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1. Chemistry Grade 12 Based on the Nelson Chemistry 12 textbook

2. Product Constant for Water, Kw • We can use the product constant for water, Kw­, to calculate the hydrogen ion concentration or the hydroxide ion concentration in an aqueous solution of a strong or weak acid or base at SATP, if the other concentration is known. • Since Kw = [H+(aq)][OH-(aq)] • Then Kw = [H+(aq)] [OH-(aq)] • And Kw = [OH-(aq)] [H+(aq)] • In neutral solutions: [H+(aq)] = [OH-(aq)] • In acidic solutions: [H+(aq)] > [OH-(aq)] • In basic solutions: [H+(aq)] < [OH-(aq)]

3. Strong Acid • For strong acids we can use the concepts that strong acids ionize quantitatively in solution and use the value of Kw to calculate [H+(aq)]and [OH-(aq)] in acidic solutions. EXAMPLE • A 0.20 mol/L solution of hydrobromic acid, HBr(aq), at SATP is found to have a hydrogen ion concentration of 0.20 mol/L. Calculate the [OH-(aq)­].

4. 0.20 mol/L • HBr (aq) H+(aq) + Br -(aq) 0.20 mol/L 0.20 mol/L • H2O(l) H+(aq) + OH-(aq) • We can ignore the contribution of H+(aq) from autoionization • The OH-(aq) can also be ignore. • Therefore the major entities in the solution are: (which are the major entities affecting acid-base characteristics) • H+(aq) and Br -(aq) • but the Br -(aq) is the conjugate base of a strong acid and therefore can be ignored as weak

5. So since HBr(aq) ionizes quantitatively (100%), [H+(aq)] = 0.20 mol/L • Now we can use the Kw expression to calculate the concentration of OH-(aq). • [OH-(aq)] = __Kw____ [H+(aq)] • [OH-(aq)] = 1.0 x 10-14/0.20 = 5.0 x 10-14 mol/L

6. Strong Base • Just as with acids, we can use the two concepts for solutions of strong bases – that they dissociate quantitatively in solution and the value of Kw to calculate the hydrogen ion concentration or hydroxide ion concentration.

7. EXAMPLE • Calculate hydrogen ion concentration in a 0.20 mol/L solution of magnesium hydroxide, a strong base. • Mg(OH)2(aq) Mg2+(aq) + 2OH-(aq) 0.20 mol/L 0.20 mol/L 2(0.20 mol/L) = 0.40 mol/L • Major entities: Mg2+(aq), 2OH-(aq), H2O(l)

8. pH – measure of the hydrogen ion concentration • pH = -log[H+(aq)] • if [H+(aq)] = 2.5 x 10-14 • pH = -log [H+(aq)] = -log (2.5 x 10 -14) pH = 13.60 • In water and any neutral solution, the pH is 7.00 (show calculations, as on p. 541).

9. To build on the previous chart: • In neutral solutions: [H+(aq)] = [OH-(aq)] pH = 7.00 • In acidic solutions: [H+(aq)] > [OH-(aq)] pH < 7.00 • In basic solutions: [H+(aq)] < [OH-(aq)] pH > 7.00

10. pOH – measure of the hydroxide ion concentration • pOH = -log[OH-(aq)] • [OH-(aq)] = 10-pOH • pH + pOH = pKw = 14.00 EXAMPLE • What is the pH of a solution whose pOH is 2.3? • [OH-(aq)] = 10-pOH = 10-2.3 • [OH-(aq)] = 5.0 x 10-3

11. The pH of Strong Acids and Bases • The pH of Strong Acids • The pH of solutions of strong monoprotic acids is calculated from the concentration of H+(aq) ions, which is assumed to be the molar concentration of the solute molecule before ionization. • The pH of Strong Bases • As with strong acids, the pOH and the pH of strong bases are determined entirely by the OH-(aq) ion contributed by the dissociation of one of the ionic hydroxide solution.

12. Weak Acids • weak electrolyte • does not ionize completely in water to form hydrogen ions • most common acids are weak acids: HF(aq),, H2CO3(aq), H2S(aq), H3BO3(aq) • A weak acid is an acid that partially ionizes in solution but exists primarily in the form of molecules.

13. Weak Bases • Arrhenius theory of bases states that bases are soluble ionic hydroxides that dissociate in water into positive metal ions and negative hydroxide ions. • There are some molecular and ionic compounds, other than hydroxides, which also dissolve in water to produce basic solutions that are called weak bases as they are not as basic as ionic hydroxide solutions of the same concentration.

14. The Brønsted-Lowry definition of a weak base is a compound that reacts non-quantitatively (incompletely) with water to form an equilibrium that includes hydroxide ions according to the following general equation: B(aq) + H2O(l)  OH-(aq) + HB+(aq)

15. Percent Ionization of Weak Acids • most weak acids ionize less than 50%, unlike strong acids that ionize close to 100% • Percent Ionization (p) is defined as follows: • p = conc. of acid ionized x 100% conc. of acid solute • For a general acid ionization reaction: HA(aq) H+(aq) + A-(aq) • p = [H+(aq)] x 100% [HA(aq)] • [H+(aq)] = p x [HA(aq)] 100 • If we know the pH of a weak solution, we can calculate the percent ionization of the acids.

16. Ionization Constants for Weak Acids • We can consider equilibrium solutions of a weak acid dissolved in water to be just like the equilibrium systems we looked at in Chapter 7 • we can use the equilibrium law expression and calculate equilibrium constants • acid ionization constant: Ka • Ka is usually determined experimentally • Percent ionization can be used to calculate the Ka­ value (using an ICE table)

17. EXAMPLE • Calculate the acid ionization constant, Ka, of acetic acid if 0.1000 mol/L solution at equilibrium at SATP has a percent ionization of 1.3%. • HC2H3O2 (aq) H+ (aq) + C2H3O2-(aq) • Ka = [H+(aq)][C2H3O2-(aq)] [HC2H3O2 (aq)] • Create ice table and substitute the values back into the Ka expression to solve. Ka = 1.7 x 10-5

18. Ionization Constants for Weak Bases • Weak bases from dynamic equillibria in aqueous solutions • The reaction of weak bases with water may be defined by the equilibria law) • OH-(aq) ions are produced and affect the acid-base characteristics of solutions • Calculated the same way as weak acids

19. Relationship between Ka and Kb • Kw = KaKb • This equation allows us to convert the Ka values of acids into the Kb values of their conjugate bases, and vice versa, given the value of Kw, which is a constant.

20. Polyprotic Acids • Can anyone remember acids with more than one ionizable proton (more than one proton to give away)? • Exs: sulfuric acid, phosphoric acid, and boric acid • Donate one proton at a time in a stepwise fashion • Each ionization reaction has its own acid ionization constant, Ka1, Ka2 … • The acid in each step is weaker, generally: Ka1 > Ka2 > Ka3

21. Acid-Base Properties of Salt Solutions

22. Salts • Solids • Composed of cations and anions • Dissolve in water • May or may not alter the pH of a solution, due to the cation, anion, or both

23. Salts that form Neutral Solutions • Salts of cations from strong bases and anions of strong acids have no effect of the pH of an aqueous solutionExs: NaCl(aq), KCl(aq), NaI(aq), NaNO3(aq)

24. Salts that form acidic solutions • The salt of a weak base (cation) and strong acid (anion) dissolves in water to form acidic solutions. • The cation reacts with water to liberate H+ • The solution has a pH less than 7 NH4Cl  dissociates  NH4+ & Cl- NH4+ + H2O  H3O+(aq) + NH3(aq)

25. Salts that form basic solutions • The salt of a strong base (cation) and weak acid (anion) dissolves in water to form basic solutions. • The anion reacts with water to liberate OH- • The solution has a pH greater than 7 NaC2H3O2dissociates Na+ & C2H3O2- C2H3O2-(aq)+ H2O  HC2H3O2(aq) + OH-(aq)

26. Salts that act as acids and bases Some salts: • contain both the cation of a weak base and the anion of a weak acid • both ions can hydrolyze • we can predict whether the solution is acidic, basic or neutral: Ka > Kb then the solution is acidic Kb > Ka then the solution is basic

27. Acid-Base Titration

28. Start by matching titration terms

29. Totally Terrific Titration Terms • Titration: a chemical analysis involving the progressive addition of a solution of known solute concentration into a solution of unknown concentration to determine the amount of a specified chemical • Titrant: solution of known concentration, usually found in a buret during titration

30. Terms • sample: the solution of unknown concentration being analyzed in a titration • primary standard: chemical that is available in a pure and stable form that can be used to produce an accurate concentration • standard solution: a stock solution that is of known concentration that is used to create the titrant

31. Terms • endpoint: point in titration when the pH indicator changes colour • equivalence point: point in titration when chemically equivalent amounts of reactants have reacted (point at which equal amounts of H3O+(aq) and OH-(aq) have been added); generally, an equilibrium is established at this point

32. Terms • Standardization: a titration used to find the concentration of the titrant using a primary standard • indicator: an acid-base indicator that will change colour at a known pH to signify to signify a specific pH in the neutralization reaction

33. Titration of a strong acid with a strong base • At the equivalence point [H+]=[OH-] or pH=7. • Phenolphthalein is a popular indicator because it is colourless in acidic solutions and pink in basic. • Remember to consider the reaction at the molar level. (convert to moles!) • C=n/V and C1V1 = C2V2

34. Titration of a strong acid with a strong base • Titration curve – a plot of the pH vs. Volume of titrant added. pH Vol. of titrant

35. Titration Websites • Titrations • http://www.chemguide.co.uk/physical/acidbaseeqia/phcurves.html (curves & notes) • http://www.chem.uoa.gr/applets/AppletTitration/Appl_Titration2.html (show different graphs) • http://www.vias.org/simulations/simusoft_titration.html (simulation)

36. Buffers • Website: • http://www.chemguide.co.uk/physical/acidbaseeqia/buffers.html#top