CH 15 Txt

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CH 15 Txt
2013-04-26 22:49:47
CHM 122

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  1. Because the neutralizaiton reaction of any strong acid with a strong base is what?
    The equilibrium constant is...
    • the reverse of the dissocation of water
    • the reciprocal of the ion-product constant for water
  2. The value of Kn for a strong acid-strong base reaction is __, which means what?
    • a very large number
    • the neutralization reaction proceeds essentially 100% to completion
  3. After neutralization of equal molar amounts of acid and base, the __ derived from a strong base and a strong acid.
    solution contains a salt
  4. Because a weak acid HA is largely undissociated, the net ionic equation for the neutralization reaction of a weak acid with a strong base involves __.
    proton transfer from HA to the strong base, OH-
  5. To obtain the equilibrium constant Kn for the neutralization of acetic acid, what do we do?
    multiply known equilibrium constants for reactiosn that add to give the net ionic equation for hte neutralization reaction
  6. Because CH3CO2H is on the left side of the equation and CH3CO2- is o the right side, one of hte reactiosn needed is __.
    Because H2O is on teh right side of the equation and OH- is on hte left side, the other reaciton needed is __.
    • the dissociation of CH3CO2H
    • the reverse of hte dissociation of H2O
  7. True or False:
    The equilirbium constant for hte net reaction equals the product of hte equilibrium constants for hte reaction added.
  8. How do we get hte Kn for a neutralization reaction?
    multiply Ka for the acid by the reciprocal of Kw
  9. What does a resulting large Kn value say?
    the neutralization reaction proceeds nearly 100% to completion
  10. As a general rule, the neutralizaton of any weak acid with a strong base will what?
    go 100% to completion because OH- has a great affinity for porotns
  11. After neutralization of equal molar amounts of CH3CO2H and NaOH, the solution contains , what?
    Therefore, the pH will be
    • Na+ (no acidic/basic properties)
    • CH3CO2- (weak base)
    • greater than seven
  12. A strong acid HA is completely dissociated into __, and its neutralization reaction with a weak base therefore involves what?
    • H3O+ and A- ions
    • proton transfer from H3O+ to the weak base B
  13. Explain the neutralizaiton of a weak base and strong acid? What will be the pH?
    • goes 1--% to competion because H3O+ is a powerful proton donor
    • less than 7= acidic
  14. Weak acid-weak base
    • largely undissociated 
    • neutralization reaction between them involves proton transfer from weak acid to weak base
  15. How can the equilibrium constant Kn be found?
    • multiplying equilibrium constants for
    • 1) acid dissociation of acetic acid (ex)
    • 2) base protonation of ammonia
    • 3) reverse of hte dissociation of water
  16. In general, weak acid-weak base what?
    have less tendency to proceed to completion than neturalizations with trong acids/ bases
  17. A solution fo a __ and its __ is an important acid-base mixture because such mixtures regulate the __ in biological systems.
    • weak acid
    • conjugate base
    • pH
  18. the decrease in [H3O+] on adding acetate ions to an acetic acid solution is an example of the __, the shift in an equilibrium on adding a substance that provides more of an ion alrady involveed in the equilibrium
    common-ion effect
  19. The __ is just another example of __, in which the stress on the equilibrium of raising one of the product concentrations is relieved by shifting the equilibrium to the reactant side.
    Le Chat princi.
  20. Solutions which contain a weak acid adn its conjugate base are called __. Why?
    • buffer solutions
    • they resist drastic changes in pH
  21. If a small amount of OH- is added to a buffer solution, what happens?
    the pH increases, but not by much because the acid component of the buffers olutioni neutralizes the added OH-;same for acids except the pH will go down
  22. What are the components of a bufer solution
    • weak acid (CH3CO2H, HF, NH4, H2PO4-)
    • +
    • conjugate base (CH3CO2-, F-, NH3, HPO2-)
  23. buffer capacity
    the buffering ability of a solution; measure of the amount of acid or base that the solution can absorb without a significant change in pH
  24. Buffer capacity is also a measure of what?
    how little the pH changes with the addition of a given amount of acid or base
  25. Buffer capacity depends on wht?
    how many moles of weak acid and conjugate base are present
  26. For equal volumes of solution, the more concentrated the solution, what?
    the greater the buffer capacity
  27. For solutions having the same concentration, the greater the volume, what?
    the greater the buffer capacity
  28. dissociation constant of a weak acid
    concentartion ratio [weak acid]/[conjugate base]
  29. Henderson-hasselbalch equation
    pH=pKa + log [Conj. Base]/[Acid]
  30. WHat does te Henderson Hasselbalch equation say?
    the pH of a buffer solution has a value clsoe to the pKa of the weak acid, differing only by the amount log [conj. base]/[acid]
  31. when [conj.base]/[acid]= 1, then log [base]/[acid]= __
    0 and the pH equals the pKa
  32. real importance of the Henderson-hasselbalch equation?
    it tells us how the pH affects the % dissociation of a weak acid
  33. What else does the Hasselbalch equation tell us?
    how to prepare a buffer soluton with a given pH. the general idea is to select a weak acid whose pKa is close to the desired pH and adjust the [conj. base]/[acid] ratio to the value specified by the HH equation
  34. The pH of a buffer solution does not depend on what?
    • the volume fo teh solution because a change in solution volume changes the concentration of the acid adn base by the same amount
    • thus, the pH and [base]/[acid] ratio remain unchanged
  35. True or False:
    Changing the volume of water used to prepare a buffer solution is critical and will change pH.
    True or False:Changing the volume of water used to prepare a buffer solution is not critical and will not change pH.
  36. The pH depends only on?
    pKa and on the relative moalr amounts of the weak acid and conjugate base
  37. With a pH meter, you can receord data to produce a __, a plot of the pH of the solution as a function of the volume of added titrant.
    pH titration curve
  38. The shape of a pH titration curve makes it possible to identify the __ in a titration, the point at which stoichiometrically equivalent quantities of acid and base have been mixed together. Knowing hte shape of it is also useful in selecting a suitable indicator to signal the equivalence point.
    equivalence point
  39. The number f millimoles of H3O+ and OH is what?
    the product of the initial volume of the acid/ base and its molarity
  40. Difference between titration curve of strong-strong and strong base-weak acid
    • initial rise in pH for weak acid is greater
    • increase in pH near equivalence point is maller in weak acid than strong 
    • the pH at the equivalence point is greater than 7 in the weak acid because the anion of a weak acid is a base
  42. Amino acids are both __ and __ and can be protonated by strong acids such as __, yielding salts such as H2A+Cl-. The protonated form of the amino acid has two dissociable protons and can react with two molar amounts of OH- to give first the __ form and then the __ form.
    • acidic 
    • basic
    • HCL
    • neutral
    • anionic
  43. Whi ch part of a polyprotic ion is neutralized first
    the acidic one; the second neutralization is of the electrically neutral group
  44. When looking at the pH of polyprotic acids, what is the pH at the first equivalence point? WHat is it called?
    • average of pKa1 and pKa2
    • isoelectric point
  45. The equilibrium equation for hte dissolution reaction is __, where the equilibrium constant __ is called the __, or simply the __. As usual for a heterogeneous equilibrium, the __ is omitted.
    • Ksp= [M^n+]^m [X^y]^x
    • Ksp
    • solubility product constant
    • solubility product
    • the concentration of hte solid
  46. What does the Ksp always equal
    the proudct of hte equilibrium concentrations fo all the ions on the right side of the chemical equation, witht eh concentrations of each ion raised to the power of its coefficient in the balanced equation
  47. The value of Ksp is unaffected by __.
    the presence of other ions in solution, such as Na+ from NaF and Cl- from CaCl2, as long as the solution is very dilute
  48. As ion concentrations increase, what happens with Ksp values?
    they are somewhat modified because of electrostatic interactions between ions
  49. If hte saturated solution is prepared by a method other than the dissolution of CaF2 in pure water, there are no what?
    separate restrictions on Ca2+ and F-; the only restriction on the ion concentrations ins that the value fo the equilibrium constant expression Ca2+ x F-^2 must equal the Ksp
  50. Once Ksp is measured, you can do uaht?
    use it to calculate the solubility of the compound--the amount of the compound that dissolves per unit volume of saturated solution
  51. What are the two complications in calculting solubilities?
    • first: Ksp values are difficult to measure and values listed in different sources might differ by as much as a factor of ten or more
    • second: calculated solubilites can be less than observed solubilities because of wide reactions
  52. When MgF2 dissolves in a solution that contains a common ion from another source--say F= from NaF-- the position of the solubility equilibrium is __
    shifted to the left by the common ion effect
  53. A smaller value of Mg2+ means that MgF2 is what
    less soluble in a sodium fluoride solution than it is in pure water
  54. An ionic compound that contains a basic anion becomes what?
    more soluble as the acidity of the solution increases
  55. The solubility of CaCO3 what?
    increases with decreasing pH becasuse the CO32- cions ocmbine with protons to give HCO3- ions
  56. The solubility of an ionic compound increases  dramatically if the solution contains a __ that can form a __ to the metal cation.
    • Lewis base
    • coordinate covalent bond
  57. What is a complex ion
    ion that contains a metal cation bonded to one or more small molecules or ions
  58. The stability of a complex ion is measured by its __, the equilibrium constant for the formation of the complex ion from the hydrated metal cation.
    formation constant Kf
  59. The large value of Kf means that the complex ion is quite __
  60. Because K is much larger than __, the solubility equilibrium for AgCl lies __
    much farther to the right
  61. In general, the solubility of an ionic compound odes what?
    increases when the metal cation is tied up in the form of a complex ion
  62. Certain ___ are soluble both in __ and __ solutions.
    • metal hydroxides
    • strongly acidic
    • strongly basic
  63. The answer of whether something is soluble depends on the value of the __, a number defined by the expression __.
    • IP (ion product)
    • IP= [Ca2+]t[F-]t^2
  64. The IP is defined as what?
    the same as Ksp, except that the concentrations in the expression for IP are initial concentrations, not equilibrium
  65. Thus, the IP is actually a __, but the term __ is more descriptive because, as usual, solid CAF2 is omitted
    • reaction quotient
    • ion prdouct
  66. IP >ksp ?
    supersaturated and precipitation will occur
  67. IP = Ksp
    solution is saturated/ equilibrium exists
  68. IP < Ksp:
    solution is unsaturated and precipitation will not occur
  69. A conveninet method for separating a mix of ions is __
    to add a solution that will precipitate some of the ions but not others
  70. What does separation depend on
    adjusting the H3O+ concentration so the reaction quotient Qc exceeds Kspa for hte very insoluble sulfides but not for the more soluble ones
  71. Why do we use Kspa for metal sulfides rather than Ksp
    • 1) the ion separations are carried out in acidic solution (so its more convenient)
    • 2) the value of Ka2 is very small, is highly basic, and is not an important species
  72. Qualitative analysis
    procedure for identifyign the ions present in an unknown solutions
  73. When aquueous HCL is added to the unknown solutoin, the cations of Group I precipitate as __.
    insoluble chlorides: AgCl, Hg2CL2, PbCl2
  74. The cations of groups II-V, which form soluble chlorides, remain in solution: __
    Because the solution is strongly acidic, only the most __. The __.
    After the insoluble chlorides have been removed, the solution is treated with __ too precipitate hte cations of group iI as __.
  75. Group III. At this point, __ is added, neutralizing teh soluton and giving an __ that is slightly __. THe decrease in H+ shifts hte metal sulfide solubility equilinrium to the __, thus precipitating the 2+ cations of group III as insoluble sulfides: 
    The 3+ cations precipitate from the basic solution, not as sulfides but as __.
    • aqueous NH3
    • aNH4+-NH3
    • basic
    • left
    • MnS, FeS, COs, NiS, Zns
    • insoluble hdroxides
  76. After hte base insolble sulfides nad hte insoluble hydroxides have been removed, the solution is treated with (NH4)2Co3 to precipitate the cations of group IV as __:
    Magnesium carbonate doesn't precipitate at htis point because [Co32-] in the NH4+-NH3 buffer is maintained at a low value.
    • CaCO3
    • BaCO3
  77. The only ions remaining in solutiona t this point are those whose __ are soluble under the conditions of previous ions. magnesium ion is spearated and identified by the addition of a solution of (HN4)2HPO4; if Mg2+ is present , a white precipitate of Mg(NH4)PO4 forms. the __ are usually identified yb the characteristic colors they impart to a Bunsen flame
    • chlorides, sulfides, carbonates
    • alkali metal ions