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Acid-Base Chemistry

Acid-Base Chemistry. Definitions of Acids and Bases. Arrhenius An acid is a substance that dissolves in water to increase the concentration of H + (aka H 3 O + ). HCl(g) + H 2 O(l)  H 3 O + (aq) + Cl - (aq)

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Acid-Base Chemistry

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  1. Acid-Base Chemistry

  2. Definitions of Acids and Bases Arrhenius An acid is a substance that dissolves in water to increase the concentration of H+ (aka H3O+). HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq) A base is a substance that dissolves in water to increase the concentration of OH-. NaOH(s) + H2O(l)  OH-(aq) + Na+(aq) + H2O(l) NH3(aq) + H2O(l)  OH-(aq) + NH4+(aq) The Arrhenius definition requires that the solvent be water.

  3. Definitions of Acids and Bases Brønsted-Lowry An acid is a substance that donates a proton H+ to a base. A base is a substance that accepts a proton H+ from an acid. Since the emphasis here is on proton transfer, an acid must have base with which to react. HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) acid base H2O(l) + NH3(aq)  OH-(aq) + NH4+(aq)

  4. Definitions of Acids and Bases Brønsted-Lowry This definition allows acid-base reactions to occur in more than just aqueous solutions. The following acid-base reaction occurs in the gas phase: HCl(g) + NH3(g)  NH4Cl(s) (NH4Cl is held together through an ionic bond)

  5. Definitions of Acids and Bases Lewis For a substance to be a proton acceptor, it must have an unshared pair of electrons: HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq) Lewis structure of water: H-Ö-H H2O(l) + NH3(aq) OH-(aq) + NH4+(aq) H Lewis structure of ammonia: H-N: H A Lewis base is an electron-pair donor. ..

  6. Definitions of Acids and Bases Lewis A Lewis base is an electron-pair donor. Any Brønsted-Lowry base is also a Lewis base. A Lewis acid is an electron-pair acceptor. While Brønsted-Lowry acids usually contain H, Lewis acids do not have to. CH3+ + OH- CH3OH Here CH3+ is a Lewis acid and OH- is a Lewis base.

  7. Definitions of Acids and Bases Brønsted-Lowry vs Lewis When the term “acid” is used, what is meant is a Brønsted-Lowry acid. We distinguish Lewis acids from Brønsted-Lowry acids by being careful always to use the term “Lewis acid” when referring to them.

  8. Brønsted-LowryConjugate Acid-Base Pairs A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor. In the forward reaction, water is the acid. In the reverse reaction, the ammonium ion is the acid. acidbase H2O(l) + NH3(aq) OH-(aq) + NH4+(aq)

  9. Brønsted-LowryConjugate Acid-Base Pairs Every acid has a conjugate base, formed by removing a single proton from the acid. Every base has a conjugate acid, formed by adding a single proton to the base. conjugate base conjugate acid acid base H2O(l) + NH3(aq) OH-(aq) + NH4+(aq) - H+ + H+

  10. Brønsted-LowryConjugate Acid-Base Pairs What is the conjugate base of each of the following acids? HCl H2SO4 HSO4- H2O Cl- HSO4- SO42- OH- acid conjugate base In order to use Appendix D to find the appropriate equilibrium constant values, you must be able to write the chemical equation for each of these acting as an acid: HSO4- (aq) H+(aq) + SO42-(aq)

  11. Brønsted-LowryConjugate Acid-Base Pairs What is the conjugate acid of each of the following bases? O2- SO42- PO43- H2O OH- HSO4- HPO42- H3O+ conjugate acid base In order to use Appendix D to find the appropriate equilibrium constant values, you must be able to write the chemical equation for each of these acting as a base: PO43- (aq) + H2O(l) HPO42-(aq) + OH-(aq)

  12. Brønsted-LowryConjugate Acid-Base Pairs What is the conjugate base of each of the following when they act as an acid? conjugate base acid NH3 NH4+ CH3OH CH3NH2 NH2- NH3 CH3O- CH3NH-

  13. Brønsted-LowryConjugate Acid-Base Pairs What is the conjugate acid of each of the following when they act as a base? conjugate acid base NH3 NH2- CH3OH CH3NH2 NH4+ NH3 CH3OH2+ CH3NH3+

  14. Brønsted-LowryConjugate Acid-Base Pairs Some substances can act as an acid in one reaction and a base in another. A substance capable of acting either as an acid or a base is amphoteric. acting as a base acting as an acid H2O HSO4- H2PO4- HPO42- OH- SO42- HPO42- PO43- H3O+ H2SO4 H3PO4 H2PO4- amphoteric

  15. Relative Strengths of Brønsted-LowryAcids and Bases The stronger the acid, the weaker its conjugate base. Knowing the strength of one member of the conjugate pair means you also know something about the strength of the other member of the pair. base strength increases acid strength increases

  16. Relative Strengths of Brønsted-LowryAcids and Bases Strong acids completely transfer their protons. Their conjugate bases have a negligible tendency to become protonated in aqueous solution. Know the 7 strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO3 base strength increases acid strength increases

  17. Relative Strengths of Brønsted-LowryAcids and Bases Weak acids only partially dissociate in aqueous solution. Be able to recognize a conjugate acid-base pair and be able to find EITHER K (Ka or Kb) using Appendix D. base strength increases acid strength increases

  18. Relative Strengths of Brønsted-LowryAcids and Bases The O2- ion is a really strong base. O2- + H2O  2OH- This means that OH- in water has NO tendency to act as an acid. base strength increases acid strength increases

  19. Using Relative Acid Strengths to Determine the Position of the Equilibrium In every acid-base reaction, the position of the equilibrium is the one which favors the transfer of the proton away from the stronger acid. HC2H3O2(aq) + HSO4-(aq) C2H3O2-(aq) + H2SO4(aq) The equilibrium lies to the left. base strength increases acid strength increases

  20. Autoionization of Water and the Ion-Product Constant Kw The following reaction occurs in water: H2O(l) + H2O(l) H3O+(aq) + OH-(aq) Kw(25°C) = [H3O+][OH-] = 1.0 x 10-14 The reaction is called the autoionization of water. The reaction can also be written: H2O(l) H+(aq) + OH-(aq) Kw(25°C) = [H+][OH-] = 1.0 x 10-14 Kw is called the ion-product constant of water. Memorize both the expression and its value at 25°C.

  21. The Ion-Product Constant Kw Kw(25°C) = [H+][OH-] = 1.0 x 10-14 This equilibrium holds not only for pure water, but for any aqueous solution. Anything that causes the hydroxide concentration to increase will also cause the hydrogen ion concentration to decrease. A solution where [H+] = [OH-] is said to be neutral. A solution where [H+] > [OH-] is said to be acidic. A solution where [H+] < [OH-] is said to be basic.

  22. pH The concentration of hydrogen ions in aqueous solutions is expressed as pH: pH = -log[H+] concentration of H+ is in M For a neutral solution, [H+] = [OH-] at 25°C, Kw = 1.0 x 10-14 = [H+]2 [H+] = 1.0 x 10-7 M pH = -log[H+] = -log(1.0 x 10-7) = 7

  23. pH = -log[H+] For a neutralsolution, pH = 7 For an acidicsolution,pH < 7 For abasicsolution,pH > 7 What is the pH of 0.1M HCl at 25°C? pH = -log[H+] = -log 0.1 = 1.0 What is the pH of 0.1M NaOH at 25°C? [H+] [OH-] = 1.0 x 10-14 [H+] = (1.0 x 10-14)/.1 = 1 x 10-13 pH = -log[H+] = -log 1.0 x 10-13 = 13.0 SFs: The number of digits after the decimal matches the SFs in the concentration of H+. Uncomfortable with logarithms? Review Appendix A.2

  24. pH of Common Substances at 25°C

  25. pH = -log[H+] The pH of blood is 7.35 – 7.45. What is the normal range of concentration of H+ in blood? 7.35 = -log[H+] [H+] = 10-7.35= 4.5 x 10-8 M 7.45 = -log[H+] [H+] = 10-7.45 = 3.5 x 10-8 M The normal range of [H+] in blood is 3.5 x 10-8 to 4.5 x 10-8 M. A student calculates the pH of 12M HCl and finds pH = -1.08. Did the student make a mistake? pH = -log 12 = -1.08. No, the student was right.

  26. pOH = -log[OH-] The pH of blood is 7.35 – 7.45. What is the normal range of pOH of blood and the normal range of concentration of OH- in blood? At 25°C, Kw = 1.0 x 10-14 = [H+][OH-] -log Kw = 14.00 = -log [H+] - log [OH-] pH + pOH = 14.00 When pH = 7.35, pOH = 14.00 - 7.35 = 6.65 When pH = 7.45, pOH = 14.00 - 7.45 = 6.55 The normal range of blood pOH is 6.55 - 6.65. [OH-] = 10-6.55 = 2.8 x 10-7 M [OH-] = 10-6.65 = 2.2 x 10-7 M The normal range of [OH-] in blood is 2.2 x 10-7 to 2.8 x 10-7 M.

  27. Measuring pH pH meter • Acid-base indicators • Pick the indicator that changes color in the pH range you desire.

  28. Calculating the pH and pOH of Strong Acids and Bases Memorize the 7 strong acids and 8 strong bases. 7 strong acids HClO4 HCl HClO3 HBr HNO3 HI H2SO4 What is the pH of 0.0500 M HNO3? pH = -log(0.0500) = 1.301 What is the pH of 0.0500 M H2SO4? pH ≈ -log(0.0500) =1.301 (why?) What is the pH of 0.0500 M KOH? pOH = -log(0.0500) = 1.301 pH = 14.00 - pOH = 14.00 - 1.301 = 12.70 What is the pH of 0.0500 M Ca(OH)2? pOH = -log(0.100) = 1.000 pH = 14.00 - 1.000 = 13.00 8 strong bases LiOH NaOH KOH Ca(OH)2 RbOH Sr(OH)2 CsOH Ba(OH)2

  29. Weak Acids and the Acid Dissociation Constant Ka Since weak acids do not dissociate completely, the calculation of pH must incorporate the acid-dissociation constant Ka HF(aq) + H2O(l) H3O+(aq) + F-(aq) Ka = [H3O+][F-] or Ka = [H+][F-] [HF] [HF]

  30. Calculating Ka from pH A 0.020 M solution of niacin (C5H4NCOOH) has a pH of 3.26. a) What % of niacin is ionized? b) What is Ka for niacin? C5H4NCOOH C5H4NCOO-(aq) + H+(aq) initial 0.020 M 0 M 0 M change - x M +x M +x M equilibrium 0.020-x M 10-3.26 M 10-3.26 M 0.020 - .00055 M 5.5 x 10-4M 5.5 x 10-4M a) % niacin ionized = 100 * 0.00055/0.020 = 2.7%

  31. Calculating Ka from pH A 0.020 M solution of niacin (C5H4NCOOH) has a pH of 3.26. a) What % of niacin is ionized? b) What is Ka for niacin? C5H4NCOOH C5H4NCOO-(aq) + H+(aq) initial 0.020 M change - 5.5 x 10-4M 5.5 x 10-4M 5.5 x 10-4M equilibrium 0.01945 M 5.5 x 10-4M 5.5 x 10-4M b) Ka= [C5H4NCOO-][H+] = (5.5 x 10-4)2= 1.6 x 10-5 [C5H4NCOOH] 0.01945

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