2. Acids, Bases, and Salts You should be able to  Understand the acid-base theories of Arrhenius, Brønsted-Lowry, and Lewis.  Identify strong acids and bases and calculate their pH’s.  Calculate the pH of a weak acid or base.  Calculate the concentration of a strong or weak acid or base from its pH.  Calculate the pH and ion concentration in a polyprotic acid.  Predict the pH of a salt from its formula and then calculate the pH of the salt.  Be familiar with titration curves and selection of an acid-base indicator.

3. pH scale 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 ACIDBASE NEUTRAL Each step on pH scale represents a factor of 10. pH 5 vs. pH 6 (10X more acidic) pH 3 vs. pH 5 (100X different) pH 8 vs. pH 13 (100,000X different) : measures acidity/basicity 10x 100x Søren Sorensen (1868 - 1939)

4. Acid Base pH = -log [H 1+ ] pH = 7 Acidic Basic Neutral [H + ] [OH - ] [H + ] = [OH - ]

5. Acid vs. Base Acid pH > 7 bitter taste does not react with metals pH < 7 sour taste react with metals Alike Different Related to H + (proton) concentration pH + pOH = 14 Affects pH and litmus paper Base Different Topic

6. Properties electrolytes turn litmus red sour taste react with metals to form H 2 gas slippery feel turn litmus blue bitter taste ChemASAP vinegar, milk, soda, apples, citrus fruits ammonia, lye, antacid, baking soda electrolytes

7. Common Acids and Bases Strong Acids (strong electrolytes) HClhydrochloric acid HNO 3 nitric acid HClO 4 perchloric acid H 2 SO 4 sulfuric acid Weak Acids (weak electrolytes) CH 3 COOHacetic acid H 2 CO 3 carbonic Strong Bases (strong electrolytes) NaOH sodium hydroxide KOH potassium hydroxide Ca(OH) 2 calcium hydroxide Weak Base (weak electrolyte) NH 3 ammonia Kotz, Purcell, Chemistry & Chemical Reactivity 1991, page 145 Weak Base (weak electrolyte) NH 4 OH ammonia NH 3 + H 2 O  NH 4 OH

8. Salt Acid + Base  Salt + Water Orange juice + milk  bad taste Evergreen shrub + concrete  dead bush Under a pine tree + fertilizer  white powder HCl + NaOH  NaCl + HOH salt water

9. Acid-Base Neutralization 1+ 1- ++ Hydronium ion Hydroxide ionWater H3O+H3O+ OH - H2OH2O Water H2OH2O Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 584

10. Formation of Sulfuric Acid Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 302 SO 2 (g) + H 2 O(l) H 2 SO 3 (aq) 2SO 2 (g) + O 2 (g) 2SO 3 (g) SO 3 (g) + H 2 O(l) H 2 SO 4 (aq) SO 2 (g) + H 2 O 2 (l) H 2 SO 4 (aq) Catalyzed by atmospheric dust Sulfuric acid + +

11. Coal Burning Power Plant

12. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

13. CO 2 (g) H 2 O (l) H 2 CO 3 (aq) Carbon dioxide Carbonic acid Water Weak acid Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

14. Common Acids Sulfuric AcidH 2 SO 4 Nitric AcidHNO 3 Phosphoric AcidH 3 PO 4 Hydrochloric AcidHCl Acetic Acid CH 3 COOH Carbonic Acid H 2 CO 3 Battery acid Used to make fertilizers and explosives Food flavoring Stomach acid Vinegar Carbonated water

15. Common Acids Formula FormulaName of AcidName of Negative Ion of Salt HFhydrofluoricfluoride HBrhydrobromicbromide HIhydroiodiciodide HClhydrochloricchloride HClOhypochloroushypochlorite HClO 2 chlorouschlorite HClO 3 chloricchlorate HClO 4 perchloricperchlorate H 2 Shydrosulfuricsulfide H 2 SO 3 sulfuroussulfite H 2 SO 4 sulfuricsulfate HNO 2 nitrousnitrite HNO 3 nitricnitrate H 2 CO 3 carboniccarbonate H 3 PO 3 phosphorousphosphite H 3 PO 4 phosphoricphosphate

16. Formation of Hydronium Ions 1+ hydronium ion H3O+H3O+ + hydrogen ion H+H+ water H2OH2O 1+ (a proton) 1+

17. Sulfuric Acid, H 2 SO 4 Sulfuric acid is the most commonly produced industrial chemical in the world. Uses: petroleum refining, metallurgy, manufacture of fertilizer, many industrial processes: metals, paper, paint, dyes, detergents Sulfuric acid is used in automobile batteries. H 2 SO 4 “oil of vitriol”

18. Nitric Acid, HNO 3 Nitric acid stains proteins yellow (like your skin). Uses: make explosives, fertilizers, rubber, plastics, dyes, and pharmaceuticals. HNO 3 “aqua fortis” O O O N H

19. Hydrochloric Acid, HCl The stomach produces HCl to aid in the digestion of food. Uses: For ‘pickling’ iron and steel. Pickling is the immersion of metals in acid solution to remove surface impurities. A dilute solution of HCl is called muriatic acid (available in many hardware stores). Muriatic acid is commonly used to adjust pH in swimming pools and in the cleaning of masonry. HCl(g) + H 2 O(l) HCl(aq) hydrogen chloride water hydrochloric acid

20. Common Bases Sodium hydroxideNaOHlye or caustic soda Potassium hydroxideKOHlye or caustic potash Magnesium hydroxideMg(OH) 2 milk of magnesia Calcium hydroxideCa(OH) 2 slaked lime Ammonia waterNH 3 H 2 Ohousehold ammonia Name Formula Common Name. NH 4 OH NH 4 1+ + OH 1- ammonium hydroxide hydroxide ion OH 1-

21. Relative Strengths of Acids and Bases perchloricHClO 4 hydrogen chlorideHCl nitricHNO 3 sulfuricH 2 SO 4 hydronium ionH 3 O + hydrogen sulfate ionHSO 4 - phosphoricH 3 PO 4 aceticHC 2 H 3 O 2 carbonicH 2 CO 3 hydrogen sulfideH 2 S ammonium ionNH 4 + hydrogen carbonate ionHCO 3 - waterH 2 O ammoniaNH 3 hydrogenH 2 Decreasing Acid Strength perchlorate ionClO 4 - chloride ionCl - nitrate ionNO 3 - hydrogen sulfate ionHSO 4 - waterH 2 O sulfate ionSO 4 2- dihydrogen phosphate ionH 2 PO 4 - acetate ionC 2 H 3 O 2 - hydrogen carbonate ionHCO 3 - hydro sulfide ionHS - ammoniaNH 3 carbonate ionCO 3 2- hydroxide ionOH - amide ionNH 2 - hydride ionH - Decreasing Base Strength Acid Formula Conjugate baseFormula Metcalfe, Williams, Catska, Modern Chemistry  1966, page 229 acid conjugate base + H +

22. Binary Hydrogen Compounds of Nonmetals When Dissolved in Water (These compounds are commonly called acids.) The prefix hydro- is used to represent hydrogen, followed by the name of the nonmetal with its ending replaced by the suffix –ic and the word acid added. Examples: *HCl HBr *The name of this compound would be hydrogen chloride if it was NOT dissolved in water. Hydrochloric acid Hydrobromic acid

23. Naming Simple Chemical Compounds Ionic (metal and nonmetal)Covalent (2 nonmetals) Metal Forms only one positive ion Forms more than one positive ion Nonmetal Use the name of element Use element name followed by a Roman numeral to show the charge First nonmetal Second nonmetal Before element name use a prefix to match subscript Use a prefix before element name and end with ide Single Negative Ion Polyatomic Ion Use the name of the element, but end with ide Use the name of polyatomic ion (ate or Ite)

24. Naming Ternary Compounds from Oxyacids The following table lists the most common families of oxy acids. one more oxygen atom most “common” one less oxygen two less oxygen HClO 4 perchloric acid HClO 3 chloric acid HClO 2 chlorous acid HClO hypochlorous acid H 2 SO 4 sulfuric acid H 2 SO 3 sulfurous acid H 3 PO 4 phosphoric acid H 3 PO 3 phosphorous acid H 3 PO 2 hypophosphorous acid HNO 3 nitric acid HNO 2 nitrous acid (HNO) 2 hyponitrous acid

25. Oxyacids  Oxysalts If you replace hydrogen with a metal, you have formed an oxysalt. A salt is a compound consisting of a metal and a non-metal. If the salt consists of a metal, a nonmetal, and oxygen it is called an oxysalt. NaClO 4, sodium perchlorate, is an oxysalt. HClO 4 perchloric acid HClO 3 chloric acid HClO 2 chlorous acid HClO hypochlorous acid NaClO 4 sodium perchlorate NaClO 3 sodium chlorate NaClO 2 sodium chlorite NaClO sodium hypochlorite OXYACID OXYSALT

26. ACID SALT per stem ic changes to per stem ate stem ic changes to stem ate stem ous changes to stem ite hyper stem ous changes to hypo stem ite HClO 3 + Na 1+ NaClO 3 + H 1+ acid cation salt

27. Arrhenius Acids and Bases Acids release hydrogen ions in water. Bases release hydroxide ions in water. An acid is a substance that produces hydronium ions, H 3 O +, when dissolved in water. Lewis Definitions A Lewis acid is a substance than can accept (and share) an electron pair. A Lewis base is a substance than can donate (and share) an electron pair. Lewis Acid Brønsted-Lowry Definitions A Brønsted-Lowry acid is a proton donor; it donates a hydrogen ion, H +. A Brønsted-Lowry base is a proton acceptor; it accepts a hydrogen ion, H +. Brønsted-Lowry Arrhenius acids Acid Definitions

28. Acid – Base Systems TypeAcidBase ArrheniusH + or H 3 O + producer OH - producer Brønsted- Lowry Proton (H + ) donor Proton (H + ) acceptor LewisElectron-pair acceptor Electron-pair donor

29. Arrhenius Bases and Their Properties According to the definition of Arrhenius a: Base Base - "a substance whose water solution yields... Are NaOH and NH 3 considered to be Arrhenius bases? 1) Bases are electrolytes Dissociation equation for NH 3 NH 3 (g) + H 2 O(l) NH 4 1+ (aq) + OH 1- (aq) Dissociation equation for NaOH NaOH(s) Na 1+ (aq) + OH 1- (aq) indicators 2) Bases cause indicators to turn a characteristic color 3) Bases neutralize acids NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(l) YES 4) Water solutions of bases tasted bitter and feel slippery. hydroxide ions (OH - ) as the only negative ions."

30. Neutralization Neutralization Neutralization is a chemical reaction between an acid and a base to produce a salt (an ionic compound) and water. NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(l) baseacidsaltwater Some neutralization reactions: H 2 SO 4 (aq) + NaOH(aq) Na 2 SO 4 +HOH sulfuric acidsodium hydroxidesodium sulfatewater HC 2 H 3 O 2 (aq) + Ca(OH) 2 (aq) Ca(C 2 H 3 O 2 ) 2 +HOH acetic acidcalcium hydroxidecalcium acetatewater 2 2 2 2

31. Neutralization ACID + BASE  SALT + WATER HCl + NaOH  NaCl + H 2 O HC 2 H 3 O 2 + NaOH  NaC 2 H 3 O 2 + H 2 O Salts can be neutral, acidic, or basic. Neutralization does not mean pH = 7. weak strong neutral basic Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

32. ACID + BASE  SALT + WATER HCl + NaOH  NaCl + H 2 O HC 2 H 3 O 2 + NaOH  NaC 2 H 3 O 2 + H 2 O Salts can be neutral, acidic, or basic. Neutralization does not mean pH = 7. weak strong neutral basic Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

33. Salt Formation NaOH HCl strong base strong acid salt of a strong base and a strong acid NaCl NaOH HC 2 H 3 O 2 strong base weak acid salt of a strong base and a weak acid NaC2H3O2C2H3O2 Note: that in each case H-OH (water) is formed NaOH + HCl  NaCl + H 2 O NaOH + HC 2 H 3 O 2  NaC 2 H 3 O 2 + H 2 O

34. Salt Formation NH 3 H 2 SO 4 weak base strong acid salt of a weak base and a strong acid (NH 4 ) 2 SO 4 NH 3 HC 2 H 3 O 2 weak base weak acid salt of a weak base and a weak acid NH 4 C2H3O2C2H3O2 Note: that in each case H-OH (water) is also formed NH 4 OH H 2 SO 4 NH 4 OH + H 2 SO 4  (NH 4 ) 2 SO 4 + H 2 O NH 4 OH + HC 2 H 3 O 2  NH 4 C 2 H 3 O 2 + H 2 O NH 4 OH

35. NH 3 H 2 SO 4 weak base strong acid salt of a weak base and a strong acid (NH 4 ) 2 SO 4 NH 4 OH H 2 SO 4 ammonium ion NH 4 + hydroxide ion OH - 1+ 1- NH 4 + OH - 1+ 1- sulfuric acid (NH 4 ) 2 SO 4 HOH 1+ HOH 1+ sulfate ion 2 NH 4 OH + H 2 SO 4 (NH 4 ) 2 SO 4 + 2 HOH water ammonium sulfate 2- H 2 SO 4 2 NH 4 OH + H 2 SO 4 (NH 4 ) 2 SO 4 + 2 H 2 O

36. phosphoric acid ammonium hydroxide ammonium phosphate Reactions that produce salt acid + base salt +water H 3 PO 4 NH 4 OH (NH 4 ) 3 PO 4 H2OH2O nitric acid magnesium hydroxide magnesium nitrate HNO 3 Mg(OH) 2 Mg(NO 3 ) 2 H2OH2O carbonic acid potassium hydroxide potassium carbonate H 2 CO 3 KOH K 2 CO 3 H2OH2O acetic acid aluminum hydroxide aluminum acetate HC 2 H 3 O 2 Al(OH) 3 Al(C 2 H 3 O 2 ) 3 H2OH2O perchloric acid barium hydroxide barium perchlorate HClO 4 Ba(OH) 2 Ba(ClO 4 ) 2 H2OH2O + + and yieldsand water

37. Brønsted-Lowry Acids and Bases 1+ + hydronium ion H3O+H3O+ 1- chloride ion Cl - (base) H2OH2O (acid) HCl ++ -- Acid = any substance that donates a proton. Base = any substance that accepts a proton.

38. Brønsted-Lowry Acids and Bases (acid) H2OH2O (base) NH 3 ++ -- 1+ + ammonium ion NH 4 + 1- hydroxide ion OH -

39. Definitions  Brønsted-Lowry HCl + H 2 O  Cl – + H 3 O + AcidsAcids are proton (H + ) donors. BasesBases are proton (H + ) acceptors. conjugate acid conjugate base baseacid Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

40. Definitions H 2 O + HNO 3  H 3 O + + NO 3 – CBCAAB Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem H H O HO O O N Base Acid

41. Definitions - can be an acid or a base.  Amphoteric - can be an acid or a base. NH 3 + H 2 O  NH 4 + + OH - CACBBA Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem H H O H N Base Acid H H

42. Definitions F - H 2 PO 4 - H2OH2O HF H 3 PO 4 H 3 O +  Give the conjugate base for each of the following: - an acid with more than one H +  Polyprotic - an acid with more than one H + Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

43. Definitions Br - HSO 4 - CO 3 2- HBr H 2 SO 4 HCO 3 -  Give the conjugate acid for each of the following: Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

44. Definitions  Lewis AcidsAcids are electron pair acceptors. BasesBases are electron pair donors. Lewis base Lewis acid Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

45. pH Scale Søren Sorensen (1868 - 1939)

46. pH Scale Acid Base 0 7 14 Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 515 [H + ] pH 10 -14 14 10 -13 13 10 -12 12 10 -11 11 10 -10 10 10 -9 9 10 -8 8 10 -7 7 10 -6 6 10 -5 5 10 -4 4 10 -3 3 10 -2 2 10 -1 1 10 0 0 1 M NaOH Ammonia (household cleaner) Blood Pure water Milk Vinegar Lemon juice Stomach acid 1 M HCl Acidic Neutral Basic

47. pH of Common Substances Timberlake, Chemistry 7 th Edition, page 335 1.0 M HCl 0 gastric juice 1.6 vinegar 2.8 carbonated beverage 3.0 orange 3.5 apple juice 3.8 tomato 4.2 lemon juice 2.2 coffee 5.0 bread 5.5 soil 5.5 potato 5.8 urine 6.0 milk 6.4 water (pure) 7.0 drinking water 7.2 blood 7.4 detergents 8.0 - 9.0 bile 8.0 seawater 8.5 milk of magnesia 10.5 ammonia 11.0 bleach 12.0 1.0 M NaOH (lye) 14.0 8 910 111214 13 34 5 6 2 1 70 acidic neutral basic [H + ] = [OH - ]

48. pH of Common Substance 14 1 x 10 -14 1 x 10 -0 0 13 1 x 10 -13 1 x 10 -1 1 12 1 x 10 -12 1 x 10 -2 2 11 1 x 10 -11 1 x 10 -3 3 10 1 x 10 -10 1 x 10 -4 4 9 1 x 10 -9 1 x 10 -5 5 8 1 x 10 -8 1 x 10 -6 6 6 1 x 10 -6 1 x 10 -8 8 5 1 x 10 -5 1 x 10 -9 9 4 1 x 10 -4 1 x 10 -10 10 3 1 x 10 -3 1 x 10 -11 11 2 1 x 10 -2 1 x 10 -12 12 1 1 x 10 -1 1 x 10 -13 13 0 1 x 10 0 1 x 10 -14 14 NaOH, 0.1 M Household bleach Household ammonia Lime water Milk of magnesia Borax Baking soda Egg white, seawater Human blood, tears Milk Saliva Rain Black coffee Banana Tomatoes Wine Cola, vinegar Lemon juice Gastric juice More basic More acidic pH [H 1+ ] [OH 1- ] pOH 7 1 x 10 -7 1 x 10 -7 7

49. Acid – Base Concentrations pH = 3 pH = 7 pH = 11 OH - H3O+H3O+ H3O+H3O+ H3O+H3O+ [H 3 O + ] = [OH - ] [H 3 O + ] > [OH - ] [H 3 O + ] < [OH - ] acidic solution neutral solution basic solution concentration (moles/L) 10 -14 10 -7 10 -1 Timberlake, Chemistry 7 th Edition, page 332

50. pH pH = -log [H 1+ ] Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 285

51. pH Calculations pH pOH [H 3 O + ] [OH - ] pH + pOH = 14 pH = -log[H 3 O + ] [H 3 O + ] = 10 -pH pOH = -log[OH - ] [OH - ] = 10 -pOH [H 3 O + ] [OH - ] = 1 x10 -14

52. pH = - log [H + ] pH = 4.6 pH = - log [H + ] 4.6 = - log [H + ] - 4.6 = log [H + ] Given: 2 nd log 10 x antilog multiply both sides by -1 substitute pH value in equation take antilog of both sides determine the [hydronium ion] choose proper equation [H + ] = 2.51x10 -5 M You can check your answer by working backwards. pH = - log [H + ] pH = - log [2.51x10 -5 M] pH = 4.6 Recall, [H + ] = [H 3 O + ]

53. Acid Dissociation monoprotic diprotic polyprotic HA(aq) H 1+ (aq) + A 1- (aq) 0.03 M pH = - log [H + ] pH = - log [0.03M] pH = 1.52 e.g. HCl, HNO 3 H 2 A(aq) 2 H 1+ (aq) + A 2- (aq) 0.3 M0.6 M0.3 M pH = - log [H + ] pH = - log [0.6M] pH = 0.22 e.g. H 2 SO 4 Given: pH = 2.1 find [H 3 PO 4 ] assume 100% dissociation e.g. H 3 PO 4 H 3 PO 4 (aq) 3 H 1+ (aq) + PO 4 3- (aq) ? Mx M pH = ?

54. Given: pH = 2.1 find [H 3 PO 4 ] assume 100% dissociation H 3 PO 4 (aq) 3 H 1+ (aq) + PO 4 3- (aq) X MX M0.00794 M Step 1) Write the dissociation of phosphoric acid Step 2) Calculate the [H + ] concentration pH = - log [H + ] 2.1 = - log [H + ] - 2.1 = log [H + ] 2 nd log - 2.1 = log [H + ] 2 nd log [H + ] = 10 -pH [H + ] = 10 -2.1 [H + ] = 0.00794 M [H + ] = 7.94 x10 -3 M 7.94 x10 -3 M Step 3) Calculate [H 3 PO 4 ] concentration Note: coefficients (1:3) for (H 3 PO 4 : H + ) 7.94 x10 -3 M 3 = 0.00265 M H 3 PO 4

55. How many grams of magnesium hydroxide are needed to add to 500 mL of H 2 O to yield a pH of 10.0? Step 1) Write out the dissociation of magnesium hydroxide Mg 2+ OH 1- Mg(OH) 2 Mg(OH) 2 (aq)Mg 2+ (aq) 2 OH 1- (aq)+ Step 2) Calculate the pOH pH + pOH = 14 10.0 + pOH = 14 pOH = 4.0 Step 3) Calculate the [OH 1- ] pOH = - log [OH 1- ] [OH 1- ] = 10 -OH [OH 1- ] = 1 x10 -4 M 1 x10 -4 M0.5 x10 -4 M5 x10 -5 M Step 4) Solve for moles of Mg(OH) 2 x = 2.5 x 10 -5 mol Mg(OH) 2 Step 5) Solve for grams of Mg(OH) 2 x g Mg(OH) 2 = 2.5 x 10 -5 mol Mg(OH) 2 1 mol Mg(OH) 2 = 0.00145 g Mg(OH) 2 58 g Mg(OH) 2

56. [ CaO ], shift [ CO 2 ], shift -- shift ; eggshells are thinner In a chicken…CaO + CO 2 CaCO 3 (eggshells) In summer, [ CO 2 ] in a chicken’s blood due to panting. How could we increase eggshell thickness in summer? -- give chickens carbonated water -- put CaO additives in chicken feed -- air condition the chicken house TOO much $$$ -- pump CO 2 gas into the chicken house would kill all the chickens! I wish I had sweat glands.

57. LeChatelier’s Principle N 2 + 3 H 2 2 NH 3 + heat Raising the temperature… …favors the endothermic reaction (the reverse reaction) in which the rise in temperature is counteracted by the absorption of heat. Increasing the pressure… …favors the forward reaction in which 4 mol of gas molecules is converted to 2 mol. Decreasing the concentration of NH 3 … …favors the forward reaction in order to replace the NH 3 that has been removed. Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532 Animation by Raymond Chang All rights reserved

58. Equilibrium Expression N 2 + 3 H 2 2 NH 3 + heat Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532 Haber Process

59. reversible reaction: H 2 SO 4 2 H 1+ + SO 4 2– Acid dissociation is a reversible reaction. Rate at which R  P Rate at which P  R = looks like nothing is happening, however… system is dynamic, NOT static equilibrium: Reactant  Product and P  R Reactant Product

60. Remove NH 3 ………………….. “ “ NH 3 ………………… “ “ H 2 ………………….. Add more N 2 ………………….. Le Chatelier’s principle N 2 (g) + 3 H 2 (g) 2 NH 3 (g) Le Chatelier’s principle: DisturbanceEquilibrium Shift no shift When a system at equilibrium is disturbed, it shifts to a new equilibrium that counteracts the disturbance. Add a catalyst………………… Increase pressure……………. Fritz Haber

61. shift to a new equilibrium: Then go inside… shift to a new equilibrium: Light-Darkening Eyeglasses AgCl + energy Ag o + Cl o “energy” Go outside… Sunlight more intense than inside light; GLASSES DARKEN (clear) (dark) “energy” GLASSES LIGHTEN

62. Maintaining Blood pH Acid entering the blood stream Carbon dioxide is exhaled HCO 3 1- + H + H 2 CO 3 H 2 O + CO 2 Bicarbonate ion circulates in the blood stream where it is in equilibrium with H + and OH -. In the lungs, bicarbonate ions combine with a hydrogen ion and lose a water molecule to form carbon dioxide, which is exhaled. Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291

63. Alkalosis If our breathing becomes too fast (hyperventilation)… Carbon dioxide is removed from the blood too quickly. This accelerates the rate of degradation of carbonic acid into carbon dioxide and water. The lower level of carbonic acid encourages the combination of hydrogen ions and bicarbonate ions to make more carbonic acid. The final result is a fall in blood H 1+ levels that raises blood pH which can result in over-excitability or death. Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291

64. Acidosis If breathing becomes too slow (hypoventilation)… …free up acid, pH of blood drops, with associated health risks such as depression of the central nervous system or death. The normal pH of blood is between 7.2 – 7.4. This pH is maintained by the bicarbonate ion and other buffers. Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291

65. H + A - H + A - HA A - H + A - H + A – H + A - H + A - H + A - HA H + A - H + A - H + A - H + HA HA HA HA HA H + A - HA HA HA HA H + A – HA H + A – HA HA H + A - H + A - H + A - H + A - HA A - H + A - H + A - H + A - H + A - H + A - HA H + A - H + A - H + A - A - H + A - H + A - H + A - H + A - H + H + A - H + A - H + A - HA H + A - A - H + A - H + A - H + A - H + A – H + A - H + A - H + A - H + A - H + A - H + A - H + A - H + A - H + A - HA A - H + A - H + A - H + A - H + HA HA H + A - HA HA HA HA HA HA HA HA H + A - H + A - HA HA HA HA HA HA HA H + A - HA HA HA HA HA HA H + A - HA HA H + A - HA HA HA HA HA HA HA HA H + A - HA HA H + A - HA HA HA HA HA HA HA H + A - HA HA HA DILUTECONCENTRATED STRONG WEAK STRONG ACIDS Dissociate nearly 100% HA H 1+ + A - WEAK ACIDS Dissociate very little HA H 1+ + A - Acids: Concentration vs. Strength

66. Comparison of Strong and Weak Acids Type of acid, HA Reversibility of reaction K a value Ions existing when acid, HA, dissociates in H 2 O Strong Not reversible K a value very large H + and A -, only. No HA present. WeakreversibleK a is smallH +, A -, and HA HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) The equilibrium expression for the reaction is K a = [H 3 O + ] [A - ] [HA] Note: H 3 O + = H +

67. Strong vs. Weak Acid Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 508

68. Concentrated vs. Dilute 0.3 M HCl 2.0 M HCl 12.0 M HCl 10.0 M CH 3 COOH Dilute, strong acid Concentrated, strong acid OR Dilute, strong, acid Concentrated, strong acid Concentrated, weak acid

69. Naming Acids _________ ide (chloride, Cl 1- ) _________ite (chlorite, ClO 2 - ) (hypochlorite, ClO - ) _________ ate (chlorate, ClO 3 - ) (perchlorate, ClO 4 - ) Hydro____ ic acid (hydrochloric acid, HCl) _________ic acid (chloric acid, HClO 3 ) (perchloric acid, HClO 4 ) ______ous acid (chlorous acid, HClO 2 ) (hypochlorous acid, HClO) AnionAcid add H + ions

70. [H 3 O + ] Equilibrium and pH Calculations HA + H 2 O H 3 O + + A - Weak acid HA + H 2 O H 3 O + + A - Strong acid acid-dissociation constant calculations K a = [A - ] [H 3 O + ] [HA] [HA] = [H 3 O + ] + pH 0 7 14 antilog (-pH) -log [H 3 O + ] [OH - ] - 1 x 10 -14 [OH - ] = 1 x 10 -14 [H 3 O + ] = Tocci, Viehland, Holt Chemistry Visualizing Matter 1996, page 525 HA H + + A - K w = [H 3 O + ][OH - ] 1 x 10 -14 = [H 3 O + ][OH - ]

71. Strengths of Conjugate Acid-Base Pairs strong medium weak very weak Acid strength increases HCl H 2 SO 4 HNO 3 H 3 O + HSO 4 - H 3 PO 4 HC 2 H 3 O 2 H 2 CO 3 H 2 S H 2 PO 4 - NH 4 + HCO 3 - HPO 4 2- H 2 O negligible very weak weak medium strong Base strength increases Cl - HSO 4 - NO 3 H 2 O SO 4 2- H 2 PO 4 - C 2 H 3 O 2 - HCO 3 - HS - HPO 4 2- NH 3 CO 3 2- PO 4 3- OH -

72. K w = [H 3 O + ][OH - ] 1 x 10 -14 = [H 3 O + ][OH - ] K eq equilibrium constant KwKw water dissociation constant KaKa acid dissociation constant KbKb base dissociation constant

73. H + + NH 3 NH 4 + NH 4 + H + + NH 3 acid CB CA base HA H + + A - strong acid weak acid 0.1 M ? M

74. Conjugate Acid Strength Very strong Strong Weak Very weak Relative acid strength Relative acid strength Relative conjugate base strength Relative conjugate base strength Very weak Very strong Weak Strong HA H + + A - pK a = [H + ] [A - ] [HA] Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 508

75. [C] [D] [Products] A (g) + 2 B (g) 3 C (g) + D (g) Weak Acids (pK a ) Weak Acids – dissociate incompletely (~20%) Strong Acids – dissociate completely (~100%) Equilibrium constant (K eq )= K eq = LeChatelier’s Principle (lu-SHAT-el-YAY’s) [Reactants] [A][B] 3 2

76. H + (aq) + C 2 H 3 O 2 1- (aq) CH 3 COOH (aq) HC 2 H 3 O 2 (aq) [Reactant] [Product] Equilibrium constant K eq = = K a = Acid dissociation constant K a = 1.8 x 10 -5 @ 25 o C for acetic acid [H + ][C 2 H 3 O 2 1- ] [HC 2 H 3 O 2 ] [H + ][C 2 H 3 O 2 1- ] [HC 2 H 3 O 2 ] KaKa = [H + ][C 2 H 3 O 2 1- ] [HC 2 H 3 O 2 ] = 1.8 x 10 -5 Assume we begin with 0.1 M acetic acid. [0.1 M ] [X ][X ] X 2 = 1.8 x 10 -6 M = 1.34 x 10 -3 M [H + ] X pH = -log[H + ] pH = -log[1.34 x10 -3 ] pH = 2.87

77. HC 2 H 3 O 2 H + + C 2 H 3 O 2 1- HCl H + + Cl 1- very large HNO 3 H + + NO 3 1- very large H 2 SO 4 H + + HSO 4 1- large 1.8 x 10 -5 H 2 S H + + HS 1- 9.5 x 10 -8 Ionization Constants for Acids KaKa

78. H 2 SO 4 2 H + + SO 4 2- in dilute solutions...occurs ~100% H 2 SO 4 H + + HSO 4 1- & HSO 4 1- H + + SO 4 2- One gram of concentrated sulfuric acid (H 2 SO 4 ) is diluted to a 1.0 dm 3 volume with water. What is the molar concentration of the hydrogen ion in this solution? What is the pH? x mol H 2 SO 4 = 1 g H 2 SO 4 Solution) First determine the number of moles of H 2 SO 4 Sample 1) = 0.010 mol H 2 SO 4 OVERALL: pH = - log [H + ] pH = 1.69 0.010 M0.020 M substitute into equation pH = - log [0.020 M] 98 g H 2 SO 4 1 mol H 2 SO 4

79. A volume of 5.71 cm 3 of pure acetic acid, HC 2 H 3 O 2, is diluted with water at 25 o C to form a solution with a volume of 1.0 dm 3. Step 2) Find the number of moles of acid. x mol acetic acid = 6.00 g HC 2 H 3 O 2 = 0.10 mol acetic acid (in 1 L) M = 0.1 molar HC 2 H 3 O 2 Step 3) Find the [H + ] K a = Step 1) Find the mass of the acid Mass of acid = density of acid x volume of acid = 1.05 g/cm 3 x 5.71 cm 3 = 6.00 g Molarity: M = mol / L Substitute into equationM = 0.10 mol / 1 L What is the molar concentration of the hydrogen ion, H +, in this solution? (The density of pure acetic acid is 1.05 g/cm 3.) (From the formula of acetic acid, you can calculate that the molar mass of acetic acid is 60 g / mol). 60 g HC 2 H 3 O 2 1 mol HC 2 H 3 O 2

80. Step 3) Find the [H + ] 1.8 x 10 -5 = K a = 1.8 x 10 -5 @ 25 o C for acetic acid K a = Substitute into equation: x 2 = 1.8 x 10 -6 M x = 1.3 x 10 -3 molar = [H + ] HC 2 H 3 O 2 H + + C 2 H 3 O 2 1- 0.1 M pH = - log[H + ] pH = - log [1.3 x10 -3 M] pH = 2.9 ? 0.1 M weak acid How do the concentrations of H + and C 2 H 3 O 2 1- compare?

81. Moles of Acid used to make 1 L of solution H+H+ pH 0.010 mol H 2 SO 4 Strong acid 0.100 mol HC 2 H 3 O 2 Weak acid Note: although the sulfuric acid is 10x less concentrated than the acetic acid... …it produces > 10x more H + H + Concentrations …Strong vs. Weak Acid pH = - log[H + ] 1.7 2.9 0.0200 M 0.0013 M

82. 1a) What is the molar hydrogen ion concentration in a 2.00 dm 3 solution of hydrogen chloride in which 3.65 g of HCl is dissolved? 1b) pH 2a) What is the molar concentration of hydrogen ions in a solution containing 3.20 g of HNO 3 in 250 cm 3 of solution? 2b) pH 3a) An acetic acid solution is 0.25 M. What is its molar concentration of hydrogen ions? 3b) pH 4) A solution of acetic acid contains 12.0 g of HC 2 H 3 O 2 in 500 cm 3 of solution. What is the molar concentration of hydrogen ions? 1a) 0.0500 M2a) 0.203 M3a) 2.1 x 10 -3 M 4) 2.7 x 10 -3 M 1b) pH = 1.32b) pH = 0.73b) pH = 2.7 Practice Problems:

83. Weak Acids Cyanic acid is a weak monoprotic acid. If the pH of 0.150 M cyanic acid is 2.32. calculate K a for cyanic acid. HCN(aq) H + (aq) + CN 1- (aq) H 3 O + (aq) 0.150 M 4.8 x 10 -3 M K a = [Products] [Reactants] K a = [H 3 O + ] [HCN] [CN 1- ] K a = [4.8 x 10 -3 M] [0.150 M] [CN 1- ][4.8 x 10 -3 M] K a = 1.54 x 10 -4 4.8 x 10 -3 M pH = -log[H 3 O + ] 10 -pH = [H 3 O + ] 10 -2.32 = [H 3 O + ] 4.8 x10 -3 M = [H 3 O + ]

84. Acid Dissociation Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 280 HCl Conjugate base Acid Conjugate pair + 1- Cl H

85. Conjugate Acid-Base Pairs HCl + H 2 O H 3 O + + Cl - acidbase acid conjugates HCl + H 2 O H 3 O + + Cl - acid base CA CB

86. Conjugate Acid-Base Pairs NH 3 + H 2 O NH 4 1+ + OH - baseacid base conjugates base acid CA CB NH 3 + H 2 O NH 4 1+ + OH -

87. Water is Amphoteric base acid CA CB NH 3 + H 2 O NH 4 1+ + OH - HCl + H 2 O H 3 O + + Cl - acid base CA CB Amphoteric or Amphiprotic substances Amphoteric or Amphiprotic substances: Substances which can act as either proton donors (acids) or proton acceptors (bases) depending on what substances are present.

88. Amphoteric 1- ++ sulfuric acid H 2 SO 4 water H2OH2O hydrogen sulfate ion HSO 4 - hydronium ion A substance that can act as either an acid or a base. H3O+H3O+ 1+ 1- ++ sulfate ion SO 4 2- water H2OH2O hydrogen sulfate ion HSO 4 - hydroxide ion OH - 1- 2-

89. 1- + + 1+ sulfuric acid H 2 SO 4 water H2OH2O hydrogen sulfate ion HSO 4 - hydronium ion H3O+H3O+ (HSO 4 - as a base) Amphoteric A substance that can act as either an acid or a base.

90. Amphoteric A substance that can act as either an acid or a base. 1- + hydrogen sulfate ion HSO 4 - hydroxide ion OH - 1- + sulfate ion SO 4 2- water H2OH2O 2- (HSO 4 - as an acid)

91. Conjugate Acid-Base Pairs HC 2 H 3 O 2 + H 2 O H 3 O 1+ + C 2 H 3 O 2 - acid 1 base 1 base 2 acid 2 conjugates acid base CA CB HC 2 H 3 O 2 + H 2 O H 3 O 1+ + C 2 H 3 O 2 - The reaction proceeds in the direction such that the stronger acid donates its proton to the stronger base.

92. Litmus Paper Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

93. pH Paper pH 0 1 2 3 4 5 6 pH 7 8 9 10 11 12 13

94. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

95. Desired Features of Sensors pH paper 1904 Detection limit Low deflection High sensitivity High selectivity Wide dynamic range Simple to use Cost-effective pH 0 1 2 3 4 5 6 pH 7 8 9 10 11 12 13

96. Range and Color Changes of Some Common Acid-Base Indicators Indicators pH Scale 1234567891011121314 Methyl orange red 3.1 – 4.4 yellow Methyl red red 4.4 6.2 yellow Bromthymol blue yellow 6.2 7.6 blue Neutral red red 6.8 8.0 yellow Phenolphthalein colorless 8.0 10.0 red colorless beyond 13.0 Bromthymol blue indicator would be used in titrating a strong acid with a strong base. Phenolpthalein indicator would be used in titrating a weak acid with a strong base. Methyl orange indicator would be used in titrating a strong acid with a weak base.

97. 1 2345 6 789101112 Indicator Phenolphthalein Methyl Red Orange IV ColorlessPinkRed OrangeYellow OrangePeachYellow pH phenolphthalein methyl red methyl orange

98. Red Cabbage Indicator Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

99. Phenolphthalein Indicator Colorless = Acidic pH Pink = Basic pH H+H+

100. 23 24 How to read a buret volume 23.45 mL (not 24.55 mL) 24.55 mL?

101. Titration TitrationTitration –Analytical method in which a standard solution is used to determine the concentration of an unknown solution. standard solution unknown solution Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

102. Equivalence point (endpoint)Equivalence point (endpoint) –Point at which equal amounts of H 3 O + and OH - have been added. –Determined by… indicator color change Titration dramatic change in pH Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

103. Titration ? M of HCl 30.0 mL of 2.0 M of NaOH If it requires 10.5 mL of ? M HCl to titrate 30.0 mL of 2.0 M NaOH to its endpoint: what is the concentration of the HCl? M 1 V 1 = M 2 V 2 M V = M V H+H+ H+H+ OH - HCl (aq)  H + (aq) + Cl - (aq) 0.1 M H 2 SO 4 (aq)  2 H + (aq) + SO 4 2- (aq) 0.1 M “0.2 M”0.1 M proper term is Normality (N) M V n = M V n H+H+ H+H+ OH - Al(OH) 3 (aq)  Al 3+ (aq) + 3 OH - (aq) 10.5 mL HCl must be ~ __x more concentrated than the NaOH. 6 (x M)(10.5 mL) = (2.0 M)(30.0 mL) X = 5.7 M 30.0 mL of NaOH with bromthymol blue indicator muriatic acid sunnyside 0.1 molar H 2 SO 4 is 0.2 normal 10.5 mL of HCl Endpoint of titration is reached…color change.

104. Titration moles H 3 O + = moles OH - M  V  n = M  V  n M:Molarity V:volume n:# of H + ions in the acid or OH - ions in the base Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

105. Solution of NaOH Solution of KOH Solution of H 2 SO 4 50.0 mL Titration 42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H 2 SO 4 Find the molarity of H 2 SO 4. H3O+H3O+ M = ? V = 50.0 mL n = 2 OH - M = 1.3M V = 42.5 mL n = 1 MV# = MV# M(50.0mL)(2) =(1.3M)(42.5mL)(1) M = 0.55M H 2 SO 4 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

106. Titration  Titration Analytical method in which a standard solution is used to determine the concentration of an unknown solution. standard solution unknown solution Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

107.  Equivalence point (endpoint) Point at which equal amounts of H 3 O + and OH - have been added. Determined by… indicator color change Titration dramatic change in pH Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

108. Titration moles H 3 O + = moles OH - M  V  n = M  V  n M:Molarity V:volume n:# of H + ions in the acid or OH - ions in the base Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

109. Titration  42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H 2 SO 4. Find the molarity of H 2 SO 4. H3O+H3O+ M = ? V = 50.0 mL n = 2 OH - M = 1.3M V = 42.5 mL n = 1 MV# = MV# M(50.0mL)(2) =(1.3M)(42.5mL)(1) M = 0.55M H 2 SO 4 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

110. Acid-Base Titration

111. Calibration Curve Acid (mL) Base (mL) 0.10 M HCl ? M NaOH 0.00 mL 1.00 mL 2.00 mL 4.00 mL 9.00 mL 17.00 mL 27.00 mL 42.00 mL 1.00 mL 2.00 mL 5.00 mL 8.00 mL 10.0 mL 15.0 mL 1)Create calibration curve of six data points 2)Using [HCl], determine concentration of NH 3 3)Determine vinegar concentration using [NaOH] determined earlier in lab Solution of NaOH Solution of NaOH Solution of HCl 5 mL Data Table

112. Titration Curve

113. Calibration Curve Acid (mL) Base (mL) pH endpoint equivalence point indicator base 7 pink - changes color to indicate pH change e.g. phenolphthalein is colorless in acid and pink in basic solution Pirate…”Walk the plank” once in water, shark eats and water changes to pink color

114. Titration Curve Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 527

115. equivalence point 14.0 12.0 10.0 8.0 6.0 4.0 2.0 0.0 10.020.030.040.0 pH Volume of 0.100 M NaOH added (mL) Titration of a Strong Acid With a Strong Base Solution of NaOH Solution of NaOH Solution of HCl H+H+ H+H+ H+H+ H+H+ Cl Cl - Na + OH - Acid-Base Titrations Adding NaOH from the buret to hydrochloric acid in the flask, a strong acid. In the beginning the pH increases very slowly. Adding additional NaOH is added. pH rises as the equivalence point is approached. Additional NaOH is added. pH increases and then levels off as NaOH is added beyond the equivalence point.

116. equivalence point 14.0 12.0 10.0 8.0 6.0 4.0 2.0 0.0 10.020.030.040.0 pH Volume of 0.100 M NaOH added (mL) Titration of a Strong Acid With a Strong Base 0.00 1.00 10.00 1.37 20.00 1.95 22.00 2.19 24.00 2.70 25.00 7.00 26.00 11.30 28.00 11.75 30.00 11.96 40.00 12.36 50.00 12.52 NaOH added (mL) pH Titration Data Solution of NaOH Solution of NaOH Solution of HCl H+H+ H+H+ H+H+ H+H+ Cl - Na + OH - 25 mL phenolphthalein - colorless phenolphthalein - pink Bromthymol blue is best indicator: pH change 6.0 - 7.6 Yellow Blue

117. Titration of a Strong Acid With a Strong Base equivalence point 14.0 12.0 10.0 8.0 6.0 4.0 2.0 0.0 10.020.030.0 pH Volume of 0.500 M NaOH added (mL) Color change methyl violet Color change bromphenol blue Color change bromthymol blue Color change phenolpthalein Color change alizarin yellow R (20.00 mL of 0.500 M HCl by 0.500 M NaOH) Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 680

118. equivalence point 14.0 12.0 10.0 8.0 6.0 4.0 2.0 0.0 10.020.030.040.0 pH Volume of 0.100 M NaOH added (mL) Titration of a Weak Acid With a Strong Base 0.00 2.89 5.00 4.14 10.00 4.57 12.50 4.74 15.00 4.92 20.00 5.35 24.00 6.12 25.00 8.72 26.00 11.30 30.00 11.96 40.00 12.36 NaOH added (mL) pH Titration Data Titration of a Weak Acid With a Strong Base Phenolphthalein is best indicator: pH change 8.0 - 9.6

119. equivalence point 14.0 12.0 10.0 8.0 6.0 4.0 2.0 0.0 10.020.030.040.0 pH Volume of 0.100 M HCl added (mL) Titration of a Weak Base With a Strong Acid 0.00 11.24 10.00 9.91 20.00 9.47 30.00 8.93 40.00 8.61 45.00 8.30 47.00 7.92 48.00 7.70 49.00 7.47 50.00 5.85 51.00 3.34 HCl added (mL) pH Titration Data Titration of a Weak Base With a Strong Acid 50.0

120. 7. What is the pH of a solution made by dissolving 2.5 g NaOH in 400 mL water? Determine number of moles of NaOH x mol NaOH = 2.5 g NaOH0.0625 mol NaOH Calculate the molarity of the solution [Recall 1000 mL = 1 L] M NaOH = 0.15625 molar NaOHNa 1+ + OH 1- 0.15625 molar pOH = -log [OH - ] pOH = -log [0.15625 M] pOH = 0.8 pOH + pH = 14 or k W = [H + ] [OH - ] 1 x 10 -14 = [H + ] [0.15625 M] [H + ] = 6.4 x 10 -14 M pH = -log [H + ] pH = 13.2 pH = -log [6.4 x 10 -14 M] 0.8 + pH = 14

121. What volume of 0.5 M HCl is required to titrate 100 mL of 3.0 M Ca(OH) 2 ? x = 600 mL of 0.5 M HCl HCl H 1+ + Cl 1- 0.3 mol HCl + Ca(OH) 2 CaCl 2 + HOH 22 x mL 0.5 M 100 mL 3.0 M M 1 V 1 = M 2 V 2 (0.5 M) (x mL) = (3.0 M) (100 mL) x = 1200 mL of 0.5 M HCl M 1 V 1 = M 2 V 2 (0.5 M) (x mL) = (6.0 M) (100 mL) Ca(OH) 2 Ca 2+ + 2OH 1- 0.3 mol 0.6 mol 0.3 mol M mol L HCl mol HCl = M x L mol = (0.5 M)(0.6 L) mol = 0.3 mol HCl Ca(OH) 2 mol = (3.0 M)(0.1 L) mol = 0.3 mol Ca(OH) 2 mol = M x L Ca(OH) 2 [H + ] = [OH - ] "6.0 M"

122. 6. 10.0 grams vinegar M mol L NaOH mol NaOH = M x L mol = (0.150 M)(0.0654 L) mol = 0.00981 mol NaOH titrated with65.40 mL of 0.150 M NaOH (acetic acid + water) moles NaOHmoles HC 2 H 3 O 2 = therefore, you have... 0.00981 mol HC 2 H 3 O 2 B) A) x g HC 2 H 3 O 2 = 0.00981 mol HC 2 H 3 O 2 0.59 g HC 2 H 3 O 2 C)% = % = 5.9 % acetic acid Commercial vinegar is sold as 3 - 5 % acetic acid

123. 49 mL 0.2 M HCl + 50 mL 0.2 M NaOH A) mol HCl = M. L mol HCl = (0.2 M). (0.049 L) mol HCl = 0.0098 mol B) mol NaOH = M. L mol NaOH = (0.2 M). (0.05 L) mol NaOH = 0.010 mol 49 mL 0.2 M HCl 50 mL 0.2 M NaOH 99 mL H 2 O 1 mL of 0.2 M NaOH 0.010 mol OH 1- 0.0098 mol H 1+ - 0.0002 mol OH 1- “net” 1) What is the pH of a solution made by combining 49 mL of 0.2 M HCl with 50 mL of 0.2 M NaOH? 1) What is the pH of a solution made by adding 1mL of 0.2 M NaOH with 99 mL H 2 O? HCl + NaOH H 2 O + NaCl

124. M mol L 49 mL 0.2 M HCl 50 mL 0.2 M NaOH 99 mL H 2 O 1 mL of 0.2 M NaOH 0.010 mol OH 1- 0.0098 mol H 1+ - 0.0002 mol OH 1- “net” 1) What is the pH of a solution made by adding 1mL of 0.2 M NaOH with 99 mL H 2 O? NaOH  Na 1+ + OH 1- Calculate the molarity of the solution [Recall 1000 mL = 1 L] M NaOH = 0.002020 molar NaOHNa 1+ + OH 1- 0.002020 molar pOH = -log [OH - ] pOH = -log [0.002020 M] pOH = 2.7 pOH + pH = 14 or k W = [H + ] [OH - ] 1 x 10 -14 = [H + ] [0.002020 M] [H + ] = 4.95 x 10 -12 M pH = -log [H + ] pH = 11.3 pH = -log [4.95 x 10 -12 M] 2.7 + pH = 14 M = mol L M = 0.0002 mol NaOH 0.0099 L

125. Carboxylic Acid HC 2 H 3 O 2 CH 3 COOH C2H4O2C2H4O2 R - COOH HCC H H O O H carboxylic acid H+H+ : : = acetic acid 1-

126. Lactic Acid H 3 C C CO 2 H H OH Lactic acid C 3 H 6 O 3

127. Titration ? M NaOH1.0 M HCl titrate with 1.00 mL 2.00 mL M 1 V 1 = M 2 V 2 (1.0 M)(1.00 mL) = (x M)(2.00 mL) X = 0.5 M NaOH ? M NaOH1.0 M H 2 SO 4 titrate with 1.00 mL 2.00 mL M 1 V 1 = M 2 V 2 (1.0 M)(1.00 mL) = (x M)(2.00 mL) X = 0.5 M NaOH 2.0 M H 1+ ?

128. Calibration Curve VinegarAmmonia 1 mL 3 mL 5 mL 10 mL 15 mL Acid Base vinegar ammonia vinegar ammonia Using 3 mL vinegar… titrate with 0.130 M NaOH solution. M mol L NaOH mol NaOH = M x L mol = (0.130 M)(0.0196 L) mol = 0.002548 mol NaOH moles NaOH moles HC 2 H 3 O 2 = therefore, you have... 0.002548 mol HC 2 H 3 O 2 B) A) x g HC 2 H 3 O 2 = 0.002548 mol HC 2 H 3 O 2 0.153 g HC 2 H 3 O 2 C) % = % = 5.1 % acetic acid Commercial vinegar is sold as 3 - 5 % acetic acid Calculate molarity (M) of acetic acid. M 1 V 1 = M 2 V 2 Calculate % acetic acid in vinegar. % = part / whole x100

129. Calibration Curve VinegarAmmonia 1 mL 3 mL 5 mL 10 mL 15 mL vinegar ammonia Using 3 mL vinegar… titrate with 0.130 M NaOH solution. Calculate molarity (M) of acetic acid. M 1 V 1 = M 2 V 2 M 1 V 1 = M 2 V 2 (M acetic acid )(3.0 mL) = (0.130 M NaOH )(19.6 mL) M acetic acid = 0.8493 molar It required 19.6 mL of NaOH to reach the endpoint. Acid Base vinegar ammonia

130. HCl H 2 SO 4 H 3 PO 4 HNO 3 CH 3 COOH HF Hydrochloric acid Sulfuric acid Phosphoric acid Nitric acid Acetic acid Hydrofluoric acid stomach acid, pickling metal battery acid, # 1 selling chemical food flavoring fertilizer, explosives vinegar etch glass NaOH Ca(OH) 2 NH 4 OH sodium hydroxide calcium hydroxideammonium hydroxide

131. pH scale 0 7 14 acid neutral base [H + ] = [OH - ] Soren Sorenson developed pH scale pOH = -log [OH - ] k W = [H + ] [OH - ] pH = -log [H + ] pH + pOH = 14 (alkalinity) Arnold Beckman invented the pH meter H + + H 2 O H 3 O + protonhydronium ion k w = 1 x 10 -14

132. Concentrated vs. Dilute Concentration: Molarity molality Normality M = mol L m = mol kg H 2 SO 4 2 H 1+ + SO 4 2- 3 M “6 M”

133. Strong / Weak Acid Strong HA H + + A - (~100% dissociation) Weak HA H + + A - (~20% dissociation) K a = [Product] [Reactant] acid dissociation constant KaKa 0.8 H 3 PO 4 3H + + PO 4 3- 0.0021 HF H + + F - H 2 A 2 H + + A - K a = [H + ] 2 [A - ] [H 2 A]

134. Acid + Base Salt + Water How would you make calcium sulfate in the lab? + CaSO 4 ACID Sour taste, litmus red Arrhenius – H + as only ion in water Brønsted-Lowry – proton donor BASE bitter taste, litmus blue Arrhenius – OH - as only ion in water Brønsted-Lowry – proton acceptor H 2 SO 4 Ca(OH) 2 + 2 H 2 O ??

135. phenolphthalein colorless pink acid baseweak strong bromthymol blue yellow blue acid basestrong universal indicator R O Y G B I V pH 4 7 12 litmus paper & pH paper Indicators

136. Buffers - salts of weak acids and weak bases that maintain a pH LeChatelier’s Principle - acidosis & alkalosis (bicarbonate ion acts as buffer) - darkening glasses - egg shells thinner in summer (warm) e.g. Aspirin (acetyl salicylic acid) vs. Bufferin low pH upsets stomach

137. Amino Acids – Functional Groups AmineCarboxylic AcidBase Pair NH 2 1- R- COOH NH 3 NH 2 1- NH 4 1+ amineammoniaammonium ion N H H H : N H H H H : 1+ N H H : 1- : H+H+ lose H +

138. Water – Amphiprotic H2OH2OOH 1- H 3 O 1+ hydroxidewaterhydronium ion N H H H : N H H H H : 1+ N H H : 1- : H+H+ lose H +

139. Water – Also Amphiprotic H2OH2OOH 1- H 3 O 1+ hydroxidewaterhydronium ion H+H+ lose H + Amphiprotic – Act as an acid (proton donor) or base (proton acceptor) O 2- H+H+ H+H+  ++

140. Range and Color Changes of Some Common Acid-Base Indicators Indicators pH Scale 1234567891011121314 Methyl orange red 3.1 – 4.4 yellow Methyl red red 4.4 6.2 yellow Bromthymol blue yellow 6.2 7.6 blue Neutral red red 6.8 8.0 yellow Phenolphthalein colorless 8.0 10.0 red colorless beyond 13.0

141. Neutralization of Bug Bites Wasp - stings with base (neutralize with lemon juice or vinegar) Red Ant - bites with acid (neutralize with baking soda)

142. Strength  Strong Acid/Base 100% ionized in water strong electrolyte - + HCl HNO 3 H 2 SO 4 HBr HI HClO 4 NaOH KOH Ca(OH) 2 Ba(OH) 2 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

143. Strength  Weak Acid/Base does not ionize completely weak electrolyte - + HF CH 3 COOH H 3 PO 4 H 2 CO 3 HCN NH 3 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

144. Ionization of Water H 2 O + H 2 O H 3 O + + OH - K w = [H 3 O + ][OH - ] = 1.0  10 -14 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

145. Why is pure water pH = 7? 11 in 500,000,000 water molecules will autoionize. HH 2 O + H 2 O  H 3 O + + OH 1- TThis yields a hydronium ion concentration of 1 x 10 -7 M H 3 O + per liter of solution ppH = -log[H 3 O + ] ppH = -log[1 x 10 -7 ] or pH = 7

146. H H O H H H H H H O H H 1- 1+ H H O H H H H O H H H H O H H H H O H H O H H H H H H O H H

147. Ionization of Water  Find the hydroxide ion concentration of 3.0  10 -2 M HCl. [H 3 O + ][OH - ] = 1.0  10 -14 [3.0  10 -2 ][OH - ] = 1.0  10 -14 [OH - ] = 3.3  10 -13 M Acidic or basic? Acidic Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

148. pH = -log[H 3 O + ] pH Scale 0 7 INCREASING ACIDITY NEUTRAL INCREASING BASICITY 14 pouvoir hydrogène (Fr.) “hydrogen power” Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

149. pH Scale pH = -log[H 3 O + ] pOH = -log[OH - ] pH + pOH = 14 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

150. pH Scale  What is the pH of 0.050 M HNO 3 ? pH = -log[H 3 O + ] pH = -log[0.050] pH = 1.3 Acidic or basic? Acidic Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

151. pH Scale  What is the molarity of HBr in a solution that has a pOH of 9.6? pH + pOH = 14 pH + 9.6 = 14 pH = 4.4 Acidic pH = -log[H 3 O + ] 4.4 = -log[H 3 O + ] -4.4 = log[H 3 O + ] [H 3 O + ] = 4.0  10 -5 M HBr Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem