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Unit 11: Acid-Base Equilibrium Chapter 16 and 17

Unit 11: Acid-Base Equilibrium Chapter 16 and 17. Problem Set Chapter 16: 17, 21, 37, 43, 45, 61, 65, 69, 77, 79, 101, 107 Chapter 17: 19, 23, 27, 31, 41, . Write the Formula. Name the Acid. nitric acid hydrofluoric acid hydrobromic acid chloric acid acetic acid carbonic acid

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Unit 11: Acid-Base Equilibrium Chapter 16 and 17

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  1. Unit 11: Acid-Base EquilibriumChapter 16 and 17 Problem Set • Chapter 16: 17, 21, 37, 43, 45, 61, 65, 69, 77, 79, 101, 107 • Chapter 17: 19, 23, 27, 31, 41,

  2. Write the Formula • Name the Acid • nitric acid • hydrofluoric acid • hydrobromic acid • chloric acid • acetic acid • carbonic acid • tetraboric acid • (tetraborate = B4O72-) • HCl • H2S • H3PO4 • HI • HNO2 • H3P • H2SO4 WarmUp

  3. Write the Formula • Name the Acid • nitric acid • hydrofluoric acid • hydrobromic acid • chloric acid • acetic acid • carbonic acid • tetraboric acid • (tetraborate = B4O72-) • HCl • H2S • H3PO4 • HI • HNO2 • H3P • H2SO4 WarmUp

  4. Arrhenius Definition • Acids produce hydrogen ions in aqueous solution. • Bases produce hydroxide ions when dissolved in water. • Limits to aqueous solutions. • Only one kind of base. • NH3 ammonia could not be an Arrhenius base.

  5. Bronsted-Lowry Definitions • And acid is an proton (H+) donor and a base is a proton acceptor. • Acids and bases always come in pairs. • HCl is an acid.. • When it dissolves in water it gives its proton to water. • HCl(g) + H2O(l) H3O+ + Cl- • Water is a base makes hydronium ion

  6. Neutralization Reactions • Remember: Acid + Base → Water + Salt • The salt’s name and formula will be based off of the cation of the base and the anion of the acid • Example: HCl + NaOH ↔ H2O + NaCl • Neutralization reactions are one way that is used to produce pure salts. • Amphoteric Substances: can act as either an acid or a base

  7. Titration • Method used to perform neutralization and to determine concentrations of an unknown acid or base. • Titrant – added standard solution • End Point – indicator color change • Equivalence Point – [H+] = [OH-]

  8. Acid Base Calculations • In neutral solutions, [H+] and [OH-] are equal • [H+] = [OH-] = 1 x 10-7 M • For other solutions, when [H+] increases, then [OH-] will decrease and vise versa • [H+] x [OH-] = 1.0 x 10-14 • Kw = [H+] x [OH-] = 1.0 x 10-14 • This is the Ion-Product Constant for Water • Acidic Solutions: [H+] > [OH-] • Basic Solutions: [H+] < [OH-]

  9. Using pH Scale • pH: Power of Hydrogen • pH = 7 Neutral [H+] = 1 x 10-7 • pH < 7 Acidic [H] > 1 x 10-7 • pH > 7 Basic [H+] = 1 x 10-7 • pH = -log [H+] • Example Calculate pH if [H+] = 1 x 10-7 M pH = -log (1 x 10-7 M) = -(log 1 + log 10-7) = -(0 +(-7)) pH = 7 * The exponent shows you the pH!*

  10. The pOH Scale • pOH : Power of Hydroxide • pOH = 7 Neutral[OH-] = 1 x 10-7 • pOH < 7 Basic [OH-] > 1 x 10-7 • pOH > 7 Acidic [OH-] < 1 x 10-7 • pOH = -log [OH-] *Same process as before, only now you are dealing with [OH-] Other Key Equations • pH + pOH = 14 • pH = 14 – pOH • pOH = 14 – pH

  11. Calculations: Complete the following chart

  12. How do we know if something is a strong or weak acid/base? We must calculate a value for the ionic dissociation for that substance. ACID Ka = acid dissociation constant: the ratio of the concentration of the dissociated form of the acid to the concentration of the undissociated form. BASE Kb = base dissociation constant: the ratio of the product of the conjugate acid and OH concentrations to the concentration of the conjugate base

  13. Strong and Weak Acids and Bases • Dependent on how they dissolve in water • Strong Acids completely ionize in aqueous solutions • Weak acids only partially ionize in aqueous solutions • Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solutions • Weak bases react with water to form OH and the conjugate acid of the base **Memorize Strong Acids and Bases**

  14. Memorize the 8 strong acids… all others are weak • HCl hydrochloric acid • HNO3 nitric acid • HBrhydrobromic acid • HIO4 periodic acid • HI hydroiodic acid • H2SO4 sulfuric acid • HClO4 perchloric acid • HClO3 chloric acid

  15. Writing the Ka Expression • Example HCl + H2O ↔ H3O+ + Cl- • First you need to write a Keq expression • Second, water can be eliminated because its concentration is a constant. • Write the Ka expression for this acid: CH3COOH + H2O ↔ H3O+ + CH3COO-

  16. Significance of Ka • Ka indicated the fraction of acid in the ionized form • Weak acids have small Ka values • Strong acids have large Ka values because their ionization is more complete

  17. Writing the Kb Expression • Example NH3 + H2O ↔ NH4+ + OH- • Remember: • Kb tells us how weak bases compete for OH- from strong bases • The smaller the value for Kb, the weaker the base

  18. Relationships • KW = [H+][OH-] • -log KW = -log([H+][OH-]) • -log KW = -log[H+]+ -log[OH-] • pKW = pH + pOH • KW = 1.0 x10-14 • 14.00 = pH + pOH • [H+],[OH-],pH and pOH Given any one of these we can find the other three.

  19. Example • Calculate the pH of 2.0 M acetic acid HC2H3O2 with a Ka 1.8 x10-5 • Calculate pOH, [OH-], [H+]

  20. A mixture of Weak Acids • The process is the same. • Determine the major species. • The stronger will predominate. • Bigger Ka if concentrations are comparable • Calculate the pH of a mixture 1.20 M HF (Ka = 7.2 x 10-4) and 3.4 M HOC6H5 (Ka = 1.6 x 10-10)

  21. Salt Hydrolysis • Most salt solutions are neutral, but some can be acidic or basic • Salts like sodium chloride and potassium sulfate are neutral WHY? • Neutral salts are made from strong acids and strong bases • Salt Hydrolysis – the cations or anions of the dissolved salt remove H+ from, or donate H+ to water • A solution is either acidic or basic based on H+ transfer.

  22. Salt Cations/Anions Effect on pH In water: • If cation and anion do not react, then pH = neutral • If cation does not react, but anion does, then hydroxide is produced, pH = basic • If cation reacts but anion does not, then H are produced, pH = acidic

  23. If both react: OH- and H3O+ both produced • Anions from strong acids will not change pH (ex: Br-) • Anions from a weak acid will increase pH (ex: CN-) • Cations from weak bases will decrease pH (ex: CH3NH3+) • Group 1A ions will increase pH • Group 2A ions will decrease pH

  24. What is the pH of a 0.140M solution of sodium acetate? C2H3O2- + HOH↔HC2H3O2 + OH- But what is Kb? KbxKa = Kw Kb = Kw/Ka = 1x10-14/1.8x10-5 Solve for x x = 8.77x10-6 =[OH-] pOH=5.06, pH = 8.94 so the solution is basic.

  25. The Common Ion Effect • When the salt with the anion of a weak acid is added to that acid, • It reverses the dissociation of the acid. • Lowers the percent dissociation of the acid. • The same principle applies to salts with the cation of a weak base.. • The calculations are the same as last chapter.

  26. Types of Acids • MonoproticAcied – one acidic hydrogen • PolyproticAcids- more than 1 acidic hydrogen (diprotic, triprotic). • Oxyacids - Proton is attached to the oxygen of an ion. • Organic acids contain the Carboxyl group -COOH with the H attached to O • Generally very weak. • H2SO3 ↔ H+ + HSO3- Ka1 = 1.7 x 10-2 • HSO3- ↔ H+ + SO32- Ka2 = 6.4 x 10-8 • Why is K smaller? • It is always easier to remove the first proton

  27. Structure and Acid-Base Properties • Any molecule with an H in it is a potential acid. • The stronger the X-H bond the less acidic (compare bond dissociation energies). • The more polar the X-H bond the stronger the acid (use electronegativities). • The more polar H-O-X bond -stronger acid.

  28. Strength of oxyacids • The more oxygen hooked to the central atom, the more acidic the hydrogen. • HClO4 > HClO3 > HClO2 > HClO • Remember that the H is attached to an oxygen atom. • The oxygens are electronegative • Pull electrons away from hydrogen

  29. Cl O H Strength of oxyacids Electron Density

  30. Cl O H Strength of oxyacids Electron Density O

  31. Cl O H Strength of oxyacids Electron Density O O

  32. Cl O H Strength of oxyacids Electron Density O O O

  33. Calculate the Concentration • Of all the ions in a solution of 1.00 M Arsenic acid H3AsO4 • Ka1 = 5.0 x 10-3 • Ka2 = 8.0 x 10-8 • Ka3 = 6.0 x 10-10

  34. Sulfuric acid is special • In first step it is a strong acid. • Ka2 = 1.2 x 10-2 • Calculate the concentrations in a 2.0 M solution of H2SO4 • Calculate the concentrations in a 2.0 x 10-3 M solution of H2SO4

  35. Highly charged metal ions pull the electrons of surrounding water molecules toward them. Make it easier for H+ to come off. Hydrated metals H O Al+3 H

  36. Acid-Base Properties of Oxides • Non-metal oxides dissolved in water can make acids. • SO3 (g) + H2O(l) H2SO4(aq) • Ionic oxides dissolve in water to produce bases. • CaO(s) + H2O(l) Ca(OH)2(aq)

  37. Lewis Acids and Bases • Most general definition. • Acids are electron pair acceptors. • Bases are electron pair donors. F H B F :N H F H

  38. Lewis Acids and Bases • Boron triflouride wants more electrons. F H B F :N H F H

  39. Lewis Acids and Bases • Boron triflouride wants more electrons. • BF3 is Lewis base NH3 is a Lewis Acid. H F F H B N F H

  40. ) ) 6 Lewis Acids and Bases ( H Al+3 + 6 O H +3 ( H Al O H

  41. Buffers • Solutions in which the pH remains relatively constant when small amounts of acid of base are added. • Acidic Buffer – composed of a weak acid and one of its salts • Basic Buffer – composed of a weak base and one of its salts • Buffer Capacity – the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs

  42. Buffers have the ability to acquire or give a way a H+, which allows it to maintain a standard pH • Example: Ethanoic Acid/Ethanoate system (Acetic Acid / Acetate) CH3COO- + H+ ↔ CHCOOH CH3COOH + OH- ↔ CHCOO- + H2O Both reactions occur simultaneously to maintain pH

  43. pH of Buffers • To determine pH of Buffers, we use the Henderson-Hasselbalch Equation • Calculate the pH of a solution that is .50 M HAc and .25 M NaAc (Ka = 1.8 x 10-5)

  44. Calculate the pH of the following mixtures • 0.75 M lactic acid (HC3H5O3) and 0.25 M sodium lactate (Ka = 1.4 x 10-4) • 0.25 M NH3 and 0.40 M NH4Cl(Kb = 1.8 x 10-5)

  45. Titration Curves

  46. Strong acid with strong Base • Equivalence at pH 7 7 pH mL of Base added

  47. Weak acid with strong Base • Equivalence at pH >7 >7 pH mL of Base added

  48. Strong base with strong acid • Equivalence at pH 7 7 pH mL of Base added

  49. Weak base with strong acid • Equivalence at pH <7 <7 pH mL of Base added

  50. Summary • Strong acid and base just stoichiometry. • Determine Ka, use for 0 mL base • Weak acid before equivalence point • Stoichiometry first • Then Henderson-Hasselbach • Weak acid at equivalence point Kb • Weak base after equivalence - leftover strong base.

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