1. Ch 15 Acids and bases

2. Acids Conduct electric current (electrolyte)  Acid + water --> H 3 O + (hydronium ion)  Reacts with metals (above hydrogen on the activity series  Acid + metal --> H 2 (g) + salt  Generate hydronium ions in water(pg 532 table 1)  Strong acid --> 100% ionization  Weak acid less than 100% ionization  Svante Arrhenius defines acids as generating H 3 O +  Polyprotic acids form more than one H 3 O + ion. 1st one ionizes more than the rest  H 2 SO4 --> H 3 O + + HSO 4 -  HSO 4 - H 3 O + + SO 4 2-

3. Bases (electrolytes) Usually solids, slippery (reacts with oils in skin => soap)  Good cleaning agent Some bases insoluble in water (Mg(OH) 2 )  Weak base, use solubility rules(pg 533 table 2) Some bases very soluble in water (NaOH)  Strong base, akaline metals with OH- Bases generate hydroxide ions (OH-)  Strong base dissociate complete  Weak base dissociate less than 100% Oxides, carbonates and phosphates are basic  Na 2 O -> 2Na + + 2OH-, CO 3 2- HCO 3 - + OH-

4. Bronsted-Lowry A/B Arrhenius definition of A/B narrow Bronsted-Lowry definition of A/B more broad  Focus on the proton (H + ) transfer between two sub. B-L acid:  Proton donor HCl + H 2 O --> H 3 O + + Cl- B-L base:  Proton acceptor NH 3 + H 2 O NH 4 + + OH- Conjugate acid/base (pg 537 table 3):  A/B reaction: one substance passes a proton(acid) to another substance(base)  Acid forms a conj. base, base forms conj. Acid

5. Amphoteric species Have both acid and basic properties Can donate and accept protons HCO 3 -(acid)+ NH 3 CO 3 2- + NH 4 + HCO 3 -(base) + HCl H 2 CO 3 + Cl-

6. Acidity, basicity & pH Self ionization of water: water is amphoteric and self ionizes  H 2 O + H 2 O H 3 O + + OH-  Pure water contains ions [H 3 O + ] = [OH-] = 1 x10 -7 Equilibrium: 2H 2 O(l) H 3 O + + OH-  Keq = Kw = [H 3 O + ][OH - ] = (1x10 -7 )(1x10 -7 )=1x10 -14 Increase[H 3 O + ](acid) decrease[OH-] increase[OH-](base) decrease[H 3 O + ] Calculate: What is [OH-] for aq 0.125M solution of HCl What is [H 3 O + ] in a solution of 0.000500M NaOH

7. Calculating pH pH and acidity: neutral solution have equal amounts, [ ], of H 3 O + and OH- ions  [ ] = 1 x10 -7 M Soren Sorensen (pg 542 table 5 range [H 3 O + ]/pH)  Used power of 10 (log scale) to represent [H 3 O + ] pH: “power of hydrogen”  Negative logarithmic scale, low ph reflect high [H 3 O + ] Calculate pH: pH = -log[H 3 O + ] Calculate: pH if [H 3 O + ] = 3.5x10 -4 M pH if [OH-] = 5.0x10 -5 M [H 3 O + ] ion that has a pH = 2.25 [H 3 O + ] & [OH-] ion that has a pH =10.65 pH 0-7 acid, pH 7 neutral, pH 7-14 basic

8. Measuring pH Indicators: quick but not very precise pH meter: more precise, yet complicated & expensive Indicators (pg 546 figure 12):  Cpd turns different colors in solution of different pH  Some have a rainbow of colors (range of pH) and some have just 2 that give above or below pH indication. pH meter:  Electronic instrument with probes  Probe: two electrodes, one sensitive to [H3O+]  ==> electric voltage develop and current measured that is convert to pH, meter has to calibrated with solutions of known pH

9. Neutralization & titrations Neutralization reaction: reaction between an acid and base to form salt and water.  H3O+ + OH- --> 2H2O(l) (net ionic reaction)  Equal amounts of each, reaction 100% Titrations:one solution is added to another solution with the use of a buret and flask  At equilibrium point equal amounts of acid and base have been added  With Sa & SB eq. pt. occurs at a pH = 7  Used to determine [ ] of a substance by adding a solution of know volume & [ ] until the reaction is complete as determine by an indicator or pH meter

10. Titration continue Titrant:solution of known [ ] used to titrate a solution of unknown [ ] Titration curve: graph showing pH changes as the titrant is added. Eq. pt. can be determined Pg 552-553 skill tool box/ performing titration pH 7 Volume titrant (mL) Eq. pt. Strong acid Titrated with a Strong base

11. Select indicator Indicators: have transition range (end pt)  pH range through which an indicator changes color due to acid or base form of indicator(pg554 figure 19 and table 6)  Eq. pt. and end pt. should be within one pH unit SA/SB eq.pt at pH = 7, SA/WB eq. pt. below pH 7, WA/SB eq. pt. above pH 7. Titration calculations:  Determine [ ] of unknown substance at eq. pt. n H3O+ =n OH- [ ] acid V acid = [ ] base V base Acid has to be mono protic!!

12. Titration calculations A volume of 20.0mL of a solution of HNO 3 that has an unknown [ ] is titrated with 34.37 mL of a 0.8220 M solution of NaOH. What is [HNO 3 ]? What volume of a 1.366M solution of NaOH would be required of NaOH would be required to titrate 47.22 mL of 2.075 M solution H 2 SO 4

13. Equilibrium of WA/WB Wa & WB: some acids are better proton donors, some bases are better proton acceptors SA form weak conj. base, SB form weak conj. acids and vise versa (pg 559 table 7). Reactions tend to favor the weaker A/B pair. Ka (equilibrium constant for weak acids)  Equilibrium constant for a reaction in which an acid donates a proton to water (pg 559 table 7)  Smaller Ka => weaker the acid. CH 3 COOH H 3 O + + CH 3 COO - Ka = [H 3 O + ] [CH 3 COO-]/[CH 3 COOH] = 1x10 -5

14. Calculations with WA What is [H 3 O + ] in a 0.250M solution of benzoic acid C 6 H 5 COOH? In a 0.025M solution of formic acid, the [H 3 O + ] = 2.03 x 10 -3 M, calculate the Ka for HCOOH?

15. Buffer solutions Made from a weak acid and its conjugate base (salt) that neutralizes small amounts of acid or base added to it. Blood pH range 7.35-7.45. Use buffer solutions to maintain that range  Acidosis or alkalosis if buffer is exceeded 2 parts of a buffer:  WA & salt of conj. base.  Acetic acid (CH 3 COOH & sodium acetate (NaCH 3 COO) How a buffers stabilizes pH:  Add acid, conj. base neutralizes H+ + CH 3 COO- -> CH 3 COOH  Add base, weak acid neutralize OH- + CH 3 COOH --> CH 3 COO- + H 2 O