1 / 74

Chapter 14 Acids and Bases

Chapter 14 Acids and Bases. 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 H 2(g) React with bases to form a salt, water, and heat

nenet
Télécharger la présentation

Chapter 14 Acids and Bases

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related