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Chemistry 281(01) Winter 2014

Chemistry 281(01) Winter 2014. CTH 277 10:00-11:15 am Instructor: Dr. Upali Siriwardane E-mail :  upali@latech.edu Office:  311 Carson Taylor Hall ; Phone: 318-257-4941; Office Hours:  MTW 8:00 am - 10:00 am; TR 8:30 - 9:30 am & 1:00-2:00 pm.

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Chemistry 281(01) Winter 2014

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  1. Chemistry 281(01) Winter 2014 CTH 27710:00-11:15 am Instructor: Dr. UpaliSiriwardane E-mail:  upali@latech.edu Office:  311 Carson Taylor Hall ; Phone: 318-257-4941; Office Hours:  MTW 8:00 am - 10:00 am; TR 8:30 - 9:30 am & 1:00-2:00 pm. January 16, 2014 Test 1 (Chapters 1&,2), February 6, 2014 Test 2 (Chapters 3 &4) February 25, 2014, Test 3 (Chapters 4 & 5), Comprehensive Final Make Up Exam: February 27, 2012 9:30-10:45 AM, CTH 311.

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

  3. Arrhenius definitions • Acid Anything that produces hydrogen ions in a water solution. • HCl(aq) H+ + Cl- • Base Anything that produces hydroxide ions in a water solution. • NaOH(aq) Na+ + OH- • Arrhenius definitions are limited to aqueous solutions. • Acid base reactions: • HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

  4. Brønsted-Lowry definitions • Expands the Arrhenius definitions • 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-

  5. Proton in water

  6. Dissociation Equilibrium HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq) H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) 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)

  7. Bronsted conjugate acid/base pairs in 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 CHEM 281 Winter 2014

  8. 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 HC2H3O2 C2H3O2- NH4+ NH3 H2SO4 HSO4- HSO4- SO42-

  9. 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

  10. Types of Acids and Bases Binary acids Oxyacid Organic acids Acidic oxides Basic oxides Amine Polyprotic acids

  11. Binary Acids Compounds containing acidic protons bonded to a more electronegative atom. e.g. HF, HCl, HBr, HI, H2S The acidity of the haloacid (HX; X = Cl, Br, I, F) Series increase in the following order: HF < HCl < HBr < HI

  12. Oxyacids Compounds containing acidic - OH groups in the molecule. Acidity of H2SO4 is greater than H2SO3 because of the extra O (oxygens) The order of acidity of oxyacids from the a halogen (Cl, Br, or I) shows a similar trend. HClO4 > HClO3 > HClO2 >HClO perchloricchloricchlorushyphochlorus

  13. Aqua Acids Acidic proton is on a water molecule coordinated to a central metal ion [Fe(OH2)6]3+,Al(OH2)63+, Si(OH)4 Acidity increase with charge Acidity increase as metal become smaller

  14. Anhydrous oxides The Lux/Flood Definition Covers things which would become acids or bases if dissolved in water. Acidic Oxides These are usually oxides of non-metallic elements such as P, S and N. E.g. NO2, SO2, SO3, CO2 They produce oxyacids when dissolved in water

  15. Basic Oxides Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water. e.g. CaO + H2O --> Ca(OH)2

  16. Protic Acids Monoprotic Acids: The form protic refers to acidity or protons. Monoprotic acids have only one acidic proton. e.g. HCl. Polyprotic Acids: They have more than one acidic proton. e.g. H2SO4 - diprotic acid H3PO4 - triprotic acid.

  17. Amines Class of organic bases derived from ammonia NH3 by replacing hydrogen by organic groups. They are defined as bases similar to NH3 by Bronsted or Lewis acid/base definitions.

  18. What acid base concepts (Arrhenius/Bronsted/Lewis) would best describe the following reactions: a) HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l) b)HCl(g) + NH3(g) ---> NH4Cl(s) c)BF3(g) + NH3(g) ---> F3B:NH3(s) d)Zn(OH)2(s) + 2OH-(aq) ---> [Zn(OH)4]2- (aq)

  19. 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.

  20. Acids and bases Acidic Basic • Citrus fruits Baking soda • Aspirin Detergents • Coca Cola Ammonia cleaners • Vinegar Tums and Rolaids • Vitamin C Soap

  21. 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]

  22. What is Ka HCl(aq) + H2O(l) <===> H3+O(aq) + Cl-(aq)

  23. E.g. Ka HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq) [H3+O][Cl-] Ka= ----------------- [HCl] [H+][Cl-] Ka= ----------------- [HCl]

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

  25. What is Kb NH3 (aq) + H2O(l) NH4+ + OH-(aq)

  26. H2SO4 Dissociation E.g. 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-]

  27. Ka and Kb E.g. 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]

  28. Acidity/Basicity of HA and F-

  29. 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.

  30. 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.

  31. Autoionization of water • Autoionization When water molecules react with one another to form ions. • H2O(l) + H2O(l) H3O+(aq) + OH-(aq) • (10-7M) (10-7M) • Kw = [ H3O+ ] [ OH- ] • = 1.0 x 10-14 at 25oC • Note: [H2O] is constant and is included in Kw. ion product of water

  32. What is Kw? H2O(l) + H2O(l) H3+O(aq) + OH-(aq) This dissociation is called autoionization of water. Autoionization of water: Kw = [H3+O][OH-] Kwis called ionic product of water Kw = 1 x 10-14

  33. Why is water important for acid/base equilibria? Water is the medium/solvent for acids and bases. Acids and bases alter the dissociation equilibrium of water based on Le Chaterlier’s principle H2O(l) + H2O(l) H3+O(aq) + OH-(aq)

  34. 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.

  35. pH and other “p” scales • 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

  36. pH scale A logarithmic scale used to keep track of the large changes in [H+]. 14 7 0 10-14 M 10-7 M 1 M Very Neutral Very Basic Acidic When you add an acid, the pH gets smaller. When you add a base, the pH gets larger.

  37. 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

  38. 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.

  39. 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

  40. 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.

  41. 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

  42. Titration curves Overtitration Indicator Transition Equivalence Point pH Buffer region % titration or ml titrant

  43. Indicators Acid-base indicators are highly colored weak acids or bases. HInIn- + H+ color 1color 2 They may have more than one color transition. Example.Thymol blue Red - Yellow - Blue One of the forms may be colorless - phenolphthalein (colorless to pink)

  44. Selection of an indicator for a titration a) strong acid/strong base b) weak acid/strong base c) strong acid/weak base d) weak acid/weak base Calculate the pH of the solution at he equivalence point or end point

  45. Common Ion Effect Weak acid and salt solutions E.g. HC2H3O2 and NaC2H3O2 Weak base and salt solutions E.g. NH3 and NH4Cl. H2O + C2H3O2- <==> OH- + HC2H3O2 (common ion) H2O + NH4+ <==> H3+O + NH3 (common ion)

  46. Buffers • Solutions that resist pH change when small amounts of acid or base are added. • Two types • weak acid and its salt • weak base and its salt • HA(aq) + H2O(l) H3O+(aq) + A-(aq) • Add OH- Add H3O+ • shift to right shift to left • Based on the common ion effect.

  47. [A-] [HA] [HA] [A-] Buffers The pH of a buffer does not depend on the absolute amount of the conjugate acid-base pair. It is based on the ratio of the two. Henderson-Hasselbalchequation. Easily derived from the Ka or Kb expression. Starting with an acid pH = pKa + log Starting with a base pH = 14 - ( pKb + log )

  48. Buffers and blood • Control of blood pH • Oxygen is transported primarily by hemoglobin in the red blood cells. • CO2 is transported both in plasma and the red blood cells. CO2 (aq) + H2O H2CO3(aq) The bicarbonate buffer is essential for controlling blood pH H+(aq) + HCO3-(aq)

  49. Main Group Acid/Bases

  50. Amphoteric Oxides

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