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Acids and Bases

Acids and Bases. The Secret Life of the Proton. Properties of acids and bases. Properties that you probably shouldn’t test Taste: acids are sour, bases bitter Feel: Bases feel slippery, acids don’t Properties you can test Reaction with metals:

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Acids and Bases

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  1. Acids and Bases The Secret Life of the Proton

  2. Properties of acids and bases • Properties that you probably shouldn’t test • Taste: acids are sour, bases bitter • Feel: Bases feel slippery, acids don’t • Properties you can test • Reaction with metals: • Acids react with active metals to release H2, bases don’t (corrosiveness)

  3. Properties of acids and bases • Bases turn organic materials into organic salts (soap) • Acids and bases conduct electricity in water solution (ionization) • Indicators change color in presence of acid or base • Acids and bases react together to form a salt and water (neutralization)

  4. Definitions • Self-ionization of water: H2O + H2O  H3O+ + OH– • Svante Arrhenius (1859-1927) • Acids dissociate in water to give H+ ion, bases give OH- H+ + H2O  H3O+

  5. Problems with Arrhenius definition • Did not take into account solvent effects • Did not explain why NH3 is a base • Did not explain basic salts • H+ does not exist in water

  6. Brønsted/Lowry definition • Acids are proton donors, and bases are proton acceptors HCl + H2O  H3O+ + Cl- acid base CH3COOH + OH- CH3COO- + H2O acid base

  7. Brønsted/Lowry definition • Some species can act as an acid or a base – amphiprotic (usually a negative ion with H in front) HCO3- + H3O+ H2CO3 + H2O base acid HCO3- + OH- CO3-2 + H2O acid base

  8. Brønsted/Lowry definition • Water is also amphiprotic H2O + NH2- NH3 + OH- acid base H2O + C7H7SO3H  C7H7SO3- + H3O+ base acid • Amphoteric substances also can act as either acids or bases, but not necessarily involving a proton

  9. Brønsted/Lowry definition • Definition does not depend on water H- + NH3 NH2- + H2 base acid C7H7SO3H + C2H5OH  C7H7SO3- + C2H5OH2+ acid base

  10. Brønsted/Lowry definition • Conjugate pairs • Any acid has a conjugate base; any base has a conjugate acid HF + H2O  H3O+ + F- acid base acid base • This is an ionization reaction

  11. Brønsted/Lowry definition C2H3O2- + H2O  HC2H3O2 + OH- base acid acid base • This is a hydrolysis reaction, responsible for basicity of some salts. • Mono- and polyprotic acids • Monoprotic acids have one acidic proton, i.e. HCl, HNO3, HC2H3O2 (or CH3COOH).

  12. Brønsted/Lowry definition • Polyprotic acids have two or more acidic protons, i.e. H2SO4, H3PO4 • In neutralization reactions all of the acidic protons react • Anhydrides • Nonmetal oxides are acid anhydrides SO3 + H2O  H2SO4 CO2 + H2O  H2CO3

  13. Brønsted/Lowry definition • Organic acid anhydrides are condensation dimers of the parent acids (missing water) acetic anhydride acetic acid

  14. Brønsted/Lowry definition • Metal oxides are basic anhydrides Na2O + H2O  2NaOH CaO + H2O  Ca(OH)2 • Metal oxides react with water to produce hydroxides

  15. Lewis acids and bases • Lewis acids are electron acceptors, and Lewis bases are electron donors • Examples: NH3 – Lewis base because of lone pair of electrons • BF3 – Lewis acid because of electron deficit on boron, and electron w/d characteristics of fluorine • NH3 + BH3 NH3BF3 base acid • Strongly associated complex formed

  16. Lewis Acids and Bases • Water can act as either because of its polarity • Hydrogen bonds are results of a Lewis acid/base reaction • Oxidation and reduction are Lewis processes Na +H2ONa+ + OH- +H2(g) base acid

  17. Acid/base strength • A strong acid or base dissociates completely in water • Examples: HCl, H2SO4, HClO4, HNO3, most mineral acids, soluble hydroxides such as NaOH HCl + H2O  H3O+ + Cl- NaOH  Na+ + OH-

  18. Acid/base strength • A weak acid or base does not dissociate very much • Examples: HF, most organic acids, NH3 • Has very little to do with corrosiveness or reactivity

  19. Acid/base strength • Conjugate pairs • The stronger the acid, the weaker the conjugate base HCl + H2O  H3O+ + Cl- strong acid weak base strong acid weak base HF + H2O  H3O+ + F- weak acid weak base strong acid strong base

  20. Acid ionization constants • Ka is an equilibrium constant • H2O + HF  H3O++F- • Ka=[H3O+][F-]/[HF] • Water is left out because it is constant • Higher Ka means stronger acid • Typical Ka values: 10-4 to10-11 for weak acids.

  21. Acid ionization constants • Degree of dissociation is related to Ka • Example: Find value of Ka for HCN if a 0.1M sample is 0.0063% ionized. Ka=[H3O+][CN-]/[HCN] [H3O+]=[CN-]=0.0063% of 0.1M=0.0000063M [HCN]=0.1-.00000630.1 Ka=0.00000632/0.1=4x10-10

  22. Acid ionization constants • Find % ionization of 0.15M carbonic acid, Ka=4.4x10-7 Assume [H2CO3]0.15M in ionized acid Ka=4.4x10-7= [HCO3-][H3O+]/[H2CO3]=y2/0.15 y=2.6x10-4M % ionization = 2.6x10-4M/0.15M =0.17%

  23. Acid base nomenclature • Bases are usually metal hydroxides and are named as salts [i.e. KOH – potassium hydroxide, Fe(OH)3 – iron (III) hydroxide] • Nonmetallic bases are usually derivatives of ammonia (NH3) – example C6H5NH2 – aniline • Hydrides are basic (NaH, LiAlH4) and are named as salts

  24. Acid-base nomenclature • Binary acids (H + another element) are named with prefix “hydro-”, suffix • “-ic” followed by “acid”. Example – HCl is hydrochloric acid • Ternary acids (oxyacids) are named after the central polyatomic anion. • “-ate” “-ic” “-ite”  “-ous” • H2SO4 – sulfuric acid • HClO – hypochlorous acid

  25. pH and pOH scale • Measures acidity, which depends on [H3O+] • pH=-log[H3O+] • Scale runs from 0-14 • pH 7=neutral • pH>7=base • pH<7=acid

  26. pH and pOH scale • Range of possible [H3O+] is 14 orders of magnitude – too large to use anything but a log scale • Derived from equilibrium constant for the self ionization of water 2H2O  H3O+ + OH- Keq=[H3O+][OH-]=Kw=1x10-14

  27. pH and pOH scale • Neutral water [H3O+]=[OH-]=10-7 so pH=7 • pOH pOH=-log[OH-] • for any solution, pH + pOH=14 (same as Kw=10-14)

  28. Indicators • Indicators are compounds containing acidic protons that have colored ion(s). They are used to identify endpoints in titrations or determine pH.

  29. pH of Weak Acid Solutions • Tend to be higher than strong acids • Find pH of 0.25M HF (Ka=6.6x10-4) • Solution: Ka=6.6x10-4= [H3O+][F-]/[HF] • [H3O+]=y=[F-], [HF]=.25-y • 6.6x10-4=y2/(.25-y)

  30. pH of Weak Acid Solutions • Assume y<<0.25, so 0.25-y0.25 • Then 6.6x10-4=y2/0.25 • y=0.0128=[H3O+] • pH=1.9 • 0.25M HCl – pH=0.6 • 20x as much hydronium in the strong acid

  31. Neutralization reactions and titration • Definition acid + base  salt + water • Only holds for Brønsted acids and bases HNO3+LiOH  LiNO3+H2O H2SO4+2KOHK2SO4+2H2O • Neutralization reaction is exothermic

  32. Titrations • Titration – use of neutralization reaction to measure amounts of acids and bases • Endpoint or equivalence point – point at which equal # of equivalents of acid and base are present • 1 equivalent = 1 mole of donated or accepted protons • HCl: 1 mole = 1 equivalent

  33. Titrations • H2SO4 1 mole=2 equivalents • H3PO4 1 mole=3 equivalents • NaOH 1 mole=1 equivalent • Ca(OH)2 1 mole=2 eq • Normality – equivalents/liter • For monoprotic acids and bases with one hydroxide, molarity = normality

  34. Titrations • For polyprotic acids and bases with more than one hydroxide (i.e. Ca(OH)2): normality = molarity x acidic protons (or hydroxides) • Titration equation: NaVa = NbVb • Equivalence point is determined by indicator color change or pH meter

  35. Titrations • Standard solution – a prepared solution of known concentration used as a standard in titrations • Sample problem – a 0.125 M standard solution of HCl is used to titrate 65 mL of an unknown solution of calcium hydroxide. The titration requires 22.3 mL HCl. What is the concentration of the base solution?

  36. Titrations • Solution: • At the endpoint, eq acid = eq base. eq acid = 0.0223L(0.125mol/L)1eq/mol =eq base = .00279eq concentration base = 0.00279eq(mol/2eq)(1/.065L) = 0.022M

  37. Titrations • Using the titration equation: • NaVa = NbVb • 0.125N(22.3) = x(65) • x = 0.043N • 0.043eq/L(mol/2eq) = 0.022M

  38. Characteristics of titrations • Strong acid/strong base – product solution is neutral, reaction goes to completion • Strong acid/weak base – result is slightly acidic NH4OH + HCl  NH4Cl + H2O • NH4+ + H2O  H3O+ + NH3 • NH4+is a good acid; Cl- is a lousy base.

  39. Characteristics of titrations • Weak acid/strong base – result is slightly basic HC2H3O2 + NaOH  NaC2H3O2 + H2O C2H3O2- + H2O  HC2H3O2 + OH- • Acetate is a decent base, but Na+ is a terrible acid. • Na+ + 2H2O  NaOH + H3O+

  40. Characteristics of titrations • Weak acid/weak base – reaction does not complete – equilibrium established. pH depends on the acid and base used. • Titration curves • Strong acid/strong base • pH stays relatively unchanged until endpoint is near, then a rapid change occurs

  41. Titration curves Titration of 50 ml of a strong acid HA 0.1000 M by NaOH 0.1000 M :

  42. Titration curves

  43. Titration curves • Titration of weak acids with strong bases • Endpoint tends to occur at pH>7 because of hydrolysis

  44. Titration curves • pH after equivalence point depends on concentration of base

  45. Titration curves • Shape of curve depends on Ka of acid – lower the Ka, the smaller the pH jump at equivalence

  46. Titration curves • Titration of weak bases with strong acids

  47. Buffers • Buffers are made from a weak acid and the salt of a weak acid – ex. HC2H3O2 + NaC2H3O2 H2O + HC2H3O2 H3O+ + C2H3O2- high low high • Since concentrations of both the acid and conjugate base are high, addition of small amounts of strong acid or base have little effect on pH

  48. Buffers • Addition of acid: H3O+ + C2H3O2- HC2H3O2 + H2O • Addition of base: OH- + HC2H3O2 H2O + C2H3O2- • pH of buffers is governed by the acid equilibrium equation for the weak acid.

  49. Buffers • Buffers with equimolar amounts of weak acid and its salt: pH = pKa • Example: Find pH of a buffer that is 0.1M in acetic acid and 0.1M in sodium acetate. Ka for acetic acid is 1.8x10-5. • Solution: 1.8x10-5 = [H3O+][C2H3O2-]/[HC2H3O2]

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