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Ch 9: Acids, Bases and Salts

Ch 9: Acids, Bases and Salts. Suggested Problems: 2, 6, 10 , 12 , 28-44, 82, 94-100, Bonus: 118. Acids and Bases in Aqueous Solution. Think back to Chapter 4 Acid Definition : A substance that produces H + when dissolved in H 2 O Examples: HCl, HNO 3 , H 2 SO 4 Base

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Ch 9: Acids, Bases and Salts

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  1. Ch 9: Acids, Bases and Salts Suggested Problems: 2, 6, 10, 12, 28-44, 82, 94-100, Bonus: 118

  2. Acids and Bases in Aqueous Solution • Think back to Chapter 4 • Acid • Definition: A substance that produces H+ when dissolved in H2O • Examples: HCl, HNO3, H2SO4 • Base • Definition: A substance that produces OH- when dissolved in H2O • Examples: KOH, NaOH, NH3

  3. Common Acids Sulfuric Acid H2SO4 Phosphoric acid H3PO4 Nitric acid HNO3 Acetic acid CH3CO2H Hydrochloric acid HCl

  4. Common Bases Sodium Hydroxide NaOH Ammonia NH3 Magnesium hydroxide Mg(OH)2 Calcium hydroxide Ca(OH)2

  5. Acids H2SO4 HCl H3PO4 HNO3 CH3CO2H Chemical Formulas Begin with H Formulas Contain CO2H Bases NaOH Ca(OH)2 Mg(OH)2 NH3 Chemical Formulas Contain OH Common Acids and Bases

  6. Arrhenius Acids and Bases • 1903 Chemistry Nobel Prize • Barely Awarded Ph.D. • Technicality issue with Arrhenius acid definition • H+ is very reactive

  7. Updated Definitions • Arrhenius acids • Substances that produce H3O+ when dissolved in H2O • Arrhenius bases • Substances that produce OH- when dissolved in H2O • What if the reactions are not in H2O?

  8. Brønsted-Lowry Acids and Bases Separately developed the same theory pertaining to acids and bases in 1923 Johannes Brønsted Thomas Lowry

  9. Brønsted-Lowry Acid • Definition: any substance that is able to give a hydrogen ion H+, to another molecule or ion • Proton donor • Not limited to reactions in H2O • Do not have to create appreciable [H3O+] • NaOH(s) + HCl(aq) CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2-(aq) NaCl(aq) + H2O(l)

  10. Brønsted-Lowry Acids • Different acids can donate different numbers of H+ monoprotic HCl 1 diprotic H2SO4 2 triprotic H3PO4 3 1 monoprotic CH3CO2H

  11. Brønsted-Lowry Bases • Definition: a substance that accepts H+ from an acid • Proton Acceptor • Not limited to reactions in H2O • Do not have to create appreciable [OH-] • NH3(g) + HCl(g) CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2-(aq) B-L Base NH4Cl(s)

  12. Identify the following as Brønsted-Lowry Acids, Bases or Neither • HCN • AlCl3 • H2CO3 • CH3CO2- • Mg2+ • CH3NH3+ Acid Neither Acid Base Neither Acid

  13. Brønsted-Lowry Acids and Bases • Summary • Acid-Base reaction is one in which a proton is transferred • Use B or B- to represent bases • Use HA to represent acids • Using the symbols B, B- and HA write two general acid-base chemical reactions B + HA BH+ + A- B- + HA BH + A-

  14. Brønsted-Lowry Acids and Bases • Consequence of Brønsted-Lowry Definition • What are species BH+, BH, and A- ? • BH+ • BH • A- • Acid-Base reactions are reversible • K is often large, resulting in a single preferred direction Acid Acid Base

  15. Conjugate Acid – Base Pairs • Definition: A pair of compounds whose formula differ only by one proton. • After an acid donates a proton, the remaining species turns into a conjugate base(CB). • After a base accepts a proton, the resulting species turns into a conjugate acid (CA).

  16. - H+ +H+ Conjugate Acid – Base Pairs Example HF(aq) + H2O(l) H3O+(aq)+ F-(aq) Acid Base Conjugate Acid Conjugate Base + H+

  17. +H+ Conjugate Acid – Base Pairs Example + H+ NH3(g) + H2O(l) NH4+(aq)+ OH-(aq) Base Acid Conjugate Acid Conjugate Base - H+

  18. Water as Both an Acid and a Base NH3(g) + H2O(l)NH4+(aq) + OH-(aq) HF(aq) + H2O(l) H3O+(aq)+ F-(aq) Base Acid • A substance that can react as an acid or a base is called amphoteric

  19. Common Acid-Base Reactions • Neutralization reaction • Acid with a metal hydroxide • Salt: anion of the acid with the cation of the base • HCl(aq) + KOH(aq) KCl(aq) + H2O(l) • Why is this called a neutralization reaction? • Net Ionic Equation • H+(aq) + OH-(aq) H2O(l) Products are always salt and H2O

  20. Acid-Base Reactions • Write the balanced chemical equation for the reaction of sulfuric acid with magnesium hydroxide.

  21. Acid-Base Reactions • Acid with bicarbonate and carbonate ions • Bicarbonate ion • HCO3- • H+(aq) + HCO3-(aq)[H2CO3(aq)] CO2(g) + H2O(l) • Carbonate ion • CO32- • 2H+(aq) + CO32-(aq)[H2CO3(aq)] CO2(g) + H2O(l) • Products of these reactions are salt, CO2 and H2O

  22. Acid-Base Reactions • Write the balanced chemical reaction for nitric acid with baking soda (sodium bicarbonate).

  23. Acid-Base Reactions • Acid with Ammonia • Products for this general reaction are ammonium salts • NH3(aq) + HNO3(aq) NH4NO3(aq) • Write the balanced chemical reaction for ammonia with sulfuric acid.

  24. Challenge Problem • An over the counter antacid has NaAl(OH)2CO3 as the active ingredient. • How many grams of this antacid are required to nuetralize 15.0 mL of 0.0955 M HCl?

  25. The Self Ionization of Water • H2O is amphoteric • But what if you have just pure water? H2O(l) + H2O(l)⇌ H3O+(aq) + OH-(aq) • This equilibrium is governed by the equilibrium constantKw

  26. Equilibrium Constant • Equilibrium constant (K) is equal to the concentration of the products divided by the reactants • aA + bB cC + dD [x] = concentration of species X in molarity

  27. KW H2O(l) + H2O(l)⇌ H3O+(aq) + OH-(aq)

  28. Strong Acids • Strong acids give away all of their hydrogen ions • For example, HCl is a strong acid, and when HCl dissolves in water: HCl  H+ + Cl-

  29. Weak Acids • Weak acids do not give away their H+ ions, and are in equilibrium with their ionized form • Most acids are weak acids • For example, acetic acid is a weak acid, and when HC2H3O2 dissolves in water: HC2H3O2⇌ H+ + C2H3O2-

  30. Strong Bases • A strong base will give away all of its hydroxide ions (OH-) • For example, NaOH is a strong base, and when NaOH dissolves in water: NaOH  Na+ + OH-

  31. Weak Bases • To think about weak bases you must think in terms of a proton acceptor not in terms of OH-. (Brønsted-Lowry Base) • Weak bases accept some H+. • Again as with weak acids there is an equilibrium present.

  32. Strong/Weak Acids and Bases • The description of strong/weak has nothing to do with concentration • Concentration is independent of it being strong or weak. • Concentration is a measure of the amount of moles per liter • You can have low concentrations of Strong Acid and Bases

  33. pH and pOH • pH informs a person about whehter or not a solution is acid or basic • pH = -log[H+] • pOH = -log[OH-] • pH of 7 is nuetral • pH less than 7 is acidic • pH greater than 7 is basic pH + pOH = 14

  34. pH and pOH calculation examples • Determine the concentration of H3O+ and OH- from the following pH values • pH = 9.0 • pH = 3.0 • pH = 11.0

  35. pH and pOH Calculations • This course will deal only with non-equilibrium acids and bases when calculating pH or pOH • Therefore the concentration of H+ and OH- will be able to be determined from the stoichiometry of the formula. • For example • What is the pH of a solution of 0.10 M HCl? • What is the pH of a solution of 0.20 M NaOH?

  36. Salts • Definition: a substance composed of the cation of a base with the anion of the acid • Need to discuss Equivalent units • This term is terribly misused by the medical and biological profession • A equivalent is the quantity of material necessary to deliver one unit of chemical reactivity • It makes no sense outside of the context of a chemical reaction! • However, in blood analysis, Equivalents = moles x charge on ion

  37. Equivalents Example • Determine the number of equivalents in the following: • 0.10 mol of NaCl • 0.10 mol of CaCl2 • Only consider either the positive or negative charges not both and the origination of the species is also important

  38. Equivalent Example • A sample of blood serum contains 0.139 eq/L of Na+ ion. Assume the Na+ comes from dissolved NaCl, and calculate the number of equivalents, number of moles and number of grams of NaCl in 250 mL of the serum.

  39. Titration Calculation • A 25 mL sample of vinegar (which contains acetic acid) is titrated with 0.100 M NaOH. If 6.75 mL of NaOH are required, what is the molarity of the acetic acid in vinegar? 0.100 M NaOH 25 mL of vinegar

  40. Titration Example • A 25.0 mL sample of H2SO4 solution requires the addition of 16.3 mL of 0.200 M NaOH solution to reach the equivalence point. What is the concentration of the acid?

  41. Buffers • A pH buffer is a solution that resists changes in pH • A pH buffer must contain a weak acid (HA) and its conjugate base (A-) HA + OH- H2O + A- A- + H+  HA

  42. Buffer Example • Carbonic acid and bicarbonate are important blood buffers

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