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Acids and Bases

Acids and Bases. Chapter 15. Overview. Acid /Base Theories Arrhenius, Bronsted-Lowry and Lewis Acid Strength pH Equilibria / Hydrolysis Common Ion effect Buffers omit carboxylic acids, sec 15.9, 15.10. Arrhenius Definitions review. Acid: H 3 0 +1 producer in water

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Acids and Bases

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  1. Acids and Bases Chapter 15

  2. Overview • Acid /Base Theories Arrhenius, Bronsted-Lowry and Lewis • Acid Strength • pH • Equilibria / Hydrolysis • Common Ion effect • Buffersomit carboxylic acids, sec 15.9, 15.10

  3. Arrhenius Definitionsreview • Acid: H30+1 producer in water • Base: OH-1 producer in water examples

  4. Bronsted-Lowry • Acid - a H +1 , proton donor • Base - a H +1 , proton acceptorExamples

  5. Amphiprotism The same ion or molecule behaves as both acid or base. example HCO3- 1 As acid: HCO3-1→CO3-2+ H+ As base: HCO3-1+ H+→H2CO3

  6. B-L Conjugate Pairs HCl + H2O → H3O+ + Cl- A B CA CB • Favored reaction→ weaker acid + weaker base i.e. strong acid → weak CB and strong base → weak CA • Acids can be cation or molecule • Bases can be anion or molecule • Protism: HCl is monoprotic; H2SO4 is diprotic

  7. Acid Strength for Binary AcidsHX Bonding Considerations • Bond strength (bond energy, BE) • Electronegativity differences ( ∆ En)NB: Kc called Ka; Kb

  8. Bond StrengthHX → H+1 + X -1 HX is a strong acid if proton is easily removed (labile) Compare HF HI Group TrendBE is decreasing; Acid strength is increasing or conversely -X-1 size is increasing, H+1 to X-1 attraction is decreasing, back reaction not likely

  9. ELECTRONEGATIVITY∆En Consider CH4 ----- > HF ∆ En is increasing and acid strength is increasingi.e.- bond polarity is increasing, closer to loosing H+1 Period TrendAcid strength increasing across period

  10. Oxy Acids H O X electron “drain” HOX if X= I, Br, Cl As H+1 lability increasing, acid strength increasing HOI to HOCl, acid strength increasing As more O atoms bonded to X, acid strength increases

  11. Water • Water auto ionizationH2O + H2O ↔ H3O+1 + OH-1acid base CA CBknow Kw = [H3O+1 ] [OH-1 ] = 1 x 10-14 • pH defined - log [H3O+1 ] = pH

  12. pH Scale 0 acid 7 base 14 most neutral mostand, pOH = - log [OH-1 ] So, pH + pOH = 14 = pKw

  13. Equilibria Computationswb or wa • See example 15.6A • Use appx C to look up Ka or pKa

  14. Assumptions for Eq Computations • At equilibrium can assume: ( M – x ) ~ M if [M / Ka ] > 100 ( test this) i.e. x < 5% M See ex. 15.8 page 635

  15. Polyprotic Acids • Stepwise Ionization • Each step has Ka • Ka is decreasing from step 1 … step n therefore, most [H3O+1 ] comes from step 1 • ex. H3PO4, H2CO3

  16. Hydrolysis Ex. CO3-2 + H2O ↔ HCO3-1 + OH-1 anion of wa = cb NH4+1 + H2O ↔ H3O+1 + NH3 cationof wb = ca Recall, acid + base → salt + waterSo, only conjugates of wa, wb hydrolyze, or wa + sb → basic solution wb + sa → acidic solution sa + sb → neutral solution ( no hydrolysis) wa + wb → ??? varies

  17. Common Ion Effect • Example: Na Ac and HAc • NaAc → Na +1 + Ac -1 100% ionized • HAc ↔ H +1 + Ac -1 Le Chatelier predicts, if [Ac -1 ] incr, shift ←IONIZATION IS SUPRESSED BY COMMON ION

  18. BUFFERS • Components: Weak acid/ salt ofex. HAc + H2O ↔ H3O+ + Ac-and NaAc 100% →Na + +Ac-So, [Ac- ] is large

  19. Buffer Action • Write reactions-acid added/ base addedKa =[H3O+ ] [Ac-] [HAc] Solve, [H3O+ ] = Ka [HAc] [Ac-]this ratiochanges only slightly in a buffered solution

  20. Henderson-Hasselbalch Equation • Convenient for finding correct wa and salt to buffer to a given pH pH = pKa + log [ X- / HX] Look up pKa to find correct acid for given pH and use equimolar concs of acid and salt ..ie, set [X- ]= [HX]

  21. Lewis Acid and Bases • Acid is an electron pair acceptor • Base is an electron pair donoracid/ base reactions as covalent bond formation; see applications in organic chem • skip 15.9, 15.10

  22. Homework Chapter 15 TEST 2 Chapters 14 and 15

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