g. HC 2 H 3 O 2 and NaC 2 H 3 O 2 n Weak base and salt solutions n E.g. NH 3 and NH 4 Cl. n H 2 O + C 2 H 3 O 2 - OH - + HC 2 H 3 O 2 (common ion) n H 2 O + NH 4 + H 3 + O + NH 3 n (common ion)

Henderson Hesselbach Equation n [ACID] n pH = pK a - log n [BASE] n n [BASE] n pH = pK a + log n [ACID]

Types of Acids and Bases Binary acids Oxyacid Organic acids Acidic oxides Basic oxides Amine Polyprotic acids

Binary Acids Compounds containing acidic protons bonded to a more electronegative atom. e.g. HF, HCl, HBr, HI, H 2 S The acidity of the haloacid (HX; X = Cl, Br, I, F) Series increase in the following order: HF < HCl < HBr < HI

Oxyacids Compounds containing acidic - OH groups in the molecule. Acidity of H 2 SO 4 is greater than H 2 SO 3 because of the extra O (oxygens) The order of acidity of oxyacids from the a halogen (Cl, Br, or I) shows a similar trend. HClO 4 > HClO 3 > HClO 2 >HClO perchloric chloric chlorus hyphochlorus

Acidic Oxides These are usually oxides of non- metallic elements such as P, S and N. E.g. NO 2, SO 2, SO 3, CO 2 They produce oxyacids when dissolved in water

Basic Oxides Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water. e.g. CaO + H 2 O --> Ca(OH) 2

Protic Acids Monoprotic Acids: The form protic refers to acidity or protons. Monoprotic acids have only one acidic proton. e.g. HCl. Polyprotic Acids: They have more than one acidic proton. e.g. H 2 SO 4 - diprotic acid H 3 PO 4 - triprotic acid.

Amines Class of organic bases derived from ammonia NH 3 by replacing hydrogen by organic groups. They are defined as bases similar to NH 3 by Bronsted or Lewis acid/base definitions.