Chapter 17 Test Flashcards
The three intermediate reactions and the final reaction of acetic acid and ammonia are shown below.
CH3COOH(aq) + NH3(aq) <==> NH4+(aq) + CH3OO-(aq)
The reactions are:
Reaction 2: CH3COOH(aq) + H2O(l) <==> CH3OO-(aq) + H3O+(aq)
Reaction 3: NH3(aq) + H2O(l) <==> NH4+(aq) + OH-(aq)
Reaction 4: H3O+(aq) + OH-(aq) <==> 2H2O(l)
Which of the following correctly describes each of the three reactions (2-4)?
Reaction 2: the dissociation of acetic acid in aqueous solution
Reaction 3: the dissociation of ammonia in aqueous solution
Reaction 4: the reverse reaction for the autoionization of water
Look at the equilibrium curve for the reaction of HA(aq) <==> H+(aq) + A-(aq)
Ka = 1*10^-10 for this reaction. Which statement about this plot is not true?
[H + ] [A− ] / [HA] at Point 2 < [H + ] [A− ] / [HA] at Point 5 because [A− ] is larger.
Our blood is able to maintain a pH of about 7.5 because our body acts like a buffer solution. Regulating pH is crucial so that our protein structures remain intact and functional. Which statement about protein in the human body is not correct?
Protein consists partially of amine groups that are neutral as a result of protonation (of one H + ion).
What is the pH of a buffer solution comprised of 3.95 M HCOOH and 0.35 M HCOO− in a 1 L solution after 0.30 mol NaOH (where Ka = 1.77 × 10^−4 ) is added?
3.00
Look at the graphical representation of a buffer solution.
In this experiment, 0.10 moles of HCl were added to 0.30 moles of NH3 to make 1.0 L of solution. Which of the following correctly lists the buffer capacity of this system?
Range A for strong bases and Range C for strong acids
Which of the following acid-base pairs would be best for preparing a buffer solution with a pH of 8.15?
pair #4
Suppose that you wanted to prepare a acetate ion / acetic acid buffer solution with a pH of 4.35. What is the value of [A− ] / [HA] for the correct buffer solution for this event? The Ka for acetic acid is 1.8 × 10^−5
0.407
Suppose that you have a 43 mL solution of 0.23 M HCl that is being titrated with 0.17 M NaOH. You stop the titration after adding 27 mL of NaOH. What is the change in pH at this point?
0.48
Suppose that we are looking at the titration of acetic acid with NaOH. At the equivalence point, we have 0.09 moles of acetate (CH3COO− ). At this point, the following reaction occurs and reaches equilibrium:
where Kb is 5.6 × 10−10.
What is the pH of the solution?
8.85
A diprotic or polyprotic acid __________________
delivers (or “donates”) more than one proton
Look at Point I of this reaction.
Oxalate anion, (Ox)2−, is in solution. It is a weak base. Two of the reactions can occur at this point.
Assume that Reaction 3 is the dominant equilibrium to consider when determining the solution’s pH. What is the pH of the solution if we start with 150 mL of 0.11 M H2(Ox) and titrate it with 0.11 M NaOH? Use the plot for the amounts of NaOH added during the reaction.
8.4
Salt Ksp HgS 7.9*10^-37 Ca(CO3) 3.8*10^-9 CaF2 1.1*10^-11 CuS 3.2*10^-28 Which statement is not correct?
Ca(CO3 ) dissociates the least in aqueous solution
The dissolution of lead(II) chloride is shown by the equilibrium equation
Which of the following statements concerning this reaction is not correct?
By placing the PbCl2(s) in an aqueous solution that contains 0.1 M of HCl, a strong acid, you do not directly affect the solubility of the PbCl2.
For which initial concentration of chromate anion CrO42− would [Ag+ ] = 6.0 × 10−6 M and cause the solution to begin to precipitate Ag2CrO4(s)?
Ag2CrO4(s0 <==> 2Ag+(aq) + CrO42-(aq)
(where Ksp = 9.0 × 10−12 )
0.25 M
Consider the following reaction between nitrous acid and ammonia:
HNO2(aq) + NH3(aq) <==> NH4NO2(aq)
Ka = 4.510^-4
Kb = 1.810^-5
What is the equilibrium constant for this reaction?
8.1 × 10^5
