1 / 48

Arhenius Definition

Arhenius Definition. Acids produce hydrogen ions in aqueous solution. Bases produce hydroxide ions when dissolved in water. Theory limited to aqueous solutions. Only one kind of base exists (ex. NaOH). NH 3 ammonia could not be an Arhenius base. Br ö nsted-Lowry Definitions.

pquinlan
Download Presentation

Arhenius Definition

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. Arhenius Definition • Acids produce hydrogen ions in aqueous solution. • Bases produce hydroxide ions when dissolved in water. • Theory limited to aqueous solutions. • Only one kind of base exists (ex. NaOH). • NH3 ammonia could not be an Arhenius base.

  2. Brönsted-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: H3O+

  3. Acid/Base Conjugate Pairs • General equation • HA(aq) + H2O(l) H3O+(aq) + A-(aq) • Acid + Base Conjugate acid + Conjugate base • This is an equilibrium. • Competitionfor H+between H2O and A- • The stronger base controls direction. • If H2O is a stronger base it takes the H+ • Equilibrium moves to right.

  4. Conjugate Acid/Base Pairs (new) • Conjugate acid-base pair: aacid and its conjugate base (HA/A-: HSO4-/SO4-2)) and base and its conjugate acid ( A-/HA: H2O/H3O+1) • Every Brönsted acid has its conjugate base • Every Brönsted base has its conjugate acid

  5. The Acid/Base Properties of Water • Water behaves as both an acid and a base. • Water auto ionizes • 2H2O(l) H3O+(aq) + OH-(aq) • KW= [H3O+][OH-]=[H+][OH-] • At 25ºC KW = 1.0 x10-14 • In EVERY aqueous solution. • Neutral solution [H+] = [OH-]= 1.0 x10-7 • Acidic solution [H+] > [OH-] • Basic solution [H+] < [OH-]

  6. Equilibrium Constant for Water • Kw called ion-product constant: product of ion concentrations of H+ and OH- at a a particular temperature • At 25 ºC Kw = 1.0 x 10-14 • [H3O+1] = [OH-] = 1.0 x 10-7 • Neutral solution: [H3O+1] = [OH-] • Acidic solution: [H3O+1] > [OH-] • Basic solution: [H3O+1] < [OH-]

  7. Properties of Water-Example • Example • The concentration of OH-1 in a certain household ammonia cleaning solution is 0.0025 M. Calculate the concentration of H+1 ions. • Answer: 4.0 x 10-12 M

  8. Amphoteric/Amphiprotic Substances • Substances behave as both acid and a base • Examples: • HSO4-1(aq) + H2O ↔ SO4-2 + H3O+1 • HSO4-1(aq) + H2O ↔ H2SO4 + OH-1 • HCO3-1 + H2O↔ • H2PO4-1 + H2O↔ • HS-1 + H2O ↔

  9. pH : Measure of Acidity • [H3O+1] = 10-pHpH= -log[H3O+] • Used because [H+] is usually very small • As pH decreases, [H+] increases exponentially • Sig figs: only the digits after the decimal place of a pH are significant • [H+] = 1.0 x 10-8 pH= 8.00 2 sig figs • pOH= -log[OH-] • pKa = -log K

  10. 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.

  11. [H+] 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 pH 0 1 3 5 7 9 11 13 14 14 13 11 9 7 5 3 1 0 pOH 10-14 10-13 10-11 10-9 Basic 10-7 10-5 10-3 10-1 100 [OH-] Acidic Neutral Basic

  12. Calculating pH of Solutions • Always write down the major ions in solution. • Remember these are equilibria. • Remember the chemistry. • Don’t try to memorize: there is no one way to do this.

  13. pH – Problem Sets (1) • Example #1: determine the pH of 0.00015 M HCl solution. • Answer: 3.82 • Example 2: What is the [H3O+] in a solution with pH = 2.19 • Ans: 6.5 x 10-3

  14. pH- Problem Set (2) • In NaOH solution [OH-] is 2.9 x 10-4 M. Calculate the pH of the solution • Answer: 10.46

  15. The Strength of Acids and Bases • Determined by the degree of ionization of the substance (HCl versus CH3COOH) • The stronger the acid, the weaker its conjugate base • Reaction favored towards the weaker species (weak acid or weak base) • H3O+ the strongest species in aqueous solution. • Weaker acids react with water to produce H3O+1, but the equilibrium is to the left (reactants)

  16. Strong Acids • Acids stronger than H3O+1 • 100% dissociation in aqueous solution • Reaction goes to completion • HCl + H2O → Cl- + H3O+

  17. Weak Acids • Acids weaker than H3O+1 react with water in a reversible reaction • Product: H3O+1 and conjugate base • HF(aq) + H2O(l) ↔ H3O+1(aq) + F-(aq) • Equilibrium primarily goes towards reactants (left)

  18. Bases • The OH- is the strongest base that can exist in H2O solution • Strong bases dissociate completely (100% dissociation): NaOH • NH2-1(aq) + H2O(l) → NH3(aq) + OH-(aq) • Weak bases: react with water in a reversible reaction • NH3(aq) + H2O(l) ↔ NH4+1 + OH-

  19. Problem Set (1) • Calculate the pH of • A 1.0 x10-3 M HCl solution (strong acid) • A 0.020 M Ba(OH)2 solution (strong base) Answers: a) 3.00 b) 12.60

  20. Problem Set (2) • Predict the direction of the following reaction in aqueous solution. Explain: Note: Use Table 15.2, page 611 HNO2(aq) + CN-(aq) ↔ HCN(aq) + NO2-1(aq) Answer: from left to right

  21. Types of Acids • Polyprotic Acids- 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.

  22. Strong Acids • HBr, HI, HCl, HNO3, H2SO4, HClO4 • ALWAYS WRITE THE MAJOR SPECIES ( H+ and Cl-) • Completely dissociated • [H+] = [HA] • [OH-] is going to be small because of equilibrium • 10-14 = [H+][OH-] • [If [HA]< 10-7 water contributes H+]

  23. Binary Acids and Strength • Depends on: • Bond dissociation energy (tendency for bond breakage) • Radius of the anion (the larger the anion, the smaller the tendency to attract the H+)

  24. 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.

  25. Binary Acids and Strength

  26. Binary Acids and Electronegativity

  27. Weak Acids • Ka will be small. • ALWAYS WRITE THE MAJOR SPECIES. • It will be an equilibrium problem from the start. • Determine whether most of the H+ will come from the acid or the water. • Compare Ka or Kw • Rest is just like last chapter.

  28. Bases • The OH-is a strong base. • Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved. • The hydroxides of alkaline earths Ca(OH)2 etc. are strong dibasic bases, but they don’t dissolve well in water. • Used as antacids because [OH-] can’t build up.

  29. Bases without OH- • Bases are proton acceptors. • NH3 + H2O NH4+ + OH- • It is the lone pair on nitrogen that accepts the proton. • Many weak bases contain N • B(aq) + H2O(l) BH+(aq) + OH- (aq) • Kb = [BH+][OH- ] [B]

  30. Strength of Bases • Hydroxides are strong. • Others are weak. • Smaller Kb weaker base. • Calculate the pH of a solution of 4.0 M pyridine(Kb = 1.7 x 10-9) N:

  31. Polyprotic acids • Always dissociate stepwise. • The first H+ comes of much easier than the second. • Ka for the first step is much bigger than Ka for the second. • Denoted Ka1, Ka2, Ka3

  32. Polyprotic acid • H2CO3 H+ + HCO3- Ka1= 4.3 x 10-7 • HCO3- H+ + CO3-2 Ka2= 4.3 x 10-10 • Base in first step is acid in second. • In calculations we can normally ignore the second dissociation.

  33. 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

  34. Cl O H Strength of oxyacids Electron Density

  35. Cl O H Strength of oxyacids Electron Density O

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

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

  38. Acid Base Reactions: Review • Strong acid – strong base: neutralization reaction • Weak acid-strong base: equilibrium HF(aq) + H2O(aq) ↔ H3O+1(aq) + F-(aq) • Strong acid- weak base: nitric acid + ammonia (equimolar quantities) H+(aq) + NH3 (aq) → NH4+1(aq) NH4+1(aq) + H2O(l) ↔ NH3(aq) + H3O+1(aq) Solution: I ACIDIC

  39. Acid Base Reactions: Review • Weak acid- weak base • NH3 (aq) + CH3COOH(aq) →CHCOO-(aq) + NH4+1(aq) • CH3COO-(aq) + H2O(l) ↔ CH3COOH + OH-(aq) • NH4+1(aq) + H2O(l) ↔ NH3(aq) + H3O+1(aq) The pH of the solution depends on the relative strengths of the acid and base: calculate – next chapter

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

  41. 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) • Amphoteric oxides: • Al(OH)3(s) + 3H+(aq) →Al+3 (aq)+ 3H2O(l) • Al(OH)3(s) + OH-(aq) ↔ Al(OH)4-1(aq)

  42. 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

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

  44. 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

  45. ) ) 6 Lewis Acids and Bases ( H Al+3 + 6 O H +3 ( H Al O H

  46. The End The End

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

More Related