CBSE Questions for Class 12 Engineering Chemistry Electrochemistry Quiz 12 - MCQExams.com

An aqueous solution of X is added slowly to an aqueous solution of Y as shown in List I. The variation in conductivity of theses reactions is given in List II. Match List I with List II and select the correct using the code given below the lists:
List-A
List-B
(P) $$\underset {X}{(C_2H_5)_3N}+\underset {Y}{CH_2COOH}$$
(1) Conductivity decreases and then increases
(Q) $$\underset {X}{KI(0.1M)}+\underset {Y}{AgNO_3(0.01M)}$$
(2) Conductivity decreases and then does not change much
(R) $$\underset {X}{CH_2COO}+\underset {Y}{KOH}$$(3) Conductivity increases and then does not change much
(S) $$\underset {X}{NaOH}+\underset {Y}{HI}$$(4) Conductivity does not change much and then increases.
  • 3, 4, 2, 1
  • 4, 3, 2, 1
  • 2, 3, 4, 1
  • 1, 4, 3, 2
$$Zn\, |\, Zn^{2+}(c_1)\, ||\, Zn^{2+}(c_2)\, |\, Zn$$. For this cell $$\Delta$$G is negative if:
  • $$c_1\, =\, c_2$$
  • $$c_1\, >\, c_2$$
  • $$c_2\, >\, c_1$$
  • None
The tarnishing of silver ornaments in atmosphere is due to:
  • $$Ag_{2}O$$
  • $$Ag_{2}S$$
  • $$Ag_{2}CO_{3}$$
  • $$Ag_{2}SO_{4}$$
The electrical conductivity of a solution serves as a means of determining the end point in a chemical reaction, involved in the titration of acids, bases, or precipitation. Which of the following conductometric titration represent the curve of HCl vs $$NaOH$$ :
x-axis $$\Rightarrow$$ Volume of alkali added
y-axis $$\Rightarrow$$ Conductivity
$$\Delta G^{\ominus}$$ for the following reaction is:
$$4A1\, +\, 3O_2\, +\, 6H_2O\, +\, 4 \overset {\ominus} OH\, \rightarrow\, 4A1(OH)_4\, ^{\ominus}$$
$$E^{\ominus}_{cell}\, =\, 2.73\, V$$
$$\Delta _fG^{\ominus}_{(\overset {\ominus} OH)}\, =\, -157\, kJ\, mol^{-1}$$
$$\Delta _fG^{\ominus}_{(H_2O)}\, =\, -237\, kJ\, mol^{-1}$$
  • $$-3.16\, \times\, 10^3\, kJ\, mol^{-1}$$
  • $$-0.79\, \times\, 10^3\, kJ\, mol^{-1}$$
  • $$-0.263\, \times\, 10^3\, kJ\, mol^{-1}$$
  • $$+0.263\, \times\, 10^3\, kJ\, mol^{-1}$$
For the electrolytic production of $$NaClO_4$$ from $$NaClO_3$$ according to the reaction $$NaClO_3 + H_2O \rightarrow NaClO_4 + H_2$$. How many faradays of electricity would be required to produce 0.5 mole of $$NaClO_4$$ ?
  • $$1$$
  • $$2$$
  • $$3$$
  • $$1.5$$
Chromium plating can involve the electrolysis of an electrolyte of an acidified mixture of chromic acid and chromium sulphate. If during electrolysis the article being plated increases in mass by 2.6 g and $$0.6\, dm^{3}$$ of oxygen are evolved at an inert anode, the oxidation state of chromiumions being discharged must be : (assuming atomic weight of Cr = 52 and 1 mole of gas at room temperature and pressure occupies a volume of 24 $$dm^{3}$$)
  • $$-1$$
  • $$0$$
  • $$+1$$
  • $$+2$$
The $$E^{\ominus}\, for\, Cu^{2+}/Cu^{\oplus},\, Cu^{\oplus} / Cu,\, Cu^{2+}/Cu,$$ are 0.15 V, 0.50 V, and 0.325 V, respectively. The redox cell showing redox reaction $$2Cu^+, \rightarrow\, Cu^{2+}\, +\, Cu$$ is made. $$ E^{\ominus}$$ of this cell reaction and $$\Delta G^{\ominus}$$ may be :
  • $$E^{\ominus}\, =\, 0.175\, V\, or\, E^{\ominus}\, =\, 0.350\, V$$
  • $$n = 1 $$or $$2$$
  • $$\Delta G^{\ominus}\, =\, -33.775\, kJ$$
  • All of these
The passage of $$1.5$$ Faradays of electricity corresponds to the flow of how many electrons?
  • $$6\times {10}^{23}$$
  • $$9\times {10}^{23}$$
  • $$12\times {10}^{23}$$
  • $$3\times {10}^{23}$$
What happens when lead storage battery is discharged?
  • $${SO}_{2}$$ is evolved
  • Lead sulphate is consumed
  • Lead is formed
  • $${H}_{2}{SO}_{4}$$ is consumed
A process in which a thin film of a metal like gold, silver, etc., is deposited on another metallic article with the help of electricity.
  • Electrorefining
  • Electroplating
  • Electrochemical process
  • Galvanization
Select the correct statement.
  • Faraday represents 96500 coulombs per second.
  • Coulomb represents one ampere for 1/2 second.
  • Coulomb represents 1/2 ampere for 1 second.
  • Faraday represents charge of one mole electron.
The rusting of iron takes place as follows:
$$\displaystyle 2H^{\oplus}\, +\, 2e^-\, +\, \frac{1}{2} O_2\, \rightarrow\, H_2O(l);\, E^{\ominus}\, =\, +1.23\, V$$
$$Fe^{2+}\, (aq)\, +\, 2e^-\, \rightarrow\, Fe(s);\, E^{\ominus}\, =\, -0.44\, V$$
Calculate $$\Delta G^{\ominus}$$ for the net process.
  • $$-322 kJ$$ $$mol^{-1}$$
  • $$-152 kJ$$ $$mol^{-1}$$
  • $$-76 kJ$$ $$mol^{-1}$$
  • $$-161 kJ$$ $$mol^{-1}$$

Separating the components of a homogeneous mixture using electricity is called:
  • electrification
  • electrolysis
  • electrolyte
  • electromagnetisation
1 Faraday can be defined as : 
  • the magnitude of the charge of 1 mole of electrons
  • the magnitude of the electric dipole
  • a fundamental constant of nature equal to $$\displaystyle 6.63\times { 10 }^{ -34 }\ J \ s/photon$$
  • a constant that accounts for the existence of ions in solution
  • the assignment of charges to individual atoms
What is the charge on an individual electron, if one $$F$$ is equivalent to $$96,487$$ $${ C }/{ mol }{ e }$$?
  • $$5.76\times { 10 }^{ 28 }{ C }/{ { e }^{ - } }$$
  • $$6.022\times { 10 }^{ 23 }{ C }/{ { e }^{ - } }$$
  • $$1.6\times { 10 }^{ 19 }{ C }/{ { e }^{ - } }$$
  • $$1.6\times { 10 }^{ -19 }{ C }/{ { e }^{ - } }$$
About an electrolyte cell, which of the following is true?
  • An electric current causes an otherwise non-spontaneous chemical reaction to occur
  • Reduction occurs at the anode
  • A spontaneous electrochemical reaction produces an electric current
  • The electrode to which electrons flow is where oxidation occurs
  • None of the above
For a reaction, where the $$\Delta { G }^{ o }=-553.91kJ$$ and $$2$$ electrons are transferred, what will be the  $${ E }_{ cell }^{ o }$$of the reaction?
  • $$-2.87 V$$
  • $$-0.00287 V$$
  • $$0.00287 V$$
  • $$2.87 V$$
Assertion : A voltaic cell spontaneously converts chemical energy into electrical energy.

Reason : A voltaic cell needs an externally applied current to work.
  • Both Assertion and Reason are true and Reason is the correct explanation of Assertion
  • Both Assertion and Reason are true but Reason is not the correct explanation of Assertion
  • Assertion is true but Reason is false
  • Assertion is false but Reason is true
  • Both Assertion and Reason are false
The below reaction is the gold-plating process reaction.
            $${ Au }^{ 3+ }\left( aq \right) +3{ e }^{ - }\rightarrow Au\left( s \right) $$
If $$0.600 g$$ of $$Au$$ is plated onto a metal, how many coulombs are used?
  • $$299 C$$
  • $$868 C$$
  • $$2,990 C$$
  • $$8,680 C$$
In an electrochemical cell, reduction takes place at anode          BECAUSE           oxidation always takes place at the cathode.
  • F, F
  • F,T
  • T,T
  • T,F
How many Faraday ($$F$$) are required for the reduction of $$1$$ mole of $${ Ni }^{ 2+ }$$ to $$Ni\left( s \right) $$?
  • $$1 F$$
  • $$2 F$$
  • $$96, 487 F$$
  • $$6.022\times { 10 }^{ 23 }F$$
Energy is required for the electrolysis of water because $${H}_{2}O$$ has less chemical potential energy than the products formed, $${H}_{2}$$ and $${O}_{2}$$.
  • True
  • False
Which of these processes could be associated with the following reaction.
$$\displaystyle 2{ H }_{ 2 }O\rightarrow 2{ H }_{ 2 }+{ O }_{ 2 }$$.
i. Electrolysis
ii. Neutralization
iii. Decomposition
  • i only
  • iii only
  • i and iii only
  • i and ii only
  • ii and iii only
A salt bridge :
  • is not used in a galvanic cell
  • completes the circuit
  • increases the rate of reaction
  • provides an easy handle to carry the cell
  • None of these 
The standard Gibbs free energy change $$(\triangle G^{\circ}$$ in $$kJ\ mol^{-1})$$, in a Daniel cell $$(E_{cell}^{\circ} = 1.1V)$$, when $$2$$ moles of $$Zn(s)$$ is oxidized at $$298\ K$$, is closest to:
  • $$-212.3$$
  • $$-106.2$$
  • $$-424.6$$
  • $$-53.1$$
Which one of the following species has maximum conductance in their aqueous solutions?
  • $$K_2PtCl_6$$
  • $$PtCl_4 \cdot 2NH_3$$
  • $$PtCl_4 \cdot 3NH_3$$
  • $$PtCl_4 \cdot 5NH_3$$
The approximate time duration in hours to electroplate 30 g of calcium from molten calcium chloride using a current of 5 amp is :
[Atomic mass of $$Ca = 40$$]
  • 8
  • 80
  • 10
  • 16
Consider a galvanic cell using solid $$Cu$$ and $$Fe$$ metals with their corresponding solutions.
What is the $$E^{\circ}_{cell}$$?
Standard Potential (V)Reduction Half-Reaction
$$2.87$$$$F_{2}(g) + 2e^{-} \rightarrow 2F^{-}(aq)$$
$$1.51$$$$MnO_{4}^{-}(aq) + 8H^{+}(aq) + 5e^{-}\rightarrow Mn^{2+}(aq) + 4H_{2}O(l)$$
$$1.36$$$$Cl_{2}(aq) + 3e^{-} \rightarrow 2Cl^{-}(aq)$$
$$1.33$$$$Cr_{2}O_{7}^{2-} (aq) + 14H^{+}(aq) + 6e^{-} \rightarrow 2Cr^{3+}(aq) + 7H_{2}O(l)$$
$$1.23$$$$O_{2}(g) + 4H^{+}(aq) + 4e^{-}\rightarrow 2H_{2}O(l)$$
$$1.06$$$$Br_{2}(l) + 2e^{-} \rightarrow 2Br^{-}(aq)$$
$$0.96$$$$NO_{3}^{-}(aq) + 4H^{+}(aq) + 3e^{-}\rightarrow NO(g) + H_{2}O(l)$$
$$0.80$$$$Ag^{+}(aq) + e^{-} \rightarrow Ag(s)$$$
$$0.77$$$$Fe^{3+} (aq) + e^{-} \rightarrow Fe^{2+}(aq)$$
$$0.68$$$$O_{2}(g) + 2H^{+}(aq) + 2e^{-}\rightarrow H_{2}O_{2}(aq)$$
$$0.59$$$$MnO_{4}^{-}(aq) + 2H_{2}O(l) + 3e^{-}\rightarrow MnO_{2}(s) + 4OH^{-}(aq)$$
$$0.54$$$$I_{2}(s) + 2e^{-}\rightarrow 2I^{-}(aq)$$
$$0.40$$$$O_{2}(g) + 2H_{2}O(l) + 4e^{-} \rightarrow 4OH^{-}(aq)$$
$$0.34$$$$Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(s)$$
$$0$$$$2H^{+}(aq) + 2e^{-}\rightarrow H_{2}(g)$$
$$-0.28$$$$Ni^{2+}(aq) + 2e^{-}\rightarrow Ni(s)$$
$$-0.44$$$$Fe^{2+}(aq) + 2e^{-}\rightarrow Fe(s)$$
$$-0.76$$$$Zn^{2+}(aq) + 2e^{-}\rightarrow Zn(s)$$
$$-0.83$$$$2H_{2}O(l) + 2e^{-}\rightarrow H_{2}(g) + 2OH^{-}(aq)$$
$$1.66$$$$Al^{3+}(aq) + 3e^{-}\rightarrow Al(s)$$
$$-2.71$$$$Na^{+}(aq) + e^{-} \rightarrow Na(s)$$
$$-3.05$$$$Li^{+}(aq) + e^{-}\rightarrow Li(s)$$
  • $$-0.78 V$$
  • $$0.10 V$$
  • $$0.78 V$$
  • $$-0.10 V$$
A positive cell potential tells you ________.
  • the reaction is not spontaneous
  • the reaction is spontaneous
  • the reaction will not occur without intervention
  • $$\Delta G$$ is negative
  • both the reaction is spontaneous and $$\Delta G$$ is negative
Consider the table of standard reduction potentials shown below.
Half-reaction$$E^{\circ}$$
$$Cl_{2} + 2e^{-}\rightarrow 2Cl^{-}$$$$1.36\ V$$
$$O_{2} + 4H^{+} +  4e^{-} \rightarrow 2H_{2}O$$$$1.23\ V$$
$$2H_{2}O + 2e^{-} \rightarrow H_{2} + 2OH^{-}$$$$-0.83\ V$$
$$Rb^{+} + e^{-} \rightarrow Rb$$$$-2.93\ V$$
Use the information from the table and your knowledge of electrochemistry to predict the CORRECT net ionic equation for the reaction that will occur when an aqueous solution of rubidium chloride undergoes electrolysis.
  • $$2Rb^{+} + 2Cl^{-} \rightarrow 2Rb + Cl_{2}$$
  • $$2H_{2}O \rightarrow 2H_{2} + O_{2}$$
  • $$H_{2} + 2OH^{-} + 2Rb^{+} \rightarrow 2Rb + 2H_{2}O$$
  • $$4Cl^{-} + O_{2} + 4H^{+} \rightarrow 2H_{2}O + 2Cl_{2}$$
A steady current of $$10.0$$ Amps is passed through a nickel production cell of $$15$$ minutes.
Which of the following is the correct expression for calculating the number of grams of nickel produced?
Note:
$$\bullet 1\ faraday = 96,500\ Coulombs$$
$$\bullet$$ The electroytic cell involves the following half-reaction:
$$Ni_{(aq)}^{2+} + 2e^{-} \rightarrow Ni_{(s)}$$
  • $$\dfrac {(10.0)(15)(96,500)}{(59)(60)} g$$
  • $$\dfrac {(10.0)(15)(59)}{(60)(96,500)} g$$
  • $$\dfrac {(10.0)(15)(60)(59)}{(96,500)(2)} g$$
  • $$\dfrac {(96,500)(59)}{(10)(15)(60)(2)} g$$
How many coulombs of electricity are required for the oxidation of one mole of water to dioxygen?
  • $$9.65 \times { 10 }^{ 4 }C$$
  • $$1.93 \times { 10 }^{ 4 }C$$
  • $$1.93 \times { 10 }^{ 5 }C$$
  • $$19.3 \times { 10 }^{ 5 }C$$
Electrolysis of dilute aqueous NaCl solution was carried out by passing 10 mA current. The time required to liberate 0.01 moles of $$H_2$$ gas at the cathode is
  • $$9.65 \times 10^4 s$$
  • $$19.3 \times 10^4 s$$
  • $$28.95 \times 10^4 s$$
  • $$38.6 \times 10^4 s$$
A $$250.0\ ml$$ sample of a $$0.20\ M\ Cr^{3+}$$ is electrolysed with a current of $$96.5\ A$$. If the remaining $$[Cr^{3+}]$$ is $$0.1\ M$$ the duration of process is ?
  • $$25\ sec$$
  • $$50\ sec$$
  • $$100\ sec$$
  • $$75\ sec$$
Two half-cells have potentials $$-0.44$$ and $$0.799$$ volt respectively. These two are coupled to make a galvanic cell. Which of the following will be true?
  • Electrode of half-cell potential $$-0.44\ V$$ will act as anode
  • Electrode of half-cell potential $$-0.44\ V$$ will act as cathode
  • Electrode of half-cell potential $$0.799\ V$$ will act as anode
  • Electrode of half-cell potential $$-0.44\ V$$ will act as a positive terminal
$$W$$ $$g$$ of $$Ag$$ is deposited at the cathode of one electrolytic cell due to the passage of $$1A$$ of current for $$1h$$. The time required for passage of current to deposit $$W$$ $$g$$ of $$Mg$$ by the same value of current is:
  • $$3.0h$$
  • $$9h$$
  • $$2.5h$$
  • $$1h$$
The reaction taking place in the cell $$Pt|\underset {1\ atm}{H_{2}(g)}|HCl(1.0\ M)|AgCl|Ag$$ is:
  • $$AgCl + (1/2)H_{2} \rightarrow Ag + H^{+} + Cl^{-}$$
  • $$Ag + H^{+} + Cl^{-} \rightarrow AgCl + (1/2) H_{2}$$
  • $$2Ag^{+} + H_{2} \rightarrow 2Ag + 2H^{+}$$
  • $$2Ag + 2H^{+} \rightarrow 2Ag^{+} + H_{2}$$
An electrolysis of a oxytungsten complex ion using $$1.10\ A$$ for $$40$$ min produces $$0.838\ g$$ of tungsten. What is the charge of tungsten in the material? (Atomic weight: $$W = 184$$)
  • $$6$$
  • $$3$$
  • $$5$$
  • $$1$$
If $$0.50\ L$$ of a $$0.60\ M\ SnSO_{4}$$ solution is electrolysed for a period of $$30.0$$ min using a current of $$4.60\ A$$. If in electrodes are used, what is the final concentration of $$Sn^{2+}$$ remaining in the solution? $$[at. wt.\ of\ Sn = 119]$$
  • $$0.257\ M$$
  • $$0.544\ M$$
  • $$0.189\ M$$
  • $$0.514\ M$$
Calculate the potential of an indicator electrode, which originally contained $$ 0.1 M \ MnO_4^- $$ and $$ 1.72 M \ H^+ , $$ and was treated with $$Fe^{2+} , $$ necessary to reduce 90% of $$ KMnO_4 $$ to $$Mn^{2+}$$. 

Given: $$[E^0_{MnO_4^-} /_{ Mn^{+2} } = 1.51 V] $$
  • $$ 1.4 V $$
  • $$ 1.5 V $$
  • $$ 1.6 V $$
  • $$ 1.3 V $$
The density of copper is $$ 8 gm/cc.$$ Number of coulumbs required to plate an area of $$ 10 cm \times 10 cm $$ on both sides to a thickness of $$10^{-2} cm $$ using $$ CuSO_4$$ solution as electrolyte is:
  • $$24,250$$
  • $$48,124$$
  • $$96,500$$
  • $$10,000$$
Consider the following half-cell reactions and associated standard half-cell potentials and determine the maximum voltage that can be obtained by combination resulting in spontaneous processes:
$$AuBr_{4}^{-}(aq) + 3e^{-} \rightarrow Au(s) + 4Br^{-}(aq); E^{\circ} = -0.86\ V$$
$$Eu^{3+}(aq) + e^{-} \rightarrow Eu^{2+}(aq); E^{\circ} = -0.43\ V$$
$$Sn^{2+}(aq) + 2e^{-} \rightarrow Sn(s); E^{\circ} = -0.14\ V$$
$$IO^{-}(aq) + H_{2}O(l) + 2e^{-}\rightarrow I^{-}(aq) + 2OH^{-}(aq); E^{\circ} = + 0.49\ V$$
  • $$+ 0.72$$
  • $$+ 1.54$$
  • $$+1.00$$
  • $$+ 1.35$$
Consider the half-cell reduction reactions :
$${ Mn }^{ 2+ }+{ 2e }^{ - }\rightarrow Mn,{ E }^{ o  }=-1.18\ V$$
$${ Mn }^{ 2+ }\rightarrow { 2n }^{ 3+ }+{ e }^{ - },{ E }^{ o }=-1.51\ V$$
The $${ E }^{ o }$$ for the reaction $${ 3Mn }^{ 2+ }\rightarrow { Mn }^{ 0 }+{ 2Mn }^{ 3+ }$$ and possibility of the forward reaction are respectively:
  • -4.18 V and yes
  • +0.33 V and yes
  • +2.69 V and no
  • -2.69 V and no
A current of $$i$$ ampere was passed for $$t$$ sec, through three calls $$P,Q$$ and $$T$$ connected in series. These contain respectively silver nitrate, mercuric nitrate, and mercurous nitrate. At the cathode of the cell $$P,0.216g$$ of $$Ag$$ was deposited. The weights of mercury deposited in the cathode of $$Q$$ and $$R$$ respectively are: (at wt. of $$Hg=200.59$$)
  • $$0.4012$$ and $$0.8024g$$
  • $$0.4012$$ and $$0.2006g$$
  • $$0.2006$$ and $$0.4012g$$
  • $$0.1003$$ and $$0.2006g$$
The heat of combustion of ethanol in a bomb calorimeter is $$-670.48Kcal\ { mol }^{ -1 }$$ at $${25}^{o}C$$. What is $$\Delta E$$ at $${25}^{o}C$$ for the reaction?
  • -269.24 kcal
  • -469.28 kcal
  • -671.07 kcal
  • +770.48 kcal
Electrolysis can be used to determine atomic masses. A current of $$0.550\ A$$ deposits $$0.55$$. If a certain metal in $$100\ minutes$$. Calculate the atomic mass of the metal if eq. mass $$-$$ mole mass $$/ 3$$.
  • $$100$$
  • $$45.0$$
  • $$48.25$$
  • $$144.75$$
Consider an electrochemical cell in which the following reaction occurs and predict which changes will decreases the cell voltage:

$$Fe^{2+} (aq) + Ag^{+}(aq)\rightarrow Ag(s) + Fe^{3+}(aq)$$

(I) decreases the $$[Ag^{+}]$$ (II) increases in $$[Fe^{3+}]$$ (III) increase the amount of $$Ag$$
  • I only
  • II and III only
  • II only
  • I and II only
What will be $$\Delta H$$ for the reaction; $$Ag(s)+\dfrac{1}{2}Hg_2Cl_2(s)\rightarrow AgCl(s)+Hg(l)$$ at $$25^0c$$, if this reaction can be conducted in a cell for which the emf = 0.0455 volt at this temperature with temperature coefficient $$3.389\times10^{-4} volt\, deg^{-1}$$?
  • +1280 cal
  • +640 cal
  • -1280 cal
  • -640 cal
$${ E }^{ 0 }$$ value of $$Mg/{ Mg }^{ +2 }$$ is $$+2.37$$ V of $$Zn/{ Zn }^{ +2 }$$ is $$+0.76$$ V and that of $$Fe/{ Fe}^{ +2 }$$ is $$+0.44$$ V, which of the following statements is correct?
  • Zinc will reduce $${ Mg }^{ +2 }$$
  • Zn will reduce $${ Fe }^{ +2 }$$
  • Mg oxidizes Fe
  • Zn oxidizes Fe
0:0:1


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