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Chemistry 102(001) Fall 2012 PowerPoint Presentation
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Chemistry 102(001) Fall 2012

Chemistry 102(001) Fall 2012

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Chemistry 102(001) Fall 2012

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  1. Chemistry 102(001) Fall 2012 CTH 328 10:00-11:15 am 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 - 10:00am.. Exams: 10:00-11:15 am, CTH 328. September 27, 2012 (Test 1): Chapter 13 October 18, 2012 (Test 2): Chapter 14 &15 November 13, 2012 (Test 3):Chapter 16 &18 Optional Comprehensive Final Exam: November 15, 2012 : Chapters 13, 14, 15, 16, 17, and 18

  2. Chapter 16.AditionalAqueous Equilibria • 17.1 Buffer Solutions • 17.2 Acid-Base Titrations • 17.3 Acid Rain • 17.4 Solubility Equilibria and the Solubility Product Constant, Ksp • 17.5 Factors Affecting Solubility / Precipitation: Will It Occur?

  3. Hydrolysis Reaction of a basic anion or acidic cation with water is an ordinary Brønsted-Lowry acid-base reaction. CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq) NH4+(aq) + H2O(l) NH3(aq) + H3O+(aq) This type of reaction is given a special name. Hydrolysis The reaction of an anion with water to produce the conjugate acid and OH-. The reaction of a cation with water to produce the conjugate base and H3O+.

  4. Acid-Base Properties of Typical Ions

  5. What salt solutions would be acidic, basic and neutral? 1) strong acid + strong base = neutral 2) weak acid + strong base = basic 3) strong acid + weak base = acidic • weak acid + weak base = neutral, basic or an acidicsolution depending on the relative strengths of the acid and the base.

  6. What pH? Neutral, basic or acidic? • a)NaCl • neutral • b) NaC2H3O2 • basic • c) NaHSO4 • acidic • d) NH4Cl • acidic

  7. 1) If the following substance is dissolved in pure water, will the solution be acidic, neutral, or basic? • a) Solid sodium carbonate-(Na2CO3): • b) Sodium chloride- (NaCl): • c) Sodium acetate- (NaC2H3O2): • d) Ammonium sulfate-((NH4)2SO4):

  8. How do you calculate pH of a salt solution? • Find out the pH, acidic or basic? • If acidic it should be a salt of weak base • If basic it should be a salt of weak acid • if acidic calculate Ka from Ka= Kw/Kb • if basic calculate Kb from Kb= Kw/Ka • Do a calculation similar to pH of a weak acid or base

  9. What is the pH of 0.5 M NH4Cl salt solution? (NH 3; Kb = 1.8 x 10-5) • Find out the pH, acidic • if acidic calculate Ka from Ka= Kw/Kb • Ka= Kw/Kb = 1 x 10-14 /1.8 x 10-5) • Ka= 5.56. X 10-10 • Do a calculation similar to pH of a weak acid

  10. Continued NH4+ + H2O H 3+O + NH3 [NH4+] [H3+O ] [NH3 ] Ini. Con. 0.5 M 0.0 M 0.00 M Change -x xx Eq. Con. 0.5 - x x x [H 3+O ] [NH3 ] Ka(NH4+) = -------------------- = [NH 4+] x2 ---------------- ; appro.:0.5 - x . 0.5 (0.5 - x)

  11. Continued x2 Ka(NH4+) = ----------- = 5.56 x 10 -10 0. 5 x2 = 5.56 x 10 -10 x 0.5 = 2.78 x 10 -10 x= 2.78 x 10 -10 = 1.66 x 10-5 [H+ ] = x = 1.66 x 10-5 M pH = -log [H+ ] = - log 1.66 x 10-5 pH = 4.77 pH of 0.5 M NH4Cl solution is 4.77 (acidic)

  12. 2) What is the pH of a 0.05 M aqueous NH4Cl solution? (Kb (NH3) = 1.8 x 10-5) • a) equilibrium reaction for the hydrolysis of salt: • b) Ka for the conjugate acid NH4+: • c) ICE set-up: • I:___________________________________________ • C:__________________________________________ • E___________________________________________ • d) Calculation of x, [H3O]+: • e) pH of the solution:

  13. 3) What is the pH of a 0.05 M aqueous NaC2H3O2 solution? (Ka (HC2H3O2) = 1.8 x 10-5) • a) equilibrium reaction for the hydrolysis of salt: • b) Kb for the conjugate base C2H3O2-: • c) ICE set-up: • I:___________________________________________ • C:__________________________________________ • E___________________________________________ • d) Calculation of x, [OH]-: • e) pOH and pH of the solution:

  14. Acid-Base Chemistryof Some Antacids

  15. 4) A 50.00-mL sample of 0.100 M KOH is being titrated with 0.100 M HNO3. Calculate the pH of the solution after 52.00 mL of HNO3 is added. • a) acid base reaction: • b) moles of KOH: c) moles of HNO3: • d) [H3O]+: • e) pH of the solution:

  16. Hydrolysis Reaction of a basic anion with water is an ordinary Brønsted-Lowry acid-base reaction. CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq) This type of reaction is given a special name. Hydrolysis The reaction of an anion with water to produce the conjugate acid and OH-. The reaction of a cation with water to produce the conjugate base and H3O+.

  17. 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)

  18. Buffers • Solutions that resist pH change when small amounts of acid or base are added. • Two types • Mixture of weak acid and its salt • Mixture of 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.

  19. Buffers [A-] [HA] [HA] [A-] 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-Hasselbalch equation. Easily derived from the Ka or Kb expression. Starting with an acid pH = pKa + log Starting with a base pH = 14 - ( pKb + log )

  20. Henderson-Hasselbalch Equation HA(aq) + H2O(l) H3O+(aq) + A-(aq) [H3O+] [A-] Ka = ---------------- [HA] [H3O+] = Ka ([HA]/[A-]) pH = pKa + log([A-]/[HA]) when the [A-] = [HA] pH = pKa

  21. Calcualtion of pH of BuffersHenderson HesselbachEquation [ACID] pH = pKa - log --------- [BASE] [BASE] pH = pKa + log --------- [ACID]

  22. 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)

  23. Buffer Capacity • Refers to the ability of the buffer to retard changes in pH when small amounts of acid or base are added • The ratio of [A-]/[HA] determines the pH of the buffer whereas the magnitude of [A-] and [HA] determine the buffer capacity

  24. Adding an Acid or Baseto a Buffer

  25. Buffer Systems

  26. 5) For the (buffer effect) of HC2H3O2/NaC2H3O2 • a) Acid dissociation reaction: • b) Salt hydrolysis reaction: • c) Common ions in both equilibria: • d) Which way salt hydrolysis equilibrium move adding H3O+: • e) Which way salt hydrolysis equilibrium move adding OH-:

  27. 6) Describe the (buffer effect) of NH3/NH4Cl • a) Buffer type: (weak acid or base)/soluble salt): • b) Base dissociation reaction: • c) Salt hydrolysis reaction: • d) common ions in both equilibria: • e) Which way salt hydrolysis equilibrium move adding H3O+: • f) Which way salt hydrolysis equilibrium move adding OH-:

  28. 7) What is the pH of a solution that is 0.2 M in acetic acid (Ka = 1.8 x 10-5) and 0.2 M in sodium acetate? • a) Is it a acid, base, salt or buffer solution? • b) Henderson-Hesselbalch equation: • c) pKa: d) • e) pH of the solution:

  29. Titrations ofAcids and Bases • Titration • Analyte • Titrant analyte + titrant=> products

  30. Indicators Acid-base indicators are highly colored weak acids or bases. HIn In- + 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)

  31. Acid-Base Indicator HIn + H2O H3O+ + In- acid base color color [H3O+][In-] Ka = [HIn] They may have more than one color transition. Example.Thymol blue Red - Yellow - Blue Weak acid that changes color with changes in pH

  32. What is an Indicator? • Indicator is an weak acid with different Ka, colors to the acid and its conjugate base. E.g. phenolphthalein • Hin H+ + In- • colorless pink • Acidic colorless • Basic pink

  33. 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

  34. pH and Color of Indicators

  35. Red Cabbage as Indicator

  36. Indicator examples Acid-base indicators are weak acids that undergo a color change at a known pH. pH phenolphthalein

  37. 8) If 50 ml of a 0.01 M HCl solution is titrated with a 0.01 M NaOH solution, what will be the concentration of salt (NaCl) the pH at the endpoint? • a) NaCl solution acidic, basic or neutral? • b) Concentration of [NaCl]: • c) pH of the solution? • d) Suitable indicator for the titration:

  38. Titration Apparatus Burette delivering base to a flask containing an acid. The pink color in the flask is due to the phenolphthalein indicator.

  39. Endpoint vs. Equivalence Point Endpoint point where there is a physical change, such as color change, with theindicator Equivalence Point # moles titrant = # moles analyte #molestitrant=(V  M)titrant #molesanalyte=(V  M)analyte

  40. 9) If 50 mL of a 0.01 M HCl solution is titrated with a 0.01 M NH3 (Kb = 1.8 x 10-5) solution, what will be • a) The initial pH (0.01 M NH3): • b) Concentration of NH4Cl at the endpoint: • c) pH at the endpoint: • d) Suitable indicator for the titration:

  41. 10) If 50 ml of a 0.01 M HC2H3O2 solution is titrated with a 0.01 M NaOH solution, what will be the • a) Molarity of NaC2H3O2at the endpoint: • b) The pH at the endpoint: • c) What indicator would be most suitable for this titration:

  42. Polyprotic Acids

  43. Organic or Carboxylic Acids

  44. Organic or Carboxylic Acids FCH2CO2H (strongest acid) > ClCH2CO2H > BrCH2CO2H (weakest acid). Acid Ka pKa HCOOH (formic acid) 1.78 X 10-43 0.75 CH3COOH (acetic acid) 1.74 X 10-54 0.76 CH3CH2COOH (propanoic acid)1.38 x 10-5 4.86

  45. Acid-Base in the Kitchen vinegar - acetic acid lemon juice (citrus juice) - citric acid baking soda - NaHCO3 milk - lactic acid baking powder - H2PO4- & HCO3-

  46. Household Cleaners

  47. Dishwashing Detergent

  48. Acid-Base Indicator Behavior acid color shows when [H3O+][In-] 1 =  [H3O+] = Ka [HIn] 10 [In-] 1 £ [HIn] 10 base color shows when [H3O+][In-] = 10[H3O+] = Ka [HIn] [In-] 1 ³ [HIn] 10

  49. Indicator pH Range acid color shows when pH + 1 = pKa and base color shows when pH - 1 = pKa Color change range is pKa = pH  1 or pH = pKa 1

  50. Titration curves Acid-base titration curve A plot of the pH against the amount of acid or base added during a titration. Plots of this type are useful for visualizing a titration. It also can be used to show where an indicator undergoes its color change.