1 / 60

9: Charge transfer rxns: acids and bases and oxidation-reduction

Learn about acids and bases, Arrhenius & Bronsted-Lowry definitions, conjugate pairs, strength, dissociation of water, pH scale, and more in chemistry. Explore key concepts with examples.

jskeete
Download Presentation

9: Charge transfer rxns: acids and bases and oxidation-reduction

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. 9: Charge transfer rxns: acids and bases and oxidation-reduction

  2. What are acids? • Acids: • sharp,sour taste, feel prickly on skin, • react with metals to release H2(g), • react with carbonates to release CO2(g), • turn litmus red, • neutralize bases.

  3. What are bases? • Bases: • bitter taste • slippery to touch • turn litmus blue

  4. Arrhenius definitions • Acid increases __ conc (hydrogen ion, proton)when dissolved in water. • HCl (aqCl-(aq) • Acid has to have __ to be acid.

  5. Arrhenius definitions • Base increases ____(hydroxide ion) conc when dissolved in water • NaOH  Na+(aq) + OH-(aq) • Base has to have __

  6. Arrhenius acid-base rxn • Arrhenius acid-base neutralization rxn yields salt + water. • Arrhenius limited to aqueous solution.

  7. But NH3 is a base!! • Arrhenius says bases have to have OH-’s and acids H+’s. • However NH3 is a base but has no OH- in its formula. • Now have enlarged defn of acid and base that’s not restricted to aqueous soln

  8. A Bronsted-Lowry (B-L) acid • A B-L base • The acid-base neutralization rxn is the transfer of a H+ from an acid to a base.

  9. Concept of base in enlarged (H+ acceptor) • Acid still has to have H+.

  10. Identify acids and bases in • SO32- + HCl HSO3- + Cl- • Note that the products of the rxn are acids and bases themselves.

  11. H3PO4 + H2O  H3O+ + H2PO4- • H3O+: hydronium ion • HCO3- + H2O  H2CO3 + OH-

  12. Water • Note that water can act both as a base and as an acid. • H3PO4 + H2O  H3O+ + H2PO4- • HCO3- + H2O  H2CO3 + OH- • Amphiprotic (amphoteric)

  13. Conjugate acid-base pairs • Conjugate acid-base pairs are species on opposite sides of the rxn that differ by one proton only. • Are these conjugate acid/base pairs? • H 2O/OH- • H2SO4/SO42- • H2SO4/ HSO3-

  14. Given base, to get conjugate acid: ___ • What is the conjugate acid of: HS- SO42- NH3 S2- • Given acid to get conjugate base: _________ • What is the conjugate base of: • HS- NH3 HSO4- S2-

  15. 9.36: Identify conjugate acid-base pairs in: • NH4+(aq) + CN--(aq)  NH3(aq) +HCN(aq) • CO32-(aq) + HCl(aq)  HCO3-(aq) + Cl-(aq)

  16. Strength of acids and bases • Some species are better proton donors than others; some species are better proton acceptors than others. • The better the species is at donating a proton to the base, the stronger the acidic properties of the species. • Strength is a measure of the degree of dissociation of the acid or base in solution. Stronger the acid(base) the larger the value of K for that dissociation.

  17. The better the species is at accepting the proton from the acid, the stronger the basic properties of the species. • We are talking about strong and weak here in relative terms.

  18. Generally talk about aqueous sol’ns. • Some acids and bases are strong acids and bases meaning that they are • HCl + H2O Cl- or HCl(aq) H+(aq) + Cl-(aq) • NaOH(aq)  Na+(aq) + OH-(aq) • Ca(OH)2(aq)Ca2+(aq) +2OH-(aq)

  19. Some acids and bases are not 100% ionized. These are weak acids (and bases). (weak electrolytes)This means that K < 1 for these rxns. Instead an equilbrium is set up: • HF + H2O  H3O+ + F- Ka= 7.1 x 10-4 • HCOOH + H2O  HCOO- + H3O+ Ka = 1.7 x 10-4 • Larger value of Ka stronger the acid

  20. How do you know what acids and bases are strong? • See fig 9.2 p239 Strong acids Strong bases • HCl OH- (metal ) • H2SO4 (1st ioniz. only) S2- • HNO3 O2- • H3O+

  21. The stronger the acid, the weaker the conj. base of the acid. The weaker the acid, the stronger the conj. base of the acid, etc. • What does this imply?

  22. Acid (dec ) Base (inc  • HCl Cl- • H3O+ H2O • HSO4- SO42- • HF F- • CH3COOH CH3COO- • NH4+ NH3 • Which is thestronger acid--HF or HSO4-? • Which is the stronger base--H2O or F-?

  23. Dissociation of water • HOH(l) + HOH(l)  H3O+ + OH- • Acid1 base2 acid2 base1 • Or H2O(l)  H+(aq) + OH-(aq) • Find from experiment that for pure water at 25oC that [H3O+] = [OH-] = • Can write the ion product for the dissociation of water as ion product = [H3O+][OH-] = ____________ at 25oC • H3O+ and H+ stand for the same thing.

  24. Can write the ion product for the dissociation of water as ion product = [H3O+][OH-] = 1 x 10-14 at 25oC • This is also called Kw. • This holds for all solutions not just pure HOH

  25. Kw = [H3O+][OH-] = 1 x10-14 • as [H3O+]  • If [H3O+] = 4.6 10-3M, [OH-] =? • as [OH-]  • If [OH-] = 1.22 x 10-4M, [H3O+]=?

  26. pH or how to avoid the exponents • define pH = -log10[H3O+] • Neutral sol’n: [H3O+] =[OH-] = 1 x10-7M pH=_ • Acidic sol’n; [H3O+] >OH-] ; [H3O+] > 1x10-7M pH<_ • Basic sol’n: [OH-]> [H3O+]; [OH-]>1x10-7M pH >_

  27. [H+]M pH • 10 -1 • 1 0 • 0. 1 1 • 0.001 3 • 1x10-7 7 • 1x10-10 10 • 1x10-14 14 • 1x10-15 15 • As pH h [H+] i

  28. What’s pH for a sol’n whose [H3O+] = 4.46 x10-3M? • If pH = 9.72, what’s [H3O+]? • If pH = 4.45, what’s [H3O+] ? • What is [OH-] in the above solutions? • Calc pH of 0.0158M HCl • Calc pH of 1 x 10-4M KOH and of 0.036M KOH

  29. 9.53 and 54. How many times more basic is a soln at pH 6 relative to pH 4? • How many times more acidic is a soln at pH 7 relative to pH 11?

  30. 9.3 Rxns btn acids and bases • Neutralization rxn HCl(aq) + NaOH(aq)NaCl(aq) + HOH(l) • acid + base salt+ water • In the dissociated form H+(aq) +Cl-(aq) +Na+(aq) + OH-(aq)Na+(aq) + Cl-(aq) + HOH(l) • H+(aq) + OH-(aq) g H2O(l) • This is the same for any strong acid-strong base neutralization

  31. Titration • Used to determine the conc of unknown acid or base solution. • 6.00mL of a 0.500M HCl soln is needed to neutralize 10.00mL of a NaOH soln. What is the concentration (molarity) of the NaOH soln?

  32. 15.00mL of a 0.200M NaOH soln is needed to neutralize 25.00mL of a HCl soln. What is the concentration (molarity) of the HCl soln? • Equivalence pt and end point, indicators

  33. Polyprotic acids • Some acids have more than one acidic hydrogen as: H2SO4, H2CO3, H3PO4 • H2CO3 +2NaOH  Na2CO3 + 2H2O • Undergo stepwise dissociation as: • H2CO3 + H2O  HCO3- + H3O+ • HCO3- + H2O  CO32- + H3O+

  34. 9.4 Acid-base buffers • Solutions that contain a • Buffers have the ability of withstanding small additions of added acid or base without large fluctuations in pH.

  35. How do buffers work? • How do buffers work? • HA H+ + A- • add acid A- + H+ HA • add base HA + OH-  HOH + A-

  36. Have 100mL of 3.6x10-5M HCl; pH = 4.44 • Add 1mL of 0.10M HCl to the above soln, the pH changes to 2.99; add 1mL of 0.10M NaOH to the above soln, pH changes to 11.99. • Have 100mL of a buffer that is 0.050M Na acetate and0.10M acetic acid, pH = 4.44 • Add 1mL of 0.10M HCl to the buffer, pH changes to 4.43; add 1mL of 0.10M NaOH to the buffer, pH changes to 4.46.

  37. Are these buffers: • NaOH/NaCl • NaOH/HCl • NH3/NH4Cl • NaF/HF • CH3COOH and CH3COONa • HNO3 and KNO3 • HBr and MgCl2 • H2CO3 and NaHCO3

  38. Calculating the pH of a buffer • HA(aq) + H2O(l) D H3O+(aq) + A-(aq) • Ka = [H3O+][A-] [HA] • Or [H3O+] = Ka[HA] [A-] • [H3O+] = Ka[acid] [salt]

  39. 9.65 and 66 What is [H3O+] for a buffer soln that is 0.200M in acid and 0.500M in the corresponding salt if the weak acid has Ka= 5.80 x 10-7? What is the pH of the solution?

  40. What is the pH of a solution that is 0.1M in CH3COOH and 0.1M in CH3COONa? Ka for acetic acid is 1.8 x 10-5.

  41. 9.49 and 50 For CH3COOH(aq)+H2O(l) D CH3COO-(aq) +H3O+(aq) • Write the K expression • Explain what happens if • Add strong acid • Dilute with water • Add strong base • Add more acetic acid

  42. Henderson-Hasselbalch equation • pH = pKa + log • Solve: What is [H3O+] for a buffer soln that is 0.200M in acid and 0.500M in the corresponding salt if the weak acid has Ka= 5.80 x 10-7? What is the pH of the solution? Usong Henderson-Hasselbalch

  43. 9.5 Oxidation-reduction processes • Oxidation (oxd’n) ; originally rxn in which O2 combined with another substance as C(s) + O2 CO2(g) • Also defined as

  44. Reduction (red’n) : originally rxn in which O2 was removed from a substance as Fe2O3(s) + 3C(s)  3CO(g) + 2 Fe(s) • Also defined as

  45. Let’s look at gain and loss of electrons more closely • Look at the reaction 2NaCl(s)  2Na(s) + Cl2(g) • We can divide this rxn into half rxns as

  46. And we identify oxidation and reduction as

  47. Na+ + e-  Na Cl- Cl2 + 2e- • The species (reactant, Cl-) that is oxidized is the _________ agent (reducing agent donates e-’s to the species that is reduced) • The species (reactant, Na+) that is reduced is the ___________ agent (oxidizing agent causes oxd’n of another species by accepting e-’s from that species)

  48. Identify species oxidized, reduced, oxidizing agent, reducing agent • Cu2+ + Mg g Mg2+ + Cu • Br2 + 2KI g 2KBr + I2 • 4Fe + 3O2 g 2Fe2O3 • Zn + Cu 2+gZn2+ + Cu

  49. Corrosion • Corrosion: eating away of metals by oxidation-reduction processes. • Rusting of iron 4Fe(s) + 3O2(g) g 2Fe2O3(s)

  50. Voltaic cells • Voltaic cells: spontaneous chemical rxn is used to generate electricity • Cu2+(aq) +Zn(s)  Zn2+(aq) + Cu(s) • Oxidation occurs at the______ and reduction at the _________.

More Related