Review 3

1. Review 3 Chapter 15,16 Acids, Bases, Salts, Buffer and Insolubles

2. Chapter 15: Acids and Bases • Concepts: H+ == H3O+ • Arrhenius: Limitation --- water as solvent • Acids are H+ generators • Bases are OH- generators • Bronsted-Lowry: Conjugate acid-base pairs • Acids are H+ Donors • Bases are H+ Acceptors • An acid upon losing H+ forms Conjugate base • Base upon gaining H+ forms Conjugate Acid • H2PO4- its conjugate Acid is H3PO4 • And H2PO4- its conjugate Base is HPO4-2

3. Higher oxidation number = MORE OXYGEN on the central atom= stronger oxyacid H2SO4 > H2SO3; HNO3 > HNO2 • More Electronegative the central atom in the oxy acid: Stronger is the acid • HClO3 > HBrO3 > HIO3 • Non Oxy acids; Weaker HA bond stronger is the acid HI >HBr > HCl > HF • Cation stronger acid than neutral molecule; neutral stronger acid than anion • H3O+ > H2O > OH−; NH4+ > NH3 > NH2− • Trend in base strength opposite • Stronger the Acid ---weaker is its conjugate base and vice versa • Lewis: Includes complex ion formation • Acids are electron pair acceptors (e.g. BF3, electron deficient molecule) • Bases are electron pair Donors (e.g. Ammonia- has lone pair)

4. Acid strength and pH = -log (H+ ); Scale 0-14, 0 to <7 acid; 7 neutral and above 7 basic/alkaline Kw = 10-14 at 25 ºC and water has [H+ ] = [OH-] = 10-7M ; pH = 7 neutral Remember Strong Acids (HCl, HBr, HI, HNO3, H2SO4, HClO4) and Strong Bases (group 1A and IIA hydroxides) • pH of Strong Acids –All Monoprotic (except H2SO4) – 100% dissociation, therefore [Acid] = [H+] –Direct calculation of pH • Ka = ∞ for strong acids • Strong bases IA –OH : [NaOH] = [OH-] – Direct pOH pH = 14-pOH IIA –OH: [OH-] = 2 * [Mg(OH) 2] – Direct pOH pH = 14-pOH • pH of WA (monoprotic) and HA +H2O ↔ A- + H3O+ • Ka = {[A- ] * [H3O+ ]}/[HA] ; Larger Ka stronger is the acid, • Ka is called acid dissociation or ionization constant; pKa = -log (Ka) • Use Ka and ice chart to calculate [H3O+] and pH • Check for % dissociation, if less than 5%, approximations are valid

5. WB (mono basic) – Use Kb and ice chart- Ammonia and ammonia based amines are weak bases; Also many metal hydroxides • Base: + H2O HBase+ + OH− • The equilibrium constant is called the base ionization constant, Kb. • Larger Kb = stronger base • -log (Kb) = pKb • Ka. Kb for conjugates is Kw, so pKa + pKb = 14 and also pH + pOH = 14 • Polyprotic acids • The ionization constants for H2SO4 are as follows: • H2SO4 + H2O  HSO4 + H3O+ strong • HSO4 + H2O  SO42 + H3O+ Ka2 = 1.2 × 10−2 • Because the first ionization is complete, use the given • [H2SO4] = [HSO4−]initial = [H3O+]initial. • Other Polyprotic acids (EXCEPT H2SO4) • Some Observations • Ka1 >> Ka2 >> Ka3 • Generally, the difference in Ka values is great enough so that the second ionization does not happen to a large enough extent to affect the pH. • Most pH problems just do first ionization. • [A2−] = Ka2as long as the second ionization is negligible.

6. Salts • Table 15.9 for summarization (page 730) – Tro • Salts are water-soluble ionic compounds • They contain Cations of Base and Anions of the Acid • Is the salt Acidic or basic? • Example NaHCO3 solutions are basic. • Na+ is the cation of the strong base NaOH. • HCO3− is the conjugate base of the weak acid H2CO3. • NH4Cl solutions are acidic. • NH4+ is the conjugate acid of the weak base NH3. • Cl− is the anion of the strong acid HCl. • NaCl solutions are neutral. • Na+ is the cation of the strong base NaOH. And Cl− is the anion of the strong acid HCl. • NaF (aq) is Basic • Na+ is the cation of the strong base NaOH. • F− is the conjugate base of the weak acid HF. • NH4F (Ka HF 10-4, Kb NH3 = 1.8*10-5) • Ka >Kb so the salt is Acidic • NH4C2H3O2 Ka = Kb for Acetic acid and Ammonia, so Neutral pH • Calculate pH of salt solutions Use the weak part of Salt and hydrolyze • Anions of weak acid hydrolyze to produce OH- • A−(aq) + H2O(l)  HA(aq) + OH−(aq) need Kb of A- use Kw/Ka = Kb • Cations of weak base hydrolyze to produce H3O+ • NH4+(aq) + H2O(l)  NH3(aq) + H3O+(aq) need Ka of NH4+ use Kw/Kb = Ka • -Cations of small, highly charged metals are weakly acidic • For Conjugate species,

7. Chapter 16 Buffers (part) and Titration • Topics • Buffers • Titration- pH curve • Choice of indicators • Solubility and Ksp • Complex Ion formation • Qualitative Analysis

8. Buffers • Buffers are solutions that resist changes in pH when an acid or base is added • They act by neutralizing acid or base that is added to the buffered solution • Many buffers are made by mixing a solution of a weak acid with a solution of soluble salt containing its conjugate base anion So has significant concentration of HA and A- Common Ion Effect HA(aq) + H2O(l) A−(aq) + H3O+(aq) Tro: Chemistry: A Molecular Approach, 2/e

9. Review for Buffers • Base buffers vs Acid Buffers • Approximations are ALWAYS valid in buffer due to common ion effect and le Chaterliers principle • Henderson Hasselbach Equation- Simple version using Approximations – FOR Both acid and base buffer use the same equation • Calculation of pH changes when small amounts of acid or base is added Alternate version of Henderson equation for a base buffer The buffering capacityis the amount of acid or base a buffer can neutralize The buffering rangeis the pH range the buffer can be effective

10. Titration Curve • A plot of pH vs. amount of added titrant • The inflection point of the curve is the equivalence point of the titration • Prior to the equivalence point, the known solution in the flask is in excess, so the pH is closest to its pH • The pH of the equivalence point depends on the pH of the salt solution • equivalence point of neutral salt, pH = 7 • equivalence point of acidic salt, pH < 7 • equivalence point of basic salt, pH > 7 • Beyond the equivalence point, the unknown solution in the burette is in excess, so the pH approaches its pH Tro: Chemistry: A Molecular Approach, 2/e 10

11. Titration Curves

12. Titration Curves Weak acid- Strong Base Conclusion:At the equivalence point of the titration, unlike the titration of a strong acid and strong base, the pH is > 7. This is due to the production of the conjugate base of a week acid.

13. Titration Curves Weak base- Strong Acid In the case of a titration of a weak base, the process follows that of a weak acid in reverse. There exists a region of buffering followed by a rapid drop in pH at the eq. point.

14. Titration of a Weak Polyprotic Acid with a Strong Base In the case of a titration of a weak polyprotic acid (HnA) there are “n” equivalence points. In the case of the diprotic oxalic acid, (H2C2O4) there are two equivalence points.

15. Solubility Product The extent of solubility can be measured by the equilibrium process of the salt’s ion concentrations in solution, Ksp. MnXm(s)  nMm+(aq) + mXn−(aq) The solubility product would be Ksp = [Mm+]n[Xn−]m For example, the dissociation reaction for PbCl2 is PbCl2(s)  Pb2+(aq) + 2 Cl−(aq) And its equilibrium constant is Ksp = [Pb2+][Cl−]2 Solubility is the amount of solute that will dissolve in a given amount of solution at a particular temperature The molar solubility is the number of moles of solute that will dissolve in a liter of solution the molarity of the dissolved solute in a saturated solution

16. Solubility & the Common Ion Effect Adding an ion “common” to an equilibrium causes the equilibrium to shift towards reactants according to Le Chatelier’s principle.