Chapter 19

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Chapter 19

  1. 1. Chapter 19Acids, Bases, and Salts
  2. 2. Acids and BasesAcids • vinegar • citrus fruits • carbonated drinks • car battery • lemon juice • teaBases • calcium hydroxide in mortar • antacids • household cleaning agents
  3. 3. Properties of Acids1. Give foods a tart or sour taste a)lemon & vinegar for example2. Aqueous solutions of acids are electrolytes (conduct electricity)3. Acids cause certain chemical indicators to change color.4. Acid + Base Salt + water
  4. 4. Properties of Bases1. Bases have a bitter taste a)soap2. Bases have a slippery feel3. Aqueous solutions of bases are electrolytes (conduct electricity)4. Bases cause certain chemical indicators to change color.5. Acid + Base Salt + water
  5. 5. Arrhenius Acids & BasesChemists recognized the properties of acids andbases, but were unable to propose a theory toexplain their behavior. In 1887, Swedish chemist Svante Arrhenius proposed a revolutionary way of defining and thinking about acids and bases• Acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution.• Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution
  6. 6. Arrhenius AcidsMonoprotic acids – acids that contain oneionizable hydrogen HNO3 – nitric acidDiprotic acids – acids that contain two ionizablehydrogens H2SO4 – sulfuric acidTriprotic acids – acids that contain threeionizable hydrogens H3PO4 – phosphoric acid
  7. 7. Arrhenius Acids• Not all compounds that contain hydrogen are acids Ex. CH4 – methane has weak polar C – H bonds and no ionizable hydrogens. Not an acid.• Not all hydrogens in an acid may be released as hydrogen ions.• Only hydrogens in very polar bonds are ionizable. In the case where hydrogen is joined to a very electronegative element. Ex. HCl hydrogen chloride very polar covalent molecule
  8. 8. Arrhenius Acids• When HCL dissolves in water, it releases hydrogen ions because the hydrogen ions are stabilized by solvation. H2O H – Cl (g) H + (aq)+ Cl- (aq) Hydrogen Hydrogen Chloride chloride ion ion Ionizes to form an aqueous solution of hydronium ions and chloride ions HCl + H 2O H 3O + + Cl-
  9. 9. Arrhenius Acids• Ethanoic acid CH3COOH is a monoprotic acid due to its structure H O H C C O H HThe three H attached to the carbon are in weak polar bonds. They do not ionize.Only the H bonded to the highly electronegative O can be ionized
  10. 10. Arrhenius BasesSodium hydroxide dissociates into sodium ionsand hydroxide ions in aqueous solution. H2ONaOH (s) Na + (aq) + OH- (aq)Sodium Sodium HydroxideHydroxide Ion ion Potassium hydroxide dissociates into sodium ions and hydroxide ions in aqueous solution. H2O KOH (s) K+ (aq) + OH- (aq) Potassium Potassium Hydroxide Hydroxide Ion ion
  11. 11. Arrhenius BasesGroup IA, the alkali metals, react with water toproduce solutions that are basic.Group IA metals are very soluble in water andcan produce concentrated solutions.Group 2A metals are not very soluble in water.Their solutions are always very dilute.
  12. 12. Bronsted-Lowry Acids and BasesArrhenius’ definition of acids and bases is not avery comprehensive one.If defines acids and bases narrowly and does notinclude certain substances that have acidic orbasic properties. Na2CO3 (aq) is basic
  13. 13. Bronsted-Lowry Acids and BasesThe Bronste-Lowry theory defines acid – a hydrogen-ion donor base – a hydrogen-ion acceptorAll acids and bases included in the Arrheniustheory are also acids and bases according to theBronsted-Lowry theory.
  14. 14. Ammonia as a Base Bronsted-Lowry Theory NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)• ammonia is the hydrogen-ion acceptor and therefore a BL base• water is the hydrogen-ion donor and therefore a BL acid.• Hydrogen ions are transferred from water to ammonia, which causes the hydroxide-ion concentration to be greater than it is in pure water.
  15. 15. Conjugate Acids and Bases NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) base acid conjugate acid conjugate base• When ammonia dissolves and reacts with water, NH4+ is the conjugate acid of the base NH3.• OH- is the conjugate base of acid H2O
  16. 16. Conjugate Acids and Bases HCl (g) + H2O (l) Ý H3O+ (aq) + Cl- (aq) acid base conjugate acid conjugate base• HCl is the hydrogen-ion donor – thus a BL acid.• Water is the hydrogen-ion acceptor – thus BL base
  17. 17. Conjugate Acid-Base PairConjugate acid – the particle formed when a base gains a hydrogen ionConjugate base – the particle that remains when an acid has donated a hydrogen ion..Conjugate acids and bases are always paired with a base or an acid, respectively.Conjugate acid-base pairs consists of two substances related by the loss or gain of a single hydrogen ion.
  18. 18. Common Conjugate Acid-Base Pairs Acid Base HCl Cl- H2SO4 HSO4- H 3O + H 2O HSO4- SO42- CH3COOH CH3COO- H2CO3 HCO3- HCO3- CO32- NH4+ NH3 H 2O OH-
  19. 19. Bronsted-Lowry Acids and BasesA water molecule that gains a hydrogen ionbecomes a positively charged hydronium ion(H3O+)Amphoteric – a substance that can act as bothan acid and a base Ex: water H2SO4 + H2O H3O+ + HSO4- NH3 + H2O NH4+ + OH-
  20. 20. Lewis Acids and BasesGilbert Lewis proposed a third Acid Base theoryAcid – accepts a pair of electrons during areactionBase – donates a pair of electrons during areactionConcept is more general than either theArrhenius theory or the Bronsted-Lowry theory.
  21. 21. Lewis Acids and BasesLewis Acid – a substance that can accept a pairof electrons to form a covalent bond.Lewis Base – a substance that can donate a pairof electrons to form a covalent bond. .. H+ + -: O–H O: : .. H H Lewis Lewis Acid Base
  22. 22. Acid Base Definitions Type Acid Base Arrhenius H+ producer OH- producerBronsted Lowry H+ H+ acceptor Electron-pair Electron-pair Lewis acceptor donor
  23. 23. End of Section 19.1
  24. 24. Hydrogen Ions From WaterWater molecules are highly polar and are incontinuous motion.Occasionally, the collisions between watermolecules are energetic enough to transfer ahydrogen ion from one water molecule toanother.Self ionization of water – the reaction in whichwater molecules produce ions
  25. 25. Hydrogen Ions From WaterA water molecule that loses a hydrogen ionbecomes a negatively charged hydroxide ionA water molecule that gains a hydrogen ionbecomes a positively charged hydronium ion H2O (l) OH- (aq) + H+ (aq) Hydroxide ion Hydroxide ionSelf ionization of water – the reaction in whichwater molecules produce ions
  26. 26. Self Ionization of WaterHydrogen ions in aqueous solution have severalnames.• Some chemists call them protons•Some chemists call them hydrogen ions orhydronium ions.For our purposes, either H+ or H3O+ will representhydrogen ions in aqueous solution. H2O + H2O H3O+ + OH-
  27. 27. Self Ionization of WaterThe self-ionization of water occurs to a very small extent.• In pure water at 25˚C, the equilibrium concentration of hydrogen ions and hydroxide ions are each only 1 x 10-7.• In other words the concentration of OH- and H+ are equal in pure water
  28. 28. Neutral SolutionsAny aqueous solution in which H+ and OH- are equal is a neutral solution.
  29. 29. Ion Product Constant for Water When [H+] increases [OH-] decreases When [H+] decreases [OH-] increases LeChatelier’s principle – when a stress is applied to a system in dynamic equilibrium, the system changes in a way that relieves the stressIf additional ions (either H+ or OH-) are added to a solution, the equilibrium shifts. The concentration of the other type of ion decreases. More water molecules are formed in the process.H+ (aq) + OH- (aq) H2O (l)
  30. 30. Ion Product Constant for WaterFor aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0 x 10-14 [H+] x [OH-] = 1.0 x 10-14 This equation is true for all dilute aqueous solutions at 25˚C.Ion-Product Constant for Water (Kw) – the product of the concentrations of the hydrogen ions and hydroxide ions in water + - -14
  31. 31. Ion Product Constant for Water Not all solutions are neutralWhen some substances dissolve in water, they release hydrogen ions.When hydrogen chloride dissolves in water, it forms hydrochloric acid. H2O HCl (g) H+ (aq) + Cl- (aq)
  32. 32. Ion Product Constant for WaterIn the previous HCl solution, the hydrogen-ion concentration is greater than the hydroxide-ion concentration.Acidic Solution – one in which [H+] is greater than [OH-]. The [H+] of an acidic solution is greater than 1 x 10-7
  33. 33. Ion Product Constant for WaterWhen sodium hydroxide dissolves in water, it forms hydroxide ions in solution. H20 NaOH(s) Na+(aq) + OH-(aq)In the above solution, the hydrogen-ion concentration is less than the hydroxide-ion concentration.Basic Solution – one in which [H+] is less than [OH-] The [H+] of a basic solution is less than 1 x 10 -7Basic solutions are also known as alkaline solutions. solutions
  34. 34. The pH ConceptThe pH scale was proposed by Danish Scientist Soren Sorensen in 1909.The pH scale is used to express [H+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14Strongly Neutral StronglyAcidic Basic
  35. 35. Calculating pHThe pH of a solution is the negative logarithm of the hydrogen-ion concentration. pH = -log[H+]
  36. 36. Calculating pHIn neutral solution, the [H+] = 1 x 10-7M. The pH is 7 pH = -log[H+] pH = -log(1 x 10-7) pH = -(log 1 + log 10-7) pH = -(0.0 + -7.0) pH = 7.0
  37. 37. Classifying SolutionsA solution in which [H+] is greater than 1 x 10-7 has a pH less than 7.0 and is acidic.A solution in which [H+] is less than 1 x 10-7 has a pH greater than 7.0 and is basic.The pH of pure water or a neutral aqueous solution is 7.0Acidic solution: pH < 7.0 [H+] > 1 x 10-7MNeutral solution: pH = 7.0 [H+] equals 1 x 10-7MBasic solution: pH > 7.0 [H+] < 1 x 10-7
  38. 38. Calculating pHpH can be read from the value of [H+] if it is written in scientific notation and has a coefficient of 1.Then the pH of the solution equals the exponent, with the sign changed from minus to plus [H+] = 1 x 10-2 has a pH of 2.0 [H+] = 1 x 10-13 has a pH of 13.0
  39. 39. Calculating pHIf the pH is an integer, it is also possible to directly write the value of [H+]. pH = 9.0 then [H+] of 1 x 10-9M pH = 4 then [H+] = 1 x 10-4M
  40. 40. Calculating pOHThe pOH of a solution equals the negative logarithm of the hydroxide-ion concentration pOH = -log [OH-] A neutral solution has a pOH of 7Acidic solution: pOH > 7.0 [OH-] < 1 x 10-7MNeutral solution: pOH = 7.0 [OH-] equals 1 x 10-7MBasic solution: pOH < 7.0 [OH-] > 1 x 10-7
  41. 41. pH and pOH Relationship pOH + pH – 14 pH= 14 – pOH pOH = 14 - pH
  42. 42. pH Significant FiguresFor pH calculation, you should express the hydrogen-ion concentration in scientific notation [H+] = 0.0010M should be written 1.0 x 10-3 0.0010M has two sig figsWrite pH = 3.00 with 2 zeros to the right of the decimal place representing the 2 sig figs
  43. 43. Problem ExampleColas are slightly acidic. If the [H+] in a solution is 1.0 X 10-5 M , is the solution acidic, basic or neutral. What is the [OH-] of this solution?[H+] = 1.0 X 10-5 M which is greater than 1.0 X 10-7 M so solution is acidic Kw = [OH-] x [H+] = 1.0 X 10-14 [OH-] = 1.0 X 10-14 / [H+] [OH-] = 1.0 X 10-14 / 1.0 X 10-5 [OH-] = 1.0 X 10-9
  44. 44. Problem ExampleWhat is the pH of a solution with a hydrogen- ion concentration of 4.2 x 10-10 M? pH = -log [H+] pH = -log (4.2 x 10-10) pH = -(9.3765) pH = 9.38
  45. 45. Problem ExamplepH of an unknown solution is 6.35. What is its hydrogen-ion concentration? pH = -log [H+] 6.35 = -log [H+] -6.35 = log [H+] Using calculator find the antilog of -6.35 4.5 x 10-7 M = [H+]
  46. 46. Problem ExampleWhat is the pH of a solution if the[OH-] =4.0X10-11M? Kw = [H+] x [OH-]= 1 x 10-14 [H+] = 1 x 10-14 / [OH-] [H+] = 1 x 10-14 / 4.0 x 10-11 [H+] =0.25 x 10-3 M [H+] = 2.5 x 10-4 M
  47. 47. Problem Example (con’t)What is the pH of a solution if the[OH-] =4.0X10-11M? pH = -log [H+] pH = -log (2.5 x 10-4) pH = - (-3.60205) pH = 3.60
  48. 48. Acid-Base IndicatorsIndicator - (HIn) is an acid or a base tht undergoes dissociation in a know pH rangeAn indicator is a valuable tool for measuring pH because its acid form and base form have different color in solution. OH- HIn (aq) H+ (aq) + In- (aq) acid H+ base form formThe acid form dominates the dissociation equilibrium at low pH (high [H+]), and the base form dominates the equilibrium at high pH (high [OH-])
  49. 49. Acid-Base IndicatorsFor each indicator, the change from dominating acid from to dominating base form occurs in a narrow range of approximately two pH units.Within this range, the color of the solution is a mixture of the colors of the acid and the base forms.Knowing the pH range over which this color change occurs, can give you a rough estimate of the pH of the solution.
  50. 50. Acid-Base IndicatorsMany different indicators are needed to span the entire pH spectrum. Indicator characteristics that limit their usefulness.• Listed pH values of indicators are usually given for 25ºC. At other temperatures, an indicator may change color at a different pH.• If the solution being tested is not colorless, the color of the indicator may be distorted.• Dissolved salts in a solution may also affect the indicator’s dissociation.Using indicator strips can help overcome these problems.
  51. 51. pH Meters A pH meter makes rapid, accurate pH measurements.• often easier to use than liquid indicators or indicator strips.• Measurements of pH obtained with a pH meter are typically accurate to within 0.01 pH unit of the true pH.• Color and cloudiness of the unknown solution do not affect the accuracy of the pH value• If the solution being tested is not colorless, the color of the indicator may be distorted.
  52. 52. End of section 19.2
  53. 53. Strong Acids Acids are classified as strong or weak depending on the degree to which they ionize in water.• In general, strong acids are completely ionized in aqueous solution. HNO3 - nitric acid HCl - hydrochloric acid H2SO4 - sulfuric acid HClO4 - perchloric acid HBr - hydrobromic acid HI - hydroiodic acid HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)
  54. 54. Weak AcidsWeak acids ionize only slightly in aqueous solution.• Some Weak Acids Acetic Acid H3COOH Boric Acid H3BO3 (all three are weak) Phosphoric Acid H3PO4 (all three are weak) Sulfuric Acid HSO4- (first ionization is strong) CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO- (aq) ethanoic acid water hydronium ethanoate ion ion
  55. 55. Acid StrengthA strong acid completely dissociates in water ([H3O+] is high).A weak acid remains largely undissociated. ([H3O+] is low).
  56. 56. Equilibrium Constant (Keq) Write the equilibrium-constant expression from the balanced chemical equation. CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)Keq = [H3O+] x [ CH3COO- ] [H3COOH] x [H2O] [H2O] constant in dilute solutions
  57. 57. Acid Dissociation Constant (Ka)Ka = Ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form.H3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)Acid Dissociation Constant Ka = [H3O+] x [ CH3COO- ] [CH3COOH]
  58. 58. Acid Dissociation Constant (Ka)Acid dissociation constant reflects the fraction of an acid in the ionized form. (Ka sometimes called ionization constant)If the value of the Ka is small, then the degree of dissociation or ionization of the acid in the solution is small.Weak acids – small Ka valuesStronger the acid – larger the Ka
  59. 59. Acid Dissociation Constant (Ka)Nitrous acid (HNO2) has a Ka of 4.4 x 10-4Acetic acid (CH3COOH) has a Ka of 1.8 x 10-5Nitrous acid is more ionized in solution and a stronger acid
  60. 60. AcidsStrong Acids• Have high [H3O+]• Large dissociation constantWeak Acids• Have low [H3O+]• Small dissociation constant
  61. 61. AcidsDiprotic and triprotic acids lose their hydrogens one at a time.Each ionization reaction has a separate dissociation constant.H3PO4 – 3 separate dissociation constants.
  62. 62. Base Dissociation Constant (Kb)Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution.• Some strong bases are not very soluble in water (calcium hydroxide and magnesium hydroxide)• Small amounts that do not dissolve dissociate completelyWeak bases react with water to form the hydroxide ion and the conjugate acid of the base. NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Ammonia Water Ammonium Ion Hydroxide ion
  63. 63. Base Dissociation Constant (Kb) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Ammonia Water Ammonium Ion Hydroxide ion Only about 1% of ammonia is present as NH4+Equilibrium Constant Keq = [NH4+] x [OH- ] [NH3] x [H2O] [H2O] constant in dilute solutionsBase Dissociation Constant Kb = [NH4+] x [OH- ] [NH3]
  64. 64. Concentration and StrengthThe words concentrated and dilute indicate how much of an acid or base is dissolved in solution.• Number of moles of the acid or base in a given volumeThe words strong and weak refer to the extent of ionization or dissociation of an acid or base• How many of the particles ionize or dissociate into ionsA sample of HCl added to a large volume of water becomes more dilute, but it is still a strong acid.Vinegar is a dilute solution of a weak acid.
  65. 65. End of section 19.3
  66. 66. Acid-Base ReactionsIf you mix a solution of a strong acid containing hydronium ions with a solution of a strong base that has an equal number of hydroxide ions, a neutral solution results.• Final solution has properties that are characteristic of neither an acidic nor a basic solution. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) H2SO4(aq) + 2KOH(aq) K2SO4(aq) + H2O(l)
  67. 67. Neutralization ReactionsReactions of weak acids and weak bases do not usually produce a neutral solution.In general, reactions with which an acid and a base react in an aqueous solution to produce a salt and water are called neutralization reactions. reactions
  68. 68. Making SaltsPrepare potassium chloride by mixing equal molar quantities of hydrochloric acid and potassium hydroxide. HCl + KOH KCl + H20Heating the solution to evaporate the water will leave the salt potassium chloride. In general, the reaction of an acid with a base produced water and salt
  69. 69. TitrationThe number of moles of hydrogen ions provided by the acid are equivalent to the number of hydroxide ions provided by the base. HCl(aq) + NaOH(aq) NaCl (aq) + H20 (l) 1 mole 1 mole 1 mole 1 mole H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l) 1 mole 2 mole 1 mole 2 moleWhen and acid & base are mixed, the Equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.
  70. 70. Sample ProblemHow many moles of sulfuric acid are required to neutralize 0.50 mol of sodium hydroxide? H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l)• Mole ratio of H2SO4 to NaOH is 1:20.50 mol NaOH 1 mol H2SO4 = 0.25 mol H2SO4 2 mol NaOH
  71. 71. Practice ProblemHow many moles of potassium hydroxide are needed to completely neutralize 1.56 mol of phosphoric acid? H3PO4(aq) + 3KOH(aq) K3PO4(aq) + 3H2O(l)1.56 mol H3PO4 3 mol KOH = 4.68 mol H3PO4 1 mol H3PO4
  72. 72. TitrationYou can determine the concentration of acid or base in a solution by performing a neutralization reaction.You must use an appropriate acid-base indicator to show when neutralization has occurred.In the lab, typically phenolphthalein for acid base neutralization reactions.Solutions that contain phenolphthalein turn from colorless to deep pink as the pH of the solution changes from acidic to basic.
  73. 73. Titrationa. Measured volume of an acid solution of unknown concentration is added to a flask
  74. 74. Titrationb. Several drops of the indicator are added to the solution while the flask is swirled
  75. 75. Titrationc. Measured volumes of the base of known concentration are mixed into the acid until the indicator just barely changes color.
  76. 76. TitrationTitration – the process of adding a known amount of solution of known concentration to determine the concentration of another solution.Standard solution – the solution of known concentrationEnd point – the point at which the indicator changes colorYou can also use titration to find the concentration of a base using a standard acid.
  77. 77. TitrationTitration – the process of adding a known amount of solution of known concentration to determine the concentration of another solution.Standard solution – the solution of known concentrationEnd point – the point at which the indicator changes color. The point of neutralizationEquivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions..
  78. 78. Sample ProblemA 25ml solution of H2SO4 is completely neutralized by 18ml of 1.0M NaOH. What is the concentration of the H2SO4 solution? H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l)0.018 L NaOH 1.0 mol NaOH 1 mol H2SO4 = 1L NaOH 2 mol NaOH 0.025L 0.36M H2SO4
  79. 79. Practice ProblemHow many milliliters of 0.45M HCl will neutralize 25.0ml of 1.00M KOH? HCl + KOH H2O + KCl0.025 L KOH 1.0 mol KOH 1 mol HCl 1L KOH 1 mol KOH 1 L HCl 1000 ml HCl = 56 ml HCl 0.45 mol HCl 1 L HCl
  80. 80. Practice ProblemWhat is the molarity of H3PO4 if 15.0 ml is completely neutralized by 38.5 ml of 0.150 M? H3PO4 + 3NaOH 3H2O + Na3PO40.0385 L NaOH 0.150 mol NaOH 1 mol H3PO4 1L NaOH 3 mol NaOH = 0.129M H3PO4 0.015L H3PO4
  81. 81. End of section 19.4
  82. 82. Salt HydrolysisA salt consists of an anion from an acid and a cation from a base.The salt forms as a result of a neutralization reactionAlthough solutions of many salts are neutral, some are acidic and others are basic...
  83. 83. Salt HydrolysisSalt Hydrolysis – the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ion to water.Hydrolyzing salts are usually derived from a strong acid and weak base or from a weak acid and a strong base.In general, salts that produce acidic solutions contain positive ions that release protons to water.Salts that produce basic solutions contain negative ions that attract protons from water.
  84. 84. Salt HydrolysisCH3COONa (aq) CH3COO- (aq) + Na+ (aq)Sodium ethanoate ethanoate ion sodium ionCH3COONa is the salt from a weak acid CH3COOH and a strong base NaOHIn solution the salt is completely ionized.
  85. 85. Salt HydrolysisSalt Hydrolysis – the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ion to water.CH3COO-(aq) + H2O(l) CH3COOH (aq) + OH- (aq) BL base BL acid makes hydrogen-ion hydrogen-ion solution acceptor donor basicThis process is called hydrolysis because it splits a hydrogen ion off a water molecule.Resulting solution contains a hydroxide-ion concentration greater than the hydrogen-ion concentration. Thus the solution is basic
  86. 86. Salt Hydrolysis NH4Cl (aq) NH4+ (aq) + Cl- (aq) Ammonium Ammonium ion Chloride ion chlorideNH4Cl is the salt from a strong acid (hydrochloric acid, HCl) and a weak base (ammonia, NH3) In solution the salt is completely ionized.
  87. 87. Salt Hydrolysis NH4+(aq) + H2O(l) NH3(aq) + H3O+(aq) BL acid BL base makes hydrogen-ion hydrogen-ion solution donor acceptor acidicThis process is also called hydrolysis because it splits a hydrogen ion off a water molecule.Resulting solution contains a hydrogen-ion concentration greater than the hydroxide-ion concentration. Thus the solution is acidic
  88. 88. Salt Hydrolysis Equivalence Point Strong Acid pH= 7 neutral Strong Base Weak Acid pH > 7 basic Strong Base Strong Acid pH < 7 acidic Weak BaseEquivalence point – the point in a titration wherethe number of moles of hydrogen ions = number ofmoles of hydroxide ions
  89. 89. BuffersBuffer – a solution in which the pH remains relatively constant when small amounts of acid or base are added.A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts.A buffer solution is better able to resist drastic changes in pH than is pure water.
  90. 90. BuffersA solution of ethanoic acid (CH3COOH) and sodium ethanoate (CH3COONa) is an example of a typical buffer.CH3COOH and CH3COO- (source is the completely ionized CH3COONa) act as reservoirs of neutralizing power.
  91. 91. Buffers CH3COO-(aq) + H+(aq) CH3COOH (aq) ethanoate ion hydrogen ion ethanoic acidWhen an acid is added to the solution, the ethanoate ions act as a hydrogen-ion sponge.CH3COOH (aq) + OH-(aq) CH3COO-(aq) + H2O (l)Ethanoic acid hydroxide ion ethanoate ion waterWhen a base is added to the solution, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion.
  92. 92. BuffersThe ethanoate ion is not strong enough base to accept hydrogen ions from water extensively.The buffer solution cannot control the pH when too much acid is added, because no more ethanoate ions are present to accept hydrogen ions.Buffer also become ineffective when too much base is added. No more ethanoic acid molecules are present to donate hydrogen ions.
  93. 93. BuffersWhen too much acid or base is added, the buffer capacity is exceeded.Buffer capacity – the amount of acid or base than can be added to a buffer solution before a significant change in pH occurs.
  94. 94. BuffersWhen a base is added to a buffered solution, the acidic form removes hydroxide ions from the solution.When an acid is added to a buffered solution, the basic form removes hydrogen ions from the solution.
  95. 95. Buffers & Your BloodYour body function properly only when the pH of your blood lies between 7.35 and 7.45.Your blood contains buffers (hydrogen carbonate ions and carbonic acid) HCO3- (aq) + H+ (aq) H2CO3 (aq) hydrogen hydrogen ions carbonic acid carbonate ionAs long as there are hydrogen carbonate ions available, the excess hydrogen ions are removed, and the pH of the blood changes very little.
  96. 96. End of Section Chapter 19

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