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Chemistry 102(01) Fall 2010

Chemistry 102(01) Fall 2010. Instructor: Dr. Upali Siriwardane e-mail : upali@latech.edu Office : CTH 311 Phone 257-4941 Office Hours : M,W 8:00-9:00 & 11:00-12:00 am; Tu,Th,F   8:00 - 10:00 am.

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Chemistry 102(01) Fall 2010

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  1. Chemistry 102(01) Fall 2010 Instructor: Dr. UpaliSiriwardane e-mail: upali@latech.edu Office: CTH 311 Phone 257-4941 Office Hours: M,W 8:00-9:00 & 11:00-12:00 am; Tu,Th,F  8:00 - 10:00 am. Test Dates: September 23, October 21, and November 16; Comprehensive Final Exam: November 18, 2010   Exam: 10:0-10:15 am, CTH 328. September 23,2010 (Test 1): Chapter 13 October 21, 2010 (Test 2): Chapters 14 & 15 November 16, 2010 (Test 3):Chapters 16, 17 & 18 Comprehensive Final Exam: November 18, 2010 :Chapters 13, 14, 15, 16, 17 and 18

  2. Chapter 16. Acids and Bases 16.1 The Brønsted-Lowry Concept of Acids and Bases 16.2 Types of acids/bases:Organic Acids and Amines 16.3 The Autoionization of Water 16.4 The pH Scale 16.5 Ionization Constants of Acids and Bases 16.6 Problem Solving Using Ka and Kb 16.7 Molecular Structure and Acid Strength 16.8 Acid-Base Reactions of Salts 16.9 Practical Acid-Base Chemistry 16.10 Lewis Acid and Bases

  3. Types of Reactions a) Precipitation Reactions. Reactions of ionic compounds or salts b) Acid/base Reactions. Reactions of acids and bases c) Redox Reactions. reactions of oxidizing & reducing agents

  4. What are Acids &Bases? Definition? a) Arrhenius b) Bronsted-Lowry c) Lewis

  5. Arrhenius Definitions • Arrhenius, Svante August (1859-1927), Swedish chemist, 1903 Nobel Prize in chemistry • Acid Anything that produces hydrogen ions in a water solution. • HCl (aq) H+ ( aq) + Cl- (aq) • Base Anything that producshydroxide ions in a water solution. • NaOH (aq) Na+ (aq) + OH- (aq) • Arrhenius definitions are limited proton acids and hydroxide bases to aqueous solutions.

  6. Brønsted-Lowry definitions • Expands the Arrhenius definitions to include many bases other than hydroxides and gas phase reactions • Acid Proton donor • Base Proton acceptor • This definition explains how substances like ammonia can act as bases. • Eg. HCl(g) + NH3(g) ------> NH4Cl(s) • HCl (acid), NH3 (base). NH3(g) + H2O(l) NH4+ + OH-

  7. Lewis Definition G.N. Lewis was successful in including acid and bases without proton or hydroxyl ions. Lewis Acid: A substance that accepts an electron pair. Lewis base: A substance that donates an electron pair. E.g. BF3(g) + :NH3(g) F3B:NH3(s) the base donates a pair of electrons to the acid forming a coordinate covalent bond common to coordination compounds. Lewis acids/bases will be discussed later in detail

  8. Dissociation Strong Acids: HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq) H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) Dissociation Equilibrium Weak Acid/base: H2O(l) + H2O(l) H3+O(aq) + OH-(aq) This dissociation is called autoionization of water. HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2-(aq) NH3 (aq) + H2O(l) NH4+ + OH-(aq) Equilibrium constants: Ka, Kb and Kw

  9. Brønsted-Lowry Definitions Conjugate acid-base pairs. Acids and bases that are related by loss or gain of H+ as H3O+ and H2O. Examples.Acid Base H3O + H2O HC2H3O2C2H3O2- NH4+ NH3 H2SO4 HSO4- HSO4- SO42-

  10. Bronsted acid/conjugate base and base/conjugate acid pairs inacid/base equilibria HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq) HCl(aq): acid H2O(l): base H3+O(aq): conjugate acid Cl-(aq): conjugate base H2O/ H3+O: base/conjugate acid pair HCl/Cl-: acid/conjugate base pair

  11. Select acid, base, acid/conjugate base pair,base/conjugate acid pair H2SO4(aq) + H2O(l) H 3+O(aq) + HSO4-(aq) acid base conjugate acid conjugate base base/conjugate acid pair acid/conjugate base pair

  12. Binary acids: HCl, HBr, HI, H2S More than two elements: HCN Oxyacid: HNO3, H2SO4, H3PO4 Polyprotic acids: H2SO4, H3PO4 Organic acids: R-COOH, R= CH3-, CH3CH2- Acidic oxides: SO3, NO2, CO2, Basic oxides: Na2O, CaO Amine: NH3. R-NH2, R= CH3-, CH3CH2- : primary R2-NH : secondary, R3-N: tertiary Lewis acids & bases: BF3 andNH3 Types of Acids and Bases

  13. Strong Acid vs. Weak Acids Strong acid completely ionized Hydrioidic HI Ka ~ 1011 pKa = -11 HydrobromicHBr Ka ~ 109pKa = -9 Perchloric HClO4 Ka ~ 107 pKa = -7 HyrdrochloricHCl Ka ~ 107pKa = -7 Chloric HClO3 Ka ~ 103pKa = -3 Sulfuric H2SO4 Ka ~ 102pKa = -2 Nitric HNO3 Ka ~ 20 pKa = -1.3 Weak acid partially ionized Hydrofluoric acid HF Ka = 6.6x10-4pKa = 3.18 Formic acid HCOOH Ka = 1.77x10-4pKa = 3.75 Acetic acid CH3COOH Ka = 1.76x10-5 pKa = 4.75 Nitrous acid HNO2 Ka = 4.6x10-4 pKa = 3.34 Acetyl Salicylic acid C9H8O4 Ka = 3x10-4pKa = 3.52 Hydrocyanic acid HCN Ka = 6.17x10-10pKa = 9.21

  14. Strong Base vs. Weak Base Strong Base completely ionized Lithium hydroxide LiOH Sodium hydroxide NaOH Potassium hydroxide KOH Kb~ 102-103 Rubidium hydroxide RbOH Cesium hydroxide CsOH Boarder-line Bases Magnesium hydroxide Mg(OH)2 Calcium hydroxide Ca(OH)2 Strotium hydroxide Sr(OH)2 Kb~ 0.01 to0.1 Barium hydroxide Ba(OH)2 Weak Base partially ionized Ammonia NH3 Kb=1.79x10-5pKb = 4.74 Ethyl amine CH3CH2NH2 Kb=5.6x10-4pKb = 3.25

  15. Acid and Base Strength • Strong acids Ionize completely in water. HCl, HBr, HI, HClO3, • HNO3, HClO4, H2SO4. • Weak acids Partially ionize in water. • Most acids are weak. • Strong bases Ionize completely in water. Strong bases are metal hydroxides - NaOH, KOH • Weak bases Partially ionize in water.

  16. Common Acids and Bases Acids Formula Molarity* nitric HNO3 16 hydrochloric HCl 12 sulfuric H2SO4 18 acetic HC2H3O2 18 Bases ammonia NH3(aq) 15 sodium hydroxide NaOH solid *undiluted.

  17. Autoionization of Water • Autoionization When water molecules react with one another to form ions. • Acids and bases alter the dissociation equilibrium of water based on Le Chaterlier’s principle • Kw = [ H3O+ ] [ OH- ] • = 1.0 x 10-14 at 25oC • Note: [H2O] is constant and is included in Kw. H2O(l) + H2O(l) H3O+(aq) + OH-(aq) (10-7M) (10-7M) ion product of water

  18. pH and other “p” scales Substance pH 1 M HCl0.0 Gastric juices 1.0 - 3.0 Lemon juice 2.2 - 2.4 Classic Coke 2.5 Coffee 5.0 Pure Water 7.0 Blood 7.35 - 7.45 Milk of Magnesia 10.5 Household ammonia 12.0 1M NaOH 14.0 • We need to measure and use acids and bases over a very large concentration range. • pH and pOH are systems to keep track of these very large ranges. • pH = -log[H3O+] • pOH = -log[OH-] • pH + pOH = 14

  19. pH scale A logarithmic scale used to keep track of the large changes in [H+]. 0 7 14 10-14 M 10-7 M 10-14 M Very Neutral Very acidic Basic When you add an acid to, the pH gets smaller. When you add a base to, the pH gets larger.

  20. pH of some common materials Substance pH 1 M HCl 0.0 Gastric juices 1.0 - 3.0 Lemon juice 2.2 - 2.4 Classic Coke 2.5 Coffee 5.0 Pure Water 7.0 Blood 7.35 - 7.45 Milk of Magnesia 10.5 Household ammonia 12.0 1M NaOH 14.0

  21. pH of Aqueous Solutions

  22. What is pH? Kw = [H3+O][OH-] = 1 x 10-14 [H3+O][OH-] = 10-7 x 10-7 Extreme cases: Basic medium [H3+O][OH-] = 10-14 x 100 Acidic medium [H3+O][OH-] = 100 x 10-14 pH value is -log[H+] spans only 0-14 in water.

  23. pH, pKw and pOH The relation of pH, Kw and pOH Kw = [H+][OH-] log Kw = log [H+] + log [OH-] -log Kw= -log [H+] -log [OH-] ; previous equation multiplied by -1 pKw = pH + pOH; pKw = 14 since Kw =1 x 10-14 14 = pH + pOH pH = 14 - pOH pOH = 14 - pH

  24. pH and pOH calculations of acid and base solutions a) Strong acids/bases dissociation is complete for strong acid such as HNO3 or base NaOH [H+] is calculated from molarity (M) of the solution b) weak acids/bases needs Ka , Kb or percent(%)dissociation

  25. pH of Strong Acid/bases HNO3(aq) + H2O(l) H3+O(aq) + NO3-(aq) Therefore, the moles of H+ ions in the solution is equal to moles of HNO3 at the beginning. [HNO3] = [H+] = 0.2 mole/L pH = -log [H+] = -log(0.2) pH = 0.699 Substance pH 1 M HCl0.0 Gastric juices 1.0 - 3.0 Lemon juice 2.2 - 2.4 Classic Coke 2.5 Coffee 5.0 Pure Water 7.0 Blood 7.35 - 7.45 Milk of Magnesia 10.5 Household ammonia 12.0 1M NaOH 14.0

  26. pH of 0.5 M H2SO4 Solution H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq) [H3+O][HSO4-] H2SO4 ; Ka1 = ------------------- [H2SO4] [H3+O][SO42-] H2SO4 ; Ka2 = ------------------- ; Ka2 ignored [HSO4-]

  27. pH of 0.5 M H2SO4 Solution • H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) • the moles of H+ ions in the solution is equal to moles of H2SO4 at the beginning. • [H2SO4] = [H+] = 0.5 mole/L • pH = -log [H+] • pH = -log(0.5) • pH = 0.30

  28. 1.5 x 10-2 M NaOH. 1.5 x 10-2 M NaOH. NaOH is also a strong base dissociates completely in water. [NaOH] = [HO- ] = 1.5 x 10-2 mole/L pOH = -log[HO-]= -log(1.5 x 10-2) pOH = 1.82 As defined and derived previously: pKw= pH + pOH; pKw= 14 pH = pKw + pOH pH = 14 - pOH pH = 14 - 1.82 ; pH = 12.18

  29. Mixtures of Strong and Weak Acids • the presence of the strong acid retards the dissociation of the weak acid

  30. Measuring pH Arnold Beckman • inventor of the pH meter • father of electronic instrumentation

  31. Equilibrium, Constant, Ka & Kb Ka: Acid dissociation constant for a equilibrium reaction. Kb: Base dissociation constant for a equilibrium reaction. Acid: HA + H2O H3+O + A- Base: BOH + H2O B+ + OH- [H3+O][ A-] [B+ ][OH-] Ka = --------------- ; Kb = ----------------- [HA] [BOH]

  32. Acid Dissociation Constant HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq) [H3+O][Cl-] Ka= ----------------- [HCl] [H+][Cl-] Ka= ----------------- [HCl]

  33. Base Dissociation Constant NH3 + H2O NH4+ + OH- [NH4+][OH-] K = [NH3]

  34. Hydrated Metal Ions as Acids

  35. Ionization Constants for Acids

  36. Comparing Kw and Ka & Kb • Any compound with a Ka value greater than Kw of water will be a an acid in water. • Any compound with a Kb value greater than Kw of water will be a base in water.

  37. WEAKER/STRONGER Acids and Bases & Ka and Kb values • A larger value of Ka or Kb indicates an equilibrium favoring product side. • Acidity and basicity increase with increasing Ka or Kb. • pKa = - log Ka and pKb = - log Kb • Acidity and basicity decrease with increasing pKa or pKb.

  38. Which is weaker? • a. HNO2    ;  Ka= 4.0 x 10-4. • b. HOCl2    ;Ka= 1.2 x 10-2. • c. HOCl     ;  Ka= 3.5 x 10-8. • d. HCN      ;  Ka= 4.9 x 10-10.

  39. What is Ka1 andKa2? • H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) • HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq)

  40. Ka Examples H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq) [H3+O][HSO4-] H2SO4 ; Ka1 = ------------------- [H2SO4] [H3+O][SO42-] H2SO4 ; Ka2 = ------------------- [HSO4-]

  41. Ka Examples HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2-(aq) [H+][C2H3O2-] H C2H3O2; Ka= ------------------ [H C2H3O2] NH3 (aq) + H2O(l) NH4+ + OH-(aq) [NH4+][OH-] NH3; Kb= -------------- [ NH3]

  42. How do you calculate pH of weak acids/bases From % dissociation From Ka or Kb What is % dissociation Amount dissociated % Dissoc. = ------------------------- x 100 Initial amount

  43. How do you calculate % dissociation from Ka or Kb 1.00 M solution of HCN; Ka = 4.9 x 10-10 What is the % dissociation for the acid?

  44. 1.00 M solution of HCN; Ka = 4.9 x 10-10 1.00 M solution of HCN; Ka = 4.9 x 10-10 First write the dissociation equilibrium equation: HCN(aq) + H 2O(l) <===> H 3+O(aq) + CN-(aq) [HCN] [H+ ] [CN- ] Ini. Con. 1.00 M 0.0 M 0.00 M Cha. Con -x xx Eq. Con. 1.0 - x xx [H 3+O ][CN-] x2 Ka = ----------- = ---------------- [HCN] 1.0 - x

  45. 1.00 M solution of HCN; Ka = 4.9 x 10-10 1.0 - x ~ 1.00 since x is small x2 Ka = -----------; Ka = 4.9 x 10-10 = x2 1.0 x = 4.9 x 10-10 = 2.21 x 10 -5 Amount disso. 2.21 x 10 -5 ----------------- x 100 =- ------------- x 100 Ini. amount 1.00 % Diss. =2.21 x 10 -5 x 100 = 0.00221 %

  46. % Dissociation gives x (amount dissociated) need for pH calculation Amount dissociated % Dissoc. = ------------------------- x 100 Initial amount/con. x % Dissoc. = --------------------------- x 100 concentration

  47. Calculate the pH of a weak acid from % dissociation 1 M HF, 2.7% dissociated Notice the conversion of % dissociation to a fraction (x): 2.7/100=0.027) x=0.027

  48. Calculate the pH of a weak acid from % dissociation • HF(aq) + H 2O(l) <===> H 3+O(aq) + F-(aq) • [H+][F-] • Ka = ----------- • [HF] • [HF] [H+ ] [F- ] • Ini. Con. 1.00 M 0.0 M 0.00 M • Chg. Con -x xx • Eq.Con. 1.0-0.027 0.027 0.027 • pH = -log [H+] • pH = -log(0.027) • pH = 1.57

  49. [H3O+][Bz-] [HBz] Weak acid Equilibria Example Determine the pH of a 0.10 M benzoic acid solution at 25 oC if Ka = 6.28 x 10-5 HBz(aq) + H2O(l) H3O+(aq) + Bz-(aq) The first step is to write the equilibrium expression Ka =

  50. Weak acid Equilibria HBz H3O+ Bz- Initial conc., M 0.10 0.00 0.00 Change, DM -xxx Eq. Conc., M 0.10 - x xx [H3O+] = [Bz-] = x We’ll assume that [Bz-] is negligible compared to [HBz]. The contribution of H3O+ from water is also negligible.

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