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Chapter 14 Acids and Bases

Chapter 14 Acids and Bases

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Chapter 14 Acids and Bases

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  1. Chapter 14Acids and Bases

  2. 14.1 – The Nature of Acids and Bases • Some properties of acids: • Sour tasting • Form a solution with pH < 7 (blue litmus  red) • React with metals to give a salt and H2(g) • React with bases to form a salt, water, and heat • React with metal carbonates and hydrogen carbonates to give a salt, water, and CO2

  3. 14.1 – The Nature of Acids and Bases • Some properties of bases (alkalis): • Bitter tasting, slippery in solution • Form a solution with pH > 7 (red litmus  blue) • React with acids to form salt, water, and heat

  4. 14.1 – The Nature of Acids and Bases

  5. 14.1 – The Nature of Acids and Bases • Some Arrhenius acids: • HX, CH3COOH, HCN, HNO3, H2CO3 • Some Arrhenius bases: • NaOH, Mg(OH)2, Ca(OH)2, KOH, NH4OH

  6. 14.1 – The Nature of Acids and Bases

  7. 14.1 – The Nature of Acids and Bases • Some Bronsted-Lowry acids: • All those mentioned in the Arrhenius category • Some Bronsted-Lowry bases: • Those mentioned in the Arrhenius category, and… • NH3, N2H4, SO42-, ClO3-,

  8. 14.1 – The Nature of Acids and Bases

  9. 14.1 – The Nature of Acids and Bases • Some Lewis acids: • Those mentioned in the other two categories, and… • BCl3, BF3, Al3+, SO3, H3C+ • Some Lewis bases: • Those mentioned in the other two categories, and… • H-, H3C-, H2O, BH4-, Xe:

  10. 14.1 – The Nature of Acids and Bases • The Arrhenius definition is the narrowest • The Lewis definition is the widest • We will focus on the Bronsted-Lowry definition

  11. 14.1 – The Nature of Acids and Bases • Bronsted-Lowry Model • To determine what makes up an acid or a base (on paper), react it with water and see what happens: Recall: Hydronium (H3O+) is sometimes written in the place of a proton, H+. This is because any free protons in solution will immediately react with water to form H3O+.

  12. 14.1 – The Nature of Acids and Bases • Bronsted-Lowry Model • For a general acid/base, we can write the dissociation constants (Ka/Kb)

  13. 14.2 – Acid strength • Strong acids: • Acids which dissociate completely in water (i.e., Ka is very large) • Commit all 6 of them to memory: HCl – hydrochloric acid HBr – hydrobromic acid HI – hydroiodic acid HNO3 – nitric acid *H2SO4 – sulfuric acid (only the first H+ is strong) HClO4 – perchloric acid

  14. 14.2 – Acid strength • Strong bases: • Bases which dissociate completely in water (i.e., Kb is very large) • They are: MOH M(OH)2 (where M is an alkali or alkali earth metal, except for beryllium and radium)

  15. 14.2 – Acid strength • Relative strength of Bronsted-Lowry acids/bases: • When an acid loses a proton, it becomes a conjugate base • When a base accepts a proton, it becomes a conjugate acid In every reaction we look at in this chapter, there is an acid, a base, and a conjugate acid, conjugate base

  16. 14.2 – Acid strength • Relative strength of Bronsted-Lowry acids/bases: • Example, hydrochloric acid: • HCl is the acid, and Cl- is the conjugate base • That makes H2O the base, and H3O+ the conjugate acid

  17. 14.2 – Acid strength • Relative strength of Bronsted-Lowry acids/bases: • Example, ammonia: • NH3 is the base, and NH4+ is the conjugate acid • That makes H2O the acid, and OH- the conjugate base

  18. 14.2 – Acid strength • Relative strength of Bronsted-Lowry acids/bases: • Acidity/basicity is determined by the extent of dissociation, given by the equilibrium constant, Ka/b: CH3COOH ⇌ CH3COO- + H+Ka = 1.85 x 10-5 HCl⇌ Cl- + H+Ka = ~ 1 x 107 • Recall a high Ka/b means that there are more products at equilibrium

  19. 14.2 – Acid strength • Relative strength of Bronsted-Lowry acids/bases: • These numbers are often very large or very small, so the pKa and pKb are often used. To convert between Ka/b and pKa/b, use the following: pKa = -log(Ka) and Ka = 10-pKa pKb = -log(Kb) and Kb = 10-pKb

  20. 14.2 – Acid strength • Relative strength of Bronsted-Lowry acids/bases: • The conjugate base of a strong acid is pH neutral, • Example, HBr • The conjugate acid of a strong base is pH neutral, • Example, NaOH

  21. 14.2 – Acid strength • Relative strength of Bronsted-Lowry acids/bases: • The conjugate base of a relatively strong weak acid, is a relatively weak, weak base. Example, HF,

  22. 14.2 – Acid strength • Relative strength of Bronsted-Lowry acids/bases: • The conjugate acid of a relatively strong weak base, is a relatively weak, weak acid. Example, NH3

  23. 14.2 – Acid strength

  24. 14.2 – Acid strength • Water as an acid and a base • Substances that can act as an acid or a base are called amphoteric. They can accept or donate protons. • Some amphoteric substances: • H2O • HCO3- • HSO4- • H2PO4-

  25. 14.2 – Acid strength • Water as an acid and a base • Water’s amphoteric properties allow it to auto-ionize: H2O(l) + H2O(l) ⇌ H3O+(aq) + OH-(aq) • The equilibrium constant (Kw) for this reaction at 25oC is Kw = [H3O+] [OH-] = 1.0 x 10-14 • Kw is called the ion-product constant, or the dissociation constant for water

  26. 14.2 – Acid strength • Water as an acid and a base • It is important to recognize that in any solution at 25oC, no matter what it contains, the product of [OH-] and [H+] must always be 1.0 x 10-14 • Example,

  27. 14.2 – Acid strength • Water as an acid and a base • Example,

  28. 14.2 – Acid strength • Examples,

  29. 14.2 – Acid strength • Examples,

  30. 14.3 – The pH scale • The pH scale is a logarithmic scale used to conveniently express [H+]. It is given by: pH = -log[H+] • Similarly, there exists a pOH scale, which is given by: pOH = -log[OH-]

  31. 14.3 – The pH scale • If you re-write the ion-product constant, Kw: And since Kw = 1.0 x 10-14, pKw = 14, so:

  32. 14.3 – The pH scale • In general, pX = -log[X] • Example, what is the pH of the following solutions: 1. [H+] = 0.001M 2. pOH = 2 3. [OH-] = 1 x 10-5M

  33. 14.4 – Calculating the pH of a strong acid solution • Recall that the Ka of a strong acid is a really high number (> 107). That means that you have virtually all H+ and conjugate base, X-. • These calculations are easy: the concentration of the given acid is equal to [H+]. • Examples, what is the pH of a 0.01M HCl solution?

  34. 14.4 – Calculating the pH of a strong acid solution • Examples,

  35. 14.5 – Calculating the pH of a weak acid solution • Harder calculations. • Must know the Ka/Kb of the substance. • Calculated using a RICE table

  36. 14.5 – Calculating the pH of a weak acid solution • Example, the pH of 0.1M CH3COOH (Ka = 1.8 x 10-5) R CH3COOH + H2O ⇌ CH3COO-(aq) + H3O+(aq) I 0.1M -- 0 0 C -x -- +x +x E 0.1 – x x x

  37. 14.5 – Calculating the pH of a weak acid solution • The pH of a mixture of weak acids • Example, What is the pH when 1M HCN (Ka = 6.2 x 10-10) mixed with 5M HNO2 (Ka = 4.0 x 10-4)? What is the [CN-]? The main contributor of H+ is HNO2:

  38. 14.5 – Calculating the pH of a weak acid solution • The pH of a mixture of weak acids • Example, What is the pH when 1M HCN (Ka = 6.2 x 10-10) mixed with 5M HNO2 (Ka = 4.0 x 10-4)? What is the [CN-]?

  39. 14.5 – Calculating the pH of a weak acid solution • The pH of a mixture of weak acids • Example, What is the pH when 1M HCN (Ka = 6.2 x 10-10) mixed with 5M HNO2 (Ka = 4.0 x 10-4)? What is the [CN-]? Now us the Ka of HCN to calculate [CN-]. Note, it doesn’t matter where the H+ comes from:

  40. 14.5 – Calculating the pH of a weak acid solution • Percent dissociation (percent ionization) • For a given weak acid, the percent dissociation increases as the acid becomes more dilute • Example, what is the percent dissociation of a 0.1M CH3COOH?

  41. 14.5 – Calculating the pH of a weak acid solution • Percent dissociation (percent ionization) • Example, what is the percent dissociation of a 1M CH3COOH (Ka = 1.85 x 10-5)? R CH3COOH + H2O ⇌ CH3COO-(aq) + H3O+(aq) I1M -- 00 C -x -- +x +x E 1 – x -- xx

  42. 14.5 – Calculating the pH of a weak acid solution • Percent dissociation (percent ionization) • Example, what is the percent dissociation of a 1M CH3COOH (Ka = 1.85 x 10-5)?

  43. 14.5 – Calculating the pH of a weak acid solution • Examples,

  44. 14.5 – Calculating the pH of a weak acid solution • Examples,

  45. 14.6 – Bases • Bases • A base does not have to contain a hydroxide ion • Almost all of the weak bases you will encounter will be a nitrogenous base (i.e., amines). Nitrogens typically have at least one lone pair of electrons that is capable of forming a bond with a proton: • Example, NH3 + H2O ⇌ NH4+ + OH-

  46. 14.6 – Bases • Bases • Kb always refers to the reaction of a base with water to form a hydroxide ion and the conjugate acid • Example, for NH3 + H2O ⇌ NH4+ + OH-

  47. 14.6 – Bases • Bases • Example calculations, • Molar mass of KOH = 56.1g/mol

  48. 14.6 – Bases • Bases • Example calculations,

  49. 14.6 – Bases • Bases • Example calculations,

  50. 14.7 – Polyprotic acids • Polyprotic acids • Some acids can donate more than one proton. They are called polyprotic acids. • The Ka for the dissociation of the second, third, fourth, etc… proton is always lower than the first. • Example, H3PO4