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ACID BASE EQUILIBRIA

ACID BASE EQUILIBRIA. CHAPTER 14. ACID/ BASE DEFINITIONS. ARRHENIUS BRONSTED-LOWRY LEWIS. 1. ARRHENIUS. Most limited definition of acids and bases. Acids supply H + in aqueous solution. Bases supply OH -1 in aqueous solution. Limited since many bases do not contain OH -1 .

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ACID BASE EQUILIBRIA

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  1. ACID BASE EQUILIBRIA CHAPTER 14

  2. ACID/ BASE DEFINITIONS • ARRHENIUS • BRONSTED-LOWRY • LEWIS

  3. 1. ARRHENIUS • Most limited definition of acids and bases. • Acids supply H+ in aqueous solution. • Bases supply OH-1 in aqueous solution. • Limited since many bases do not contain OH-1. • Ex. NH3

  4. 2. BRONSTED-LOWRY • More general definition. • Acids are proton (H+) donors. • When an acid loses a proton, it becomes a conjugate base. • Bases are proton acceptors. • Once the base accepts a proton, it becomes a conjugate acid. • General form of a BL acid base reaction: HA + H2O  A-1 + H3O+

  5. EXAMPLES: Formic acid dissociates (HCOOH) Perchloric acid dissociates   Acetic acid dissociates

  6. Relationship between acids and their conjugates: • The stronger the acid, the weaker its conjugate base.

  7. 3. LEWIS • Most general definition. • Acids are electron pair acceptors. Ex. BF3 • Bases are electron pair donors. Ex. NH3

  8. ACID DISSOCIATION EQUILIBRIUM • Ka expressions can be written for an acid dissociation reaction. • Ka = [ products]power [reactants]power Write the Ka expressions for the sample BL reactions:

  9. Higher Ka values, more dissociation, stronger acids. • Strong acids do not have a Ka value s(extremely large) since the equilibrium lies so far to the right due to complete dissociation of a strong acid. • The strong acids are: Sulfuric Nitric Perchloric Hydrochloric Hydrobromic Hydroiodic • Other Ka values are given in the appendix. • Take note of acetic acid’s Ka value and memorize it for the AP exam.

  10. WATER • Water’s Ka value is 1.0 x 10-14 How? • Consider two waters reacting…this is called auto-ionization of water. • What is the concentration of H+ and OH-­ in a sample of water.

  11. MEASURING ACID and BASE STRENGTH • pH scale • ranges from 0-14 • pH = - log [H+] • pOH = -log [OH-1] • pH + pOH = 14

  12. For water @ 25 °C pH = 7 [H+] = 10-pH pOH = 7 [OH-1] = 10-pH Kw = [H+][OH-1] and pKw = pH + pOH = 14

  13. ACID STRENGTH and CHEMICAL STRUCTURE •  When a substance is dissolved in water, it may behave as an ACID, a BASE, or produce a NEUTRAL solution. • How does the chemical structure of a substance determine such behaviors?

  14. FACTORS AFFECTING ACID STRENGTH • Strength measured by: • Ka value • Recall Ka meaning. • pH pH = -log [H+] measures how much an acid dissociates

  15. 1. BOND POLARITY • a proton is transferred only when the H is the positive pole of the compound H + X • in hydrides, ex. NaH • the H is negatively charged, so no H+ could result. • in non-polar molecule, ex. CH4 • the electrons are not being pulled from the H, so no H+ could result.

  16. 2. BOND STRENGTH also affects acid strength • strong bond – less likely to dissociate, weaker acid • weak bond – stronger acid

  17. 3. CONJUGATE STRENGTH • stable conjugates – come from a strong acid

  18. TYPES of ACIDS and STRENGTH TRENDS • Acids that can donate more than one proton are called POLYPROTIC acids • 2 protons, ex. H2SO4= diprotic • 3 protons, ex. H3PO4 = triprotic

  19. 1. BINARY ACIDS made of H and one other element. General form H-X. • H-X bond strength determines strength of the acid. • Bond strength DECREASES as the size of X increases. • Changes more drastically down a group.

  20. Going across a period, look at ELECTRONEGATIVITY.

  21. 2. OXYACIDS • contain OH groups bound to a central atom. • OH groups acting as bases: • When OH is bound to a group with extremely low ELECTRONEGATIVITY. • Ex. metals • Ca(OH)2 • KOH

  22. OH groups acting as acids: • Bound to a NON-METAL. • Does not readily lose OH. • As the ELECTRONEGATIVITY of Y increases, so will the acidity. • The O-H bond becomes more POLAR and loss of H+ is favored. • When an additional OXYGEN is bound to the central Y, further increases the POLARITY of the OH bond favoring loss of the H+.

  23. Rules for comparing oxyacid strength: 1. for oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom 2. for oxyacids with the same central atom, acid strength increases as the number of oxygens attached to Y increases.

  24. Example: Place the following oxyacids in order of increasing strength: HClO HClO2 HClO4 HClO3 HBrO

  25. 3. CARBOXYLLIC ACIDS • contain a carboxyl group • additional oxygens attached to the carboxyl group – draws electron density from the OH group, increasing its polarity. • strength of the carboxylic acid increases when the number of electronegative atoms increases.

  26. Example: • CH3OH – is this an acid? • CH3COOH – what acid is this? How does its strength compare to CH3OH?

  27. MORE PRACTICE • Arrange the compounds in each of the following series in order of increasing acid strength: • AsH3 • HI • NaH • H2O

  28. Arrange the compounds in each of the following series in order of increasing acid strength: • H2SeO3 • H2SeO4 • H2O

  29. Explain why a. HCl is stronger than H2S as an acid. b. benzoic acid (C6H5COOH) is a stronger acid than phenol (C6H5OH) c. H2SO4 is stronger than HSO4-1

  30. Calculating pH of strong acids • Large Ka, so products are favored. • Assume the [acid]0 = [ H+ ] due to complete dissociation. • What else contributes [ H+ ] in solution? • Auto- ionization of water • Assumed to be negligible for strong acids (unless they are dilute >10-6 M) • So, pH = - log [acid]o

  31. pH example (strong) • Calculate the pH and [OH-1] of a 5 x 10_3 M solution of HClO4.

  32. pH of weak acid solutions • Ka is small • Consider all sources of H+ in solution • ICE table • Ka expression • Solve, approximate whenever possible (check <5%)

  33. pH example (weak) • Calculate the pH of a 0.500 M solution of formic acid HCOOH (Ka = 1.77 x 10-4).

  34. Percent dissociation • Shows how much of the initial acid is turned into H+ in solution. % dissociation = amount dissociated x 100 initial concentration • Calculate by comparing H+ at equilibrium to [acid]0. • Ex. What is the percent dissociation in the formic acid problem?

  35. pH of a weak acid mixture • Consider all sources of H+ and determine the one that is most significant. Assume the others to be negligible. • LeChatlier’s principle predicts that acids with lower Ka values will not be able to dissociate due to an abundance of H+ from the stronger acid’s dissociation, so the weaker acid’s dissociation is suppressed.

  36. pH example (mixture of acids) • Calculate the pH of a mixture of 2.00 M formic acid and 1.50 molar hypobromous acid (Ka = 2.06 x 10-9). • Write dissociation reactions for all species in solution. HCOOH HOBR H2O

  37. 2. Determine the source of H+ in solution 3. ICE

  38. 4. Ka expression 5. H+ at equilibrium, pH

  39. Ka from percent dissociation • A 0.500 M solution of uric acid is 1.6% dissociated. Calculate the value of Ka for uric acid.

  40. BASES • Basically, the problems are the same as acids. • Focus on [OH-] at equilibrium. • Consider all sources of [OH-] including auto-ionization of water. • pOH = - log [OH-]

  41. pH example (strong base) • Calculate the pH of a solution made by putting 4.63 grams of LiOH into water and diluting to a total volume of 400 ml.

  42. pH example (weak base) • Calculate the pH of a 0.350 M solution of methylamine, CH3NH2 (Kb = 4.38 x 10-4)

  43. Polyprotic acids • Supply more than one proton to the solution. • Each step has its own Ka value. • Ka1 >> Ka2 > Ka3 , so many times it is possible to ignore the contribution of H+ from the second (and third)dissociation. • Sulfuric acid is the exception to this rule!!!

  44. pH example (polyprotic acid) • Calculate the pH, [PO4-3] and [OH-] of a 6.0 M phosphoric acid solution. 1. Write all relevant dissociation reactions and find Ka values for each.

  45. 2. ICE, assumptions, Ka1 3. Ka2 to find [HPO4-2]

  46. 4. Ka3 to find [PO4-3] 5. [OH-], pOH, pH

  47. ACID/ BASE properties of SALTS • Ionic compounds dissociate in water and the resulting ions can make the solution acidic, basic, or neutral. • Na+ and other alkali and alkaline earth metals do NOT exhibit acid or base properties. What about Na (s)? • Conjugate of strong acids or bases do NOT exhibit acid or base properties. • Kw = Ka x Kb = 1x10 -14 • If Ka> Kb, acidic • If Kb> Ka, basic

  48. Salt examples • Predict whether each of the following will create an acid, base, or neutral aqueous solution. • Na3PO4 • KI • NH4F

  49. pH example (of a salt) • Calculate the pH of a 0.500 M NaNO2 solution (Ka = 4.0 x 10-4) Reaction Kb ICE

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