141-Maj1-Int_Shuffled

Using activities, calculate the solubility of Ag₂SO₄ (silver sulfate) in 0.025 M aqueous KNO₃ solution at 25^oC. You may ignore the contribution of Ag⁺ and SO₄^2- to the ionic strength.

[K_sp of Ag₂SO₄= 1.5 x 10^-5 and

]

0.0041 M

0.052 M

0.011 M

0.0018 M

0.023 M

Which of the following inert electrolytes produces the largest increase in the degree of dissociation of acetic acid (using DHLL)?

A.	0.10 m CsCl

B.	0.001 m NaCl

C.	0.001 m KBr

D.	0.10 m CuCl₂

E.	0.01 m CaCl₂

Calculate the activity coefficients of K⁺ and SO₄^2- in an 0.002 M aqueous K₂SO₄ solution assuming complete dissociation and that DHLL is valid.

A.	0.918 and 0.71

B.	0.913 and 0.69

C.	0.811 and 0.62

D.	0.811 and 0.71

E.	0.918 and 0.62

What is the pH of a buffer solution that is 0.250 m in CH₃COOH and 0.150 m in CH₃COONa, using the Davies equation (ignore the contributions from acetic acid dissociation in the ionic strength) to calculate

?
pK_a = 4.76 and log γ_±= -0.51 |z₊z_-| [ (√I)/{1 + (√I)) - 0.30 I ]

4.52

4.63

4.21

4.81

4.39

Calculate the mean ionic molality and the mean ionic activity of a 0.115 m K₃PO₄ solution for which the mean ionic activity coefficient is 0.125.

A.	0.0531 m and 0.108

B.	0.155 m and 0.0819

C.	0.262 m and 0.0328

D.	0.295 m and 0.0411

E.	0.105 m and 0.0215

An electrochemical cell is made by connecting an aluminum strip in 0.40 M Al₂(SO₄)₃ solution to a copper strip in 1.00 M CuSO₄ solution at 25 ^oC. Given the following standard potentials: E^o_Cu2+/Cu = 0.34 V and E^o_Al3+/Al = - 1.67 V determine the cell voltage, assuming that the copper strip is gaining mass and that the aluminum strip is losing mass.

2.01 V

2.20 V

1.02 V

1.11 V

0.00 V

Determine the cell voltage for the following electrochemical reaction at 25 ^oC:
Sn(s) + 2 Ag⁺(aq) → Sn²⁺(aq) + 2 Ag(s) E^o_cell = 0.94 V
given that aqueous 0.10 M AgNO₃ and aqueous 0.050 M Sn(NO₃)₂ solutions are used to set up the the cell.

0.90 V

0.94 V

0.92 V

–0.92 V

–0.94 V

The correct Nernst equation for the reaction
Al(s) + 4 OH^-(aq) → Al(OH)₄^-(aq) + H₂(g)
is:

A.	E = E^o - (RT/nF) ln Q with Q = [H₂][Al(OH)₄^-]/{[Al(s)][OH^-]⁴}

B.	E = E^o - (RT/nF) ln Q with Q = P(H₂)[Al(OH)₄^-]/[OH^-]⁴

C.	E = E^o - (RT/nF) ln Q with Q = [Al(OH)₄^-]P(OH^-)/[OH^-]⁴

D.	E = E^o - (RT/nF) ln Q with Q = P(H₂)[Al][(OH)^-]⁴/[Al(s)]

E.	E = E^o - (RT/nF) ln Q with Q = P(H₂)[Al(OH)₄^-]/[OH^-]

Assume that scientists agreed that the following reaction:

Fe²⁺(aq) + 2 e^- → Fe(s) E^o = -0.447 V

versus the standard hydrogen electrode will have a zero voltage. Then the voltage of the following reaction:

I₂(s) + 2 e^- → 2 I^-(aq) E^o = 0.536 V versus the standard hydrogen electrode

versus this new standard will become:

A.	0.259 V versus Fe²⁺/Fe

B.	- 0.536 V versus Fe²⁺/Fe

C.	0.447 V versus Fe²⁺/Fe

D.	0.983 V versus Fe²⁺/Fe

E.	0.931 V versus Fe²⁺/Fe

10.

Which statement is true?

A.	The Gibbs energy change of an electrochemical reaction at equilibrium is equal to zero.

B.	The Gibbs energy change of a spotaneous electrochemical reaction at equili-brium is negative.

C.	The sign of the Gibbs energy change of an electrochemical reaction at equilibrium can be changed by changing the ionic strength.

D.	The Gibbs energy change of an electrochemical reaction at equlibrium can be changed by an applied electric potential.

E.	The sum of the stoichiometric coefficients multiplied by the charges of each reactant and product in an electrochemical reaction is always negative.

11.

Which of the following statements is not correct?

A.	Deviations from ideal behavior begin at lower concentrations for electrolyte solutions than for nonelectrolyte solutions.

B.	The attractive and repulsive interactions between ions depend only on their chemical identity.

C.	In electrolyte solutions the main interaction is the long-range electrostatic Coulomb interaction between ions.

D.	The attractive and repulsive interactions between ions depend on their charge and size.

E.	To make sure that the activity of an ion is equal to its concentration at infinite dilution it is assumed in the hypothetical standard state used that Henry's law is obeyed up to a concentration of unit molality for that ion.

12.

In the relation

the value of

is:

13.

Enough of the salt

is dissolved in a certain amount of water to bring it into equilibrium with the ions

and

. If the equilibrium concentration of the

ions is

and the Debye Hückel Limiting Law (DHLL) is obeyed at this concentration,

for the salt is:

14.

With respect to the variation of

, the mean ionic activity coefficient, with

, the square root of the ionic strength, for an electrolyte which of the following statements is correct?

A.	At all electrolyte concentrations decreases as increases.

B.	At all electrolyte concentrations increases as increases.

C.	Over electrolyte concentrations where increases as increases, lowering lowers γ_± and in turn solubility must increase to keep the solubility product constant, resulting in the salting out effect.

D.	Over electrolyte concentrations where decreases as increases lowering lowers γ_± and in turn solubility must decrease resulting in the salting out effect.

E.	Over electrolyte concentrations where increases as increases lowering raises γ_± and in turn solubility must increase resulting in the salting in effect.

15.

Which of the following statements is incorrect?

A.	At all electrolyte concentrations, the mean ionic activity coefficient is less than 1.

B.	The mean ionic activity coefficient of an electrolyte may be experimentally measured.

C.	At very low electrolyte concentrations, the experimentally measured mean ionic activity coefficient,, is close to thevalue calculated using the Debye-Huckel limiting law (DHLL).

D.	The activities and for the positive and negative ions of an electrolyte cannot be measured experimentally.

E.	At very high electrolyte concentrations, the mean ionic activity coefficient may reach values greater than 1.

16.

Consider the cell
Co(s)|Co²⁺(aq, a = 0.832)||Cr³⁺(aq, a = 0.342)|Cr(s)
at 25 ^oC. Given the following standard potentials:
E^o_Co2+/Co = -0.280 V and E^o_Cr3+/Cr = -0.744 V,
what happens spontaneously in the cell?

A.	Electrons flow from Cobalt to Chromium

B.	Chromium is deposited.

C.	Cobalt dissolves.

D.	Chromium dissolves.

E.	Cannot be determined.

17.

Consider the following cell at 25 ^oC:
Ag(s)|AgNO₃(aq)|AgI(s)|Ag(s)
Given the following standard potentials: E^o_Ag+/Ag = 0.799V and E^o_AgI/Ag = 0.061V,
calculate pK_sp (negative log10 of K_sp) of AgI.

–12.5

12.5

14.5

-14.5

E.	3.379 x 10^-13

18.

For fluorine, F₂, the standard potential is E^o_F2/F- = 2.87 V. Given that
E^o_Zn2+/Zn = -0.761 V, E^o_Ag+/Ag = 0.800, and E^o_Au+/Au = 1.68 V and assuming that fluorine

can be bubbled over a platinum strip without attacking it, what will happen in a cell made

of fluorine bubbling over a platinum strip and one of these metals?

A.	Fluorine oxidizes Zn, but not Ag and Au.

B.	Flourine oxidizes Zn and Ag, but not Au.

C.	Fluorine oxidizes Zn, Ag, and Au.

D.	Fluorine oxidizes none of these metals.

E.	Fluorine reduces all those metals.

19.

In an electrochemical cell the oxidation of oxalic acid (HOOCCOOH) to
carbon dioxide and water by permanganate ions, MnO₄^-, in acidic solution

has an equilibrium constant for which lnK = 778. What can be concluded from that fact?

A.	Almost complete conversion from reactants into products happens.

B.	No reaction happens.

C.	Almost only reactants can be found in the cell.

D.	Permanganate cannot oxidize oxalic acid.

E.	Carbon dioxide and water are reduced to oxalic acid.

20.

Consider the cell (acidic solutions)
Pt(s)|Cl₂(g)|ClO₂^-(aq)||Cr₂O₇^2-(aq)|Cr(s)
where left oxidation and right reduction happens. How many electrons are
exchanged in the cell reaction, balanced with the smallest possible set of
integer coefficients?

A.	0.931 V versus Fe²⁺/Fe

B.	- 0.536 V versus Fe²⁺/Fe

C.	0.447 V versus Fe²⁺/Fe

D.	0.259 V versus Fe²⁺/Fe

E.	0.983 V versus Fe²⁺/Fe

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