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ELECTROCHEMISTRY
[Type the document subtitle]
[Pick the date] Neeraj Kumar
ELECTROCHEMISTRY
1.
I
NTRODUCTION:
Electrochemistry is the branch of chemistry which deals with inter conversion of chemical energies and electrical energies. In electrochemistry, we discuss electrolytic reactions (reactions occurring due to passage of electricity through solutions) as well as electrogenetic reactions (reactions producing electric energy).2. ELECTROLYTIC CONDUCTION (ELECTROLYSIS):
(i) Electrolytic conduction is exhibited principally by molten form and by aqueous solutions of salts or substances. It takes place due to free movement of ions.
(ii) Consider electrolytic conduction of molten NaCl between inert electrodes in electrolytic cells as shown in figure. Na+ Cl– NaCl Molten + – anode Cathode DC e – TABLE – 1
Phenomenon At cathode At anode
(a) Ions attracted Cations Anions
(b) Direction of movement of electrons
Into cell Out of cell
(c) Half cell reaction Reduction Oxidation
(Na+ + e → Na) (Cl–→ ½ Cl2 + e)
(d) Sign Negative, since it is
attached to the negative end of battery.
Positive, since it is attached to the positive end of the battery.
(iii) Within the cell, current is carried by the movements of ions, cations towards the negative electrode (called the cathode) and anions towards the positive electrode (called the anode). This movement of ions is called electrolytic conduction.
(iv) Factors, that influence the electrolytic conductivity of solutions of electrolytes are inter ionic attraction, solvation of ions, viscosity of solvent, etc. and these factors depends on the attraction of solute – solute, solute – solvent and solvent – solvent respectively.
(v) The average kinetic energy of the solute ions increases as the temperature is raised and, therefore, the conductance of electrolytic conductors generally increases.
2.1 FARADAY’S LAWS OF ELECTROLYSIS:
FIRST LAW: For the same electrolyte, the mass of a substance produced or consumed at an electrode is directly proportional to the quantity of electricity passed through the electrolytic cell.
W α Q
or, W = Z Q = Z .I.t
Where Z = Electrochemical equivalent of the substantce. When I = 1 ampere and t = 1 second, w = Z, so, electrochemical equivalent (ECE) of any substance may be defined as the mass of the substance produced or consumed when a current of 1 ampere for 1 second (i.e. 1 coulomb) is passed through the electrolyte.
SECOND LAW: When same quantity of electricity is passed through different electrolytes, the amounts of products obtained are proportional to their equivalent weights.
W ∝ E
3. ELECTROLYTIC CONDUCTION:Conductance or conductivity is the property of the conductor which facilitates the flow of electricity through it. It is equal to the reciprocal of resistance,
i.e., Conductance (G) = 1/R Unit: ohm–1 or mho (Ω–1) 3.1 SPECIFIC CONDUCTANCE,λλλλsp:
Reciprocal of specific resistance is called Specific Conductance. We Know that, R α
a
ℓ
(where a : area of the electrode, ℓ : distance between the electrode) ∴ R = ρ
a
ℓ
where ρ = specific resistance of substance.
+ – a a ℓ Now, λsp. = 1 1 R a = ρ ℓ or, λsp = 1 ρ = G.a ℓ
For electrolytic cells, a
ℓ
is known as CELL CONSTANT. Unit of λsp: Ω
–1
cm–1 or Ω–1m–1
3.2 EQUIVALENT CONDUCTANCE:It is the conductance of all ions produced by one gm equivalent of an electrolyte. It is measured by
eq. sp
λ = λ ×V
Where, λeq. = Equivalent conductance λsp = Specific conductance in Ω
–1 cm–1
V = Volume of solution of ml containing 1 gram equivalent of electrolyte Unit of λeq.: Ω
–1
cm2 equiv–1
3.3 MOLAR CONDUCTANCE:It is the conductance of all the ions produced by ionization of one gram mole of an electrolyte. m sp λ = λ ×V Unit of λm: Ω –1 cm2 mole–1
3.4 Conductance of the electrolytes is known to depend upon: nature of electrolyte, i.e., strong or weak, temperature and concentration of electrolyte in solution.
3.5 Variation of molar conductance with concentration by strong electrolyte is given by Huckel Onsagar equation:λm −λm = b C, where b is a constant depending on the nature of the solvent and temperature and λm is molar conductance at infinite dilution. At infinite dilution the concentration, C tends to be zero. At this large dilution λm = λm° as C → 0
3.6 KOHLRAUCH’S LAW:At infinite dilution, when dissociation is complete, each ion makes a definite contribution towards equivalent conductance of the electrolyte irrespective of the nature of the ion with which it is associated and the value of equivalent conductance at infinite dilution for any electrolyte is the sum of contributions of its constituent ions, i.e., λ∞ = λ+ + λ–
3.7 DEGREE OF DISSOCIATION:At infinite dilution the ionization becomes 100%. At lower dilution the ionization (dissociation into ions) is less than 100% and equivalent conductance become lower, i.e.,
λeq<λ°eq ∴Degree of dissociation, α =
dilution
infinite
at
e
conductanc
Equivalent
ion
concentrat
given
a
at
e
conductanc
Equivalent
λ
λ
eq eq=
3.8 THE DRIFT SPEED AND MOBILITY OF IONS:Due to electrical force the ions gets accelerated but due to viscosity of the medium, it is retarded. Therefore it is accelerated only to some limited velocity which depends upon the external field applied, E and the viscosity of the solvent, η. This terminal limiting velocity is called the drift velocity of the ion in solution and is denoted by v. The drift velocity determines the rate of conduction of current. The drift velocity being a vector quantity, instead of v, the term drift speed (denoted as s) having no direction but magnitude equal to the |v| is used to that s = v. The drift speed of an ion is proportional to the applied field E. Thus,
s ∝ E or s = u E or s = u. V.ℓ
Where u is called the mobility of the ion and V is the potential difference across two electrodes at the distance l. From above equation, the mobility of an ion may be defined as the speed of the ion in an electrical field of unit strength (1 V m–1). Conductivity or mobility of ion depends upon: viscosity of the solution (ηηηη), charge on the ion, size of the ion, etc.
4. ELECTRODE POTENTIAL: The tendency to gain or lose electrons by an electronic conductors is called electrode potential. The tentency to lose electron is called electrode oxidation potential and the tendency to gain electron is called electrode reduction potential. Normally, electrode potential means electrode reduction potential. Electrode potential depends on the nature of conductor as well as the concentration of electrolytic solution.
4.1 STANDARD ELECTRODE POTENTIAL:As we know electrode potential depends on concentration of metal ions in solution in which metal electrode is dipped and the temperature of the half cell. The potential difference developed between metal electrode and the solution of its ions of unit molarity
(activity) at
298K is called standard electrode potential.
Standard reduction potential (SRP) = –Standard Oxidation Potential (SOP)
4.2 REFERENCE ELECTRODE: The potential of a single electrode can not be determined but the potential difference between the two electrodes can be accurately measured. Thus we must choose one electrode arbitrarity as a standard and measure the potential difference between this standard electrode and any other electrode we require. By general agreement normal hydrogen reference electrode is takes as the standard reference electrode. Its electrode potential has arbitrarily, been assumed to be 0.00 volt.
4.3 ELECTRO-CHEMICAL SERIES:When the elements are arranged in the order of their standard reduction potentials on the hydrogen scale we get electromotive series or electrochemical series of the elements.
STANDARD ELECTRODE POTENTIALS OF SOME ELECTRODES
Electrode reaction E°(V) Half – Cell representation
Li+ + e– ⇌Li – 3.045 Li+ | Li K+ + e– ⇌K –2.925 K+ | K Na+ + e– ⇌Na –2.714 Na+ | Na Mg++ + 2e– ⇌Mg –2.37 Mg++|Mg H2 + 2e – ⇌2H– –2.25 OH– | H2 | Pt Al3+ + 3e– ⇌Al –1.66 Al3+ | Al 2H2O + 2e – ⇌H2 + 2OH – –0.828 OH– | H2 |Pt Zn2+ + 2e– ⇌Zn –0.763 Zn2+ | Zn Fe2+ + 2e– ⇌Fe –0.44 Fe2+ | Fe Cd2+ + 2e– ⇌Cd –0.40 Cd2+ | Cd Ag(CN)2– + e– ⇌Ag + 2CN– –0.31 Ag(CN)2 – , CN– | Ag Cu(OH)2 + 2e – ⇌Cu + 2OH– –0.224 OH–, Cu(OH)2 | Cu AgI + e– ⇌Ag + I– –0.151 I– | AgI | Ag Sn2+ + 2e– ⇌Sn –0.136 Sn2+ | Sn Pb2+ + 2e– ⇌Pb –0.126 Pb2+ | Pb Cu(NH3)4 2+ + 2e– ⇌Cu + 4NH3 –0.12 Cu(NH3)4 2+ , NH3 | Cu Fe3+ + 3e– ⇌Fe –0.036 Fe3+ | Fe 2H+ + 2e– ⇌H2 0.000 H + | H2 | Pt AgBr + e– ⇌Ag + Br– +0.095 Br– | AgBr | Ag Cu2+ + e– ⇌Cu+ +0.153 Cu2+, Cu+ | Pt Sn4+ + 2e– ⇌Sn2+ +0.15 Sn4+, Sn2+ | Pt AgCl + e– ⇌Ag + Cl– +0.222 Cl– | AgCl | Ag Hg2Cl2 + 2e – ⇌2Hg + 2Cl– +0.2676 Cl– | Hg2Cl2 | Hg(Pt) Cu2+ + 2e– ⇌Cu +0.337 Cu+ | Cu Ag(NH3)2 + + e– ⇌Ag + 2NH3 +0.373 Ag(NH3)2 + , NH3 | Ag O2(g) + 2H2O + 4e – ⇌4OH– +0.401 OH– | O2 | Pt Cu+ + e– ⇌Cu +0.521 Cu2+ | Cu ½ I2 + e – ⇌I– +0.535 I2, I – | Pt Fe3+ + e– ⇌Fe2+ +0.771 Fe3+, Fe2+ | Pt ½ Hg2 2+ + e– ⇌Hg +0.789 Hg2 2+ | Hg(Pt) Ag+ + e– ⇌Ag +0.799 Ag+ | Ag O2 + 4H + + 4e– ⇌2H2O +1.229 H + | O2 | Pt Br2(ℓ) + 2e – ⇌2Br– +1.0562 Br–, Br2 | Pt Cr2O7 2– + 14H+ + 6e⇌2Cr3+ + 7H2O +1.33 Cr2O7 2-, Cr3+, H+ | Pt ½ Cl2(g) + e – ⇌Cl– +1.359 Cl– | Cl2 | Pt Ce4+ + e– ⇌Ce3+ (1mol dm3 H2SO4) +1.44 Ce 4+ , Ce3+ | Pt Au3+ + 3e– ⇌ Au +1.455 Au3+ | Au MnO4 – + 8H+ + 5e ⇌Mn2+ + 4H2O +1.51 MnO4, Mn 2+ , H+ | Pt Co3+ + e– ⇌Co2+ +1.82 Co3+, Co2+ | Pt
4.4 APPLICATION OF ELECTROCHEMICAL SERIES:
4.4.1 REACTIVITY OF METALS: The activity of the metal depends on its tendency to lose electrons. This tendency depends on the magnitude of SRP (Standard Reduction Potential). The metal which has high –ve value (or smaller +ve value) of SRP, readily loses the electrons and converted into cations. Such a metal is said to be chemically active. The chemical reactivity of metals decreases from top to bottom in the series. For example:
(a) Alkali metals and alkaline earth metals having high –ve values of SRP are chemically active. These react with cold water and evolve H2 gas. These readily dissolve in acids forming corresponding salts and combine with those substances which accept electrons.
(b) Metals like Fe, Pb, Sn, Ni, Co etc., which lie little down in series do not react with cold water but react with steam to evolve hydrogen.
(c) Metals like Cu, Ag, An etc. which lie below hydrogen are less reactive and do not evolve hydrogen from water.
4.4.2 ELECTROPOSITIVE CHARACTER OF METALS:The electropositive character also depends on the
tendency to lose electrons. Like reactivity, the electropositive character of metal decreases from top to bottom. On the basis of SRP values, metals are divided into three groups.
(A) Strongly electropositive metals (having SRP near about –2.0V or more negative) (B) Moderately electropositive metals (having SRP between 0.0 and –0.2V)
(C) Weakly electropositive metals (lie below hydrogen and having +ve SRP)
4.4.3 DISPLACEMENT REACTIONS:
(i) A metal higher in the series (lower SRP) displace the metal from its salt solution which is lower in the series (higher SRP)
(ii) A non metal having high value of SRP will displace another non metal with lower SRP (i.e. occupying position above in the series).
Ex: Fe + Cu2+ →Fe2+ + Cu Cl2 + 2I
– →
2Cl– + I2
5.
E
LECTROCHEMICAL ORG
ALVANICC
ELLS:
Electrochemical or Galvanic cell is a device for converting chemical energy into electrical energy. The electromotive force (EMF) of such cell is directly proportional to the intensity of chemical reaction taking place in it. The chemical reaction responsible for production of electricity normally takes place in two separate compartments. These compartments are called half cells. Each half cell consist electrode and electrolytic solution. When the two compartments are connected by a “salt bridge” and electrodes are joined by a wire through galvanometer, the electricity begins to flow. One the example of Galvanic cell is Daniell Cell.5.1 DANIELL CELL:
Salt bridge Copper(II)
sulphate solution (CuSO4) Copper electrode e– Zinc electrode Zinc sulphate Solution (ZnSO4) e– A
(i) When zinc and copper electrodes are joined by a wire the following observations are made: (a) There is a flow of current through the external circuit.
(c) The concentration of ZnSO4 solution increases, while the concentration of CuSO4 solution decreases.
(d) The solution in both the compartments remains electrically neutral. (ii) The indirect redox reaction take place in the cell will be:
Anode or LHS compartment: Zn →Zn2+ (aq) + 2e– Cathode or RHS compartment: Cu2+ (aq) + 2e– →Cu Net cell reaction is: Zn + Cu2+(aq) →Zn2+(aq) + Cu
• The half cell in which oxidation occurs is called oxidation half cell and the half cell in which reduction occurs is called reduction half cell.
• The electrode where oxidation takes place is called Anode and where reduction occurs is termed Cathode.
• During the passage of electric current through external circuit, electrons flow from the anode to cathode and their polarity is negative and positive, respectively.
• Chemical energy is converted into electrical energy. • The net cell reaction is the sum of two cell reactions.
5.2 SALT BRIDGE AND ITS SIGNIFICANCE: Salt bridge is usually an inverted U-tube filled with concentrated solution of inert electrolyte (viz. KCl, KNO3, NH4NO3 etc.). Inert electrolyte is one whose ions are neither involved in any electrochemical change nor do the react chemically with the electrolytes in the two half cells. Gelatin or Agar-Agar (Plant Gel), is dissolved in a hot concentrated aqueous solution of inert electrolyte and the solution thus formed is filled in the U-tube. The ends of the U-tube are plugged with cotton wool as to minimize diffusion. The following are the functions of the salt bridge:
(a) It connects the solutions of two half cells and completes the cell circuit.
(b) It prevents transference or diffusion of the solutions from one half cell to the other. (c) It keeps the solutions in two half cells electrically neutral.
(d) It prevents liquid-liquid junction potential, i.e., the potential difference which arises between two solutions when in contact with each other.
• Salt bridge can be replaced by a porous partition which allows the migration of ions without allowing the solutions to intermix.
5.3 REPRESENTATION OF GALVANIC CELL:The two half cells required to form an electrochemical cell can be represented by following rules:
(a) Reduction half cell is represented in right and oxidation half cell, in left. (b) The separation of two phases (state of matter) is shown by a vertical line.
(c) The various materials present in the same phase are shown together with the help of commas. (d) The two half cells are joined with the help of double vertical lines.
(e) The significant features of the substances viz. fugacity (pressure) of gas, activity (concentration) of ion etc. are indicated in brackets drawn immediately after writing the substance.
For example: In Daniell cell, the cell reaction is Zn + Cu2+→ Zn2+ + Cu. The cell diagram is Zn | Zn2+ | | Cu2+ | Cu
5.4 ELECTROMOTIVE FORCE OR EMF: The difference in potential which causes a current to flow from the electrode of higher potential to the one of lower potential is called EMF of the cell. It can be represented as
Ecell = Higher reduction potential – Lower reduction potential Or E0cell = SRP of Cathode – SRP of Anode
Or E0cell = SOP of Anode + SRP of Cathode Or E0cell = SOP of Anode – SOP of Cathode
5.5 FREE ENERGY CHANGE IN GALVANIC CELLS:When a cell operates, work is done by the system, as electrical energy flows through external circuit. The quantity of electrical work done is measured in terms of free energy change, ∆G. This value is equal to the product of number of moles of electrons and emf of the cell.
∆G = –nFEcell
If reactants and products are in the standard states, ∆∆∆∆G°°°° = –nFE°°°°cell
5.6 NERNST EQUATION: The electrode potential and EMF of the cell depend upon the nature of the electrode, temperature and the activities (concentration) of the ions in solution. The dependence of electrode and cell potential on activities of the ions in solution can be obtained from thermodynamic.
∆G = ∆G° + RT ln Q
or, –nFEcell = –nF Ecell + RT ln Q or, Ecell Ecell RTln Q
nF = − Nernst equation At 25°C, ECell = Ecell – 0.059 log Q n
(iii) ∆G° refers to free energy change for the reaction when the various reactants and products are present at standard conditions. (298K is the temperature, 1M is the concentration and 1 atm. is the pressure).
6. COMMERCIAL CELLS: Some of the electrochemical cells which are commonly employed as source of electrical energy can be classified as follows:
6.1 THE PRIMARY CELLS:Such cells can be used only so long the reactive materials are present. Once they exhausted, the cell can not be recharged by passing current through it.
Example: Leclanche or dry cell: represented as: Zn | NH4Cl (20%), ZnCl2 | MnO2 | Graphite The reactions involved are:
Anode: Zn(s) →Zn+ + 2e
Cathode: 2MnO2(s) + H2O(l)+ 2e →Mn2O3(s) + 2OH –
The OH–- further causes the following secondary reactions: NH4Cl + OH – → NH3 + Cl – + H2O Zn2+ + 2NH3 + 2Cl – → Zn(NH3)2Cl2 The EMF of the cell is 1.5 Volts
6.2 THE SECONDARY CELLS:Such type of cells can be used again and again by recharging the cell. One of the examples is the Lead storage battery. In lead storage battery Pb acts as anode and lead impregnated with lead oxide (PbO2) acts as cathode. The electrolyte is a nearly 37% H2SO4 solution with specific gravity of about 1.15 at 298K. The reactions are:
Anode: Pb(s) + SO24−(aq) →PbSO4(s) + 2e Cathode: PbO2(s) + 4H
+
(aq.) + 2e →PbSO4(s) + 2H2O(l)
The EMF of the cell depends on the concentration of H2SO4 viz. 1.9V at 7.4%, 2.0V at 21.4% and 2.14V at 39.2% of H2SO4. To recharge the cell, it is connected with a source of potential higher than that of the cell in such a way, that the cell now acts as the electrolytic cell, Pb is deposited on the cathode, PbO2 is formed at the anode and H2SO4 is regenerated.
6.3 FUEL CELL:Fuel cells are electrochemical devices, which convert the energy of fuel by oxidation reaction into electrical energy. A fundamental example of fuel cell is the Hydrogen – Oxygen cell.
H2 O2 O2 H2 NaOH Solution Fuel cell Θ Θ Ammeter
It consists of two electrodes made of porous graphite. Platinum is coated on the surface of the electrodes. The electrodes are placed in aqueous solution of KOH or NaOH. H2 and O2 are bubbled into the cell under 50 atm pressure. Electrods reactions are:
Anode: H2 + 2OH – →
2H2O + 2e Cathode: ½ O2 + H2O + 2e →2OH
–
ELECTRODE POTENTIAL
H
ASEEN
S
HURUAT
(DPP–I)1. A metal having negative reduction potential when dipped in the solution of its own ions has a tendency to
(A) remain neutral (B) become electrically positive (C) go into the solution (D) be deposited from the solution 2. A metal rod is dipped in a solution of its ions. Its electrode potential is independent of
(A) temperature of the solution (B) concentration of the solution (C) area of the metal exposed (D) nature of the metal
3. Indicator electrode is
(A) SHE (B) Calomel electrode
(C) Ag/AgCl electrode (D) Quinhydrone electrode
4. The following reactions represent the reduction of IO3− ion into I− ion in acidic and basic medium. Predict in which medium IO3− ion will act as a better oxidising egent?
IO3− + 6 H+ + 6 e → I− + 3H2O; Eº = + 0.907 V IO3− + 3H2O + 6e → I−+ 6 OH¯ ; Eº = + 0.260 V
(A) Acid medium (B) Basic medium (C) Equally in both (D) none of these 5. The position of some metals in the electrochemical series in decreasing electropositive character is
given as: Mg > Al > Zn > Cu > Ag. What will happen if a copper spoon is used to stir a solution of aluminium nitrate?
(A) the spoon will get coated with aluminium (B) an alloy of copper and aluminium is formed (C) the solution become blue
(D) there is no reaction
6. The metal that can not be obtained on reduction of its oxide by Al is
(A) K (B) Mn (C) Cr (D) Fe
7. The standard electrode potentials of four elements A, B, C and D are -3.05, -1.66, -0.40 and +0.80 volts. The highest chemical reactivity will be shown by
(A) A (B) B (C) C (D) D
8. The calomel electrode is reversible with respect to (A) Hg2
2+
(B) H+ (C) Hg2+ (D) Cl–
9. Which one of the following does not get oxidised by bromine water? (A) Fe2+ to Fe3+ (B) Cu+ to Cu2+ (C) Mn2+ to MnO
4
– (D) Sn2+ to Sn4+
10. The standard electrode potential values of the elements A, B and C are 0.68, –2.50 and –0.50V respectively. The order of their reducing power is
P
ATHRILA
R
ASTA
(DPP–II)1. We have an oxidation – reduction system: [Fe (CN)6]3– + e ⇌[Fe(CN) 6]
4– ; Eº = + 0.36 V. At what ratio of the concentration of oxidised and reduced from will the potential of the system be 0.28 V?
(A) 22.7 : 1 (B) 1 : 22.7 (C) 1 : 3.9 (D) 1 : 15
2. The standard reduction potential of oxygen in acidic solution is 1.23V (O2 + 4 H3O+ + 4e → 6H2O). What is the standard reduction potential of oxygen in basic solution?
(A) +0.404 V (B) -0.404 V (C) +2.056 V (D) -2.056 V 3. The electrode potential of hydrogen electrode in neutral solution and 298K is
(A) - 0.41 V (B) zero (C) - 0.49 V (D) - 0.82 V
4. The standard reduction potential for the reaction [Co(H2O)6]3+ + e → [Co(H 2O)6]
2+ is about 1.8V. The standard reduction potential for the reaction [Co(NH3)6]3+ + e → [Co(NH
3)6]
2+ is 0.1V. Which of the complex ion, [Co(H2O)6]2+ or [Co(NH
3)6]
2+ can be oxidised to the corresponding cobalt (III) complex, by oxygen.
2
O / OH
E −= 0.4V.
(A) [Co(H2O)6]2+ (B) [Co(NH 3)6]
2+ (C) both (D) none of these 5. The standard potentials of OCl–/Cl– and Cl–/Cl
2are 0.94V and 1.36V respectively. The Eº value of OCl–/Cl
2 will be
(A) 1.04 V (B) – 0.42 V (C) –2.20 V (D) 0.52 V
6. Calculate the degree of hydrolysis and hydrolysis constant of aniline hydrochloride in M/32 solution at 25ºC. Given:
6 5 2 2 0
C H NH HCl|H (1atm)
E = −0.188 V
7. Calculate the reduction potential at 25ºC for Fe3+/Fe2+ electrode if the concentration of Fe2+ ion is five times that of Fe3+ ion.
3 2
0 Fe |Fe
E + += 0.77 volt.
8. Calculate electrode potential for the half cell Pt | H2 (1 atm) | 0.357M-CH3COOH. Ka for CH3COOH = 1.74 × 10–5.
9. Estimate the standard reduction potential of CuS/Cu electrode. Ksp for CuS = 8.5 × 10–36 and 2
0 Cu |Cu E + =
+ 0.34 volt.
10. What is the standard electrode potential for the electrode, MnO4−/MnO2 in acid solution? – 2 4 MnO | Mn Eº + = 1.51V, 2 2 MnO / Mn Eº + = 1.23V.
GALVANIC CELL
H
ASEEN
S
HURUAT
(DPP–III)1. The e.m.f. of a galvanic cell composed of two hydrogen electrodes is 272 mV. What is the PH of the solution in which the anode is immersed if the cathode is in contact with a solution of PH = 3?
(A) 3 (B) 6.7 (C) 7.6 (D) 1.6
2. For a cell reaction involving a two-electron change the standard emf of the cell is found to be 0.295V at 25ºC. The equilibrium constant of the reaction at 25ºC will be
(A) 1 × 1010 (B) 1 × 10–10 (C) 29.5 × 10– 2 (D) 10
3. Tell which of the following statements accurately describes the effect of adding CN- to the cathode of a cell with a cell reaction: Cd + 2Ag+ → 2Ag + Cd2+, Eº = 1.2V
(A) Eº increases because Cd(CN)4 2−
forms (B) Eº decreases because Cd(CN)4 2−
forms (C) Eº increases because Ag(CN)2− forms (D) Eº decreases because Ag(CN)2− forms 4. The ECell for Ag(s) | AgI (satd) || Ag+ (0.10M) | Ag (s) is + 0.417 V. What is Ksp of AgI?
(A) 2.7 × 10-8 (B) 7.32 × 10-13 (C) 8.55 × 10-9 (D) 7.32 × 10-17 5. Assuming that hydrogen behaves as an ideal gas, what is the e.m.f. of the cell at 25ºC if P1 = 640 mm
and P2 = 425 mm: Pt | H2 (P1) | HCl | H2 (P2) | Pt?
(A) -0.005 V (B) -0.01 V (C) +0.005 V (D) +0.01 V
6. A depolariser used in dry cell batteries is
(A) NH4Cl (B) MnO2 (C) KOH (D) Na3PO4
7. For a electrochemical cell Zn | Zn2+ (C1 M) || Cu2+ (C2M) | Cu, the decrease in free energy at a given temperature is a function of
(A) ln C1 (B) ln C2 (C) ln C2/C1 (D) ln C1/C2
8. At what PH does the potential (emf) for the disproportionation of chlorine change from a negative value to a positive value, assuming 1.0 atm pressure and 1.0 M – Concentration for all species except hydrogen ion? Given:
Cl2 + 2e → 2 Cl–; Eº = 1.36 volt OCl– + 2H+ + e → Cl
2 + H2O; Eº = 1.63 volt.
(A) 4.58 (B) 1.14 (C) 2.29 (D) 9.15
9. The emf of the cell at 298K Hg(l) | Hg2Cl2 (s), KCl sol. (1.0 N) | Quinohydrone | Pt, is 0.212V. What is the PH of the quinohydrone solution, the potential of the normal calomel electrode is 0.2812V and Eº for the quinohydrone electrode is 0.6996V, both at the same temperature.
(A) 3.5 (B) 7.0 (C) nearly zero (D) nearly 14
10. The potential of a cell consisting of an anode of silver in 0.10M-AgNO3 solution and a cathode of Pt immersed in a solution of 1.5M-Cr2O72¯ , 0.75M-Cr3+ and 0.25M-H+ is (E0
Ag+|Ag = 0.80V and E0 Cr2O72¯ |Cr 3+ = 1.33 V) (A) 0.41 V (B) 0.61 V (C) 0.51 V (D) 0.71 V
11. Two electrochemical cells are assembled in which the following reactions occur: V2+ + VO2+ + 2H+→ 2V3+ + H 2O; EºCell = 0.616 V V3+ + Ag+ + H 2O → VO 2+ + Ag (s) + 2H+; Eº Cell = 0.439 V If EºAg+
|Ag = 0.799 V, what is EºV3+|V2+?
(A) -0.256 V (B) +0.256 V (C) +1.854 V (D) -1.854 V
12. The emf of the cell: H2 (g) | Buffer || Normal calomel Electrode is 0.6885 V at 25º C, when barometric pressure is 760 mm. What is the PH of the buffer solution? Eº
Calomel = 0.28 V.
(A) 3.46 (B) 6.92 (C) tending to zero (D) tending to 14
13. The emf of the cell obtained by combining Zn and Cu electrodes of a Daniel cell with normal calomel elctrodes are 1.083V and – 0.018V, respectively at 25º C. If the potential of normal calomel electrode is – 0.28V, find emf of Daniel cell.
(A) 1.065 V (B) 1.101 V (C) 0.803 V (D) 0.262 V
14. From the following E° values for the half cells
(i) D → D2+ + 2e–; E° = –1.5 V (ii) B+ + e– → B; E° = – 0.5 V
(iii) A3–→ A2– + e–; E° = 1.5 V (iv) C2+ + e– → C+; E° = + 0.5 V
Which combination of two half cells would result in a cell with largest potential?
(A) i and iii (B) i and iv (C) iii and iv (D) ii and iv
15. A Tl+/Tl couple was prepared by saturating 0.10M-KBr with TlBr and allowing Tl+ ions form the insoluble bromide to equilibrate. This couple was observed to have a potential –0.443V with respect to Pb2+/Pb couple in which Pb2+ was 0.10M. What is the Ksp of TlBr. Given 2
0 Pb |Pb
E + : = – 0.126V
[Antilog(-4.44) = 3.6 ×10–5, Antilog(-4.54) = 2.88 ×10–5]
P
ATHRILA
R
ASTA
(DPP–IV)1. Write the cell diagram for each of the following reactions and calculate Eº for each cell. (a) AgBr (s) + 1/2 H2 (g) →Ag(s) + H+ + Br–
E0AgBr|Ag|Br -= + 0.10 volt. (b) PbSO4 + 2e →Pb(s) + SO4 2− ; Eº = – 0.359 volt Pb2+ + 2e →Pb(s); Eº = – 0.126 volt
(c) AgI (s) + e →Ag(s) + I–; Eº = – 0.152 volt Ag+ + e →Ag (s); Eº = + 0.80 volt
(d) Fe3+ (0.1 M) + e →Fe2+ (1M); Eº = 0.77 volt
AgCl (s) + e →Ag (s) + Cl– (0.001 M); Eº = 0.22 volt
2. The standard reduction potential at 25ºC of the reaction 2H2O + 2e– ⇌H2 + 2OH– is – 0.8277 volt. Calculate equilibrium constant for the reaction 2H2O ⇌H3O+ + OH– at 25ºC.
3. Calculate [Fe3+] at equilibrium when potassium iodide is added to a solution of Fe3+ initially 0.50 M until [I–] = 1.0 M. Given 3 2 0 Fe |Fe E + += 0.770 V, 2 0 I |I E − = 0.535 V.
4. If it is desired to construct the following voltaic cell to have ECell = 0.086V, what [Cl–] must be present in the cathodic half cell to achieve the desired emf. Given Ksp of AgCl and AgI are 1.8 × 10–10 and 8.5 ×10–17, respectively. Ag(s) | Ag+ (sat. AgI) || Ag+ (sat. AgCl, x M- Cl–) | Ag (s)
5. The voltage of the cell given below is – 0.46V.
Pt | H2 | NaHSO3 (0.4 M), Na2SO3 (6.44 × 10–3 M) || Zn2+ (0.3 M) | Zn If EºZn
2+
|Zn = – 0.763V, Calculate Ka2 of H2SO4.
6. For the reaction, H2(g) + 2AgCl(s) + 2H2O(ℓ)→2Ag(s) + 2H3O+(aq) + 2Cl–(aq) at 25ºC, the standard free energy of formation of AgCl(s), H2O(ℓ) and (H3O+ + Cl–) (aq) are –109.7, –237.2 and –368.4 KJ/mol. Calculate what will be the cell voltage if this reaction is run at 25ºC and 1 atm in a cell in which H2 activity is unity and H3O+ and Cl– activities are each at 0.01M?
7. An alloy weighing 1.05g of Pb-Ag was dissolved in desired amount of HNO3 and volume was made 350 ml. An silver electrode was dipped in solution and Ecell of the cell Pt, H2(1atm) | H+(1M) ||Ag+ |Ag was 0.503V at 298K. Calculate the percentange of lead in alloy. Given EºAg+|Ag = 0.80V.
8. Estimate the cell potential of a Daniel cell having 1.0M-Zn2+ and originally having 1.0M-Cu2+ after
sufficient ammonia has been added to the cathode compartment to make the NH3 concentration 2.0M. Given: EºZn 2+ |Zn= – 0.76 V, EºCu 2+ |Cu= + 0.34 V, Kf for Cu (NH3)4 2+ = 1 × 1012.
9. Calculate the minimum mass of NaOH required to be added in RHS to consume all the H+ present in RHS of the cell of emf + 0.701V at 25ºC, before its use. Also report the emf of cell after addition of NaOH. Zn | Zn2+ (0.1M) || HCl | H
2 (1 atm), Pt; EºZn 2+
|Zn = – 0.76 V
10. The EMF of the cell: Ag, AgCl in 0.1M–KCl || satd. NH4NO3 || 0.1M–AgNO3, Ag is 0.45 volt at 25ºC. Calculate solubility product of AgCl. 0.1M–KCl is 85% dissociated and 0.1M–AgNO3 is 82% dissociated.
11. Given E0 = 0.08V for the Fe3+ (cyt b) | Fe2+ (cyt b) couple and E0
= 0.22V for the Fe3+ (cyb c 1) | Fe
2+ (cyt c1) couple, where E0 represents the standard state reduction potentials at PH = 7.0 at 25ºC and cyt is an abbreviation for cytochromes. Determine E0 and K for the equation
Fe3+ (cyt c 1) + Fe
2+ (cyt b) ⇌Fe2+ (cyt c 1) + Fe
3+ (cyt b)
12. When metallic copper is shaken with a solution of a copper salt, the reaction Cu + Cu2+⇌2Cu+
proceeds. When equilibrium is established at 298K, [Cu2+]/[Cu+]2 = 1.66 ×106M–1. If the standard
potential of the Cu2+ | Cu half cell is + 0.337V, what is the standard potential of Cu+| Cu half cell?
13. When silver chloride is dissolved in a large excess of ammonia, practically all silver ion can be assumed to exist in form of a single ionic species [Agx(NH3)y]
x+
. Compute the values of x and y using the two following cells:
(A) Ag | 0.4 x 10-3 M AgCl, 1M-NH3 || 40 x 10 -3
M AgCl, 1M-NH3 | Ag; Ecell = 0.1185V at 298K. (B) Ag | 3 x 10-3 M AgCl, 1M-NH3 || 3 x 10
-3
M AgCl, 0.1M-NH3 | Ag; Ecell = 0.1263V at 298K.
14. For the reaction: H2 (1atm) + 2AgCl(s) ⇌2Ag(s) + 2H+ (0.1M) + 2Cl– (0.1M); ∆Gº = – 42927 J at 25ºC. Calculate the e.m.f. of the cell in which the given reaction takes place.
15. The emf of the cell: Cd(s) | CdCl2.5H2O (sat.) || AgCl(s) | Ag(s) for the reaction is Cd(s) + 2AgCl(s) + aq → CdCl2.5H2O (sat.) + 2Ag(s)
is +0.6915V at 0ºC and +0.6753V at 25ºC. Calculate the free energy change, enthalpy change and entropy change at 25ºC.
ELECTROLYSIS
H
ASEEN
S
HURUAT
(DPP–V)1. Electrochemical equivalent is more for
(A) Hydrogen (B) Silver (C) Copper (D) Zinc
2. A current of 3.7 ampere is passed for 6 hours between Nickel electrodes in 0.50 litre of 2M solution of Ni(NO3)2 . The molarity of Ni2+ at the end of electrolysis is
(A) 1.172 M (B) 0.172 M (C) 0.586 M (D) 2 M
3. The elctrochemical equivalent of two substances are E1 and E2 The current that must pass to deposit the same amount at the cathodes in the same time must be in the ratio of
(A) E1 : E2 (B) E2 : E1 (C) E1 − E2 : E2 (D) E1 : E2 − E1 4. The same quantity of electricity is passed through one molar solution of H2SO4 and one molar solution
of HCl. The amount of hydrogen evolved from H2SO4 as compared to that from HCl is
(A) the same (B) twice as such
(C) one half as such (D) dependent on concentration
5. Two electrolytic cells, one containing acidified ferrous chloride and another acidified ferric chloride are connected in series. The ratio of iron deposited at cathodes in the two cells will be
(A) 3 : 1 (B) 2 : 3 (C) 1 : 1 (D) 3 : 2
6. A galvanic cell is set up from a zinc bar weighing 100g and 1.0L of 1.0M copper sulphate solution. How long would the cell run if it is assumed to deliver a steady current of 1.0 ampere?
(A) 53.6 hr (B) 26.8 hr (C) 20.5 hr (D) 41.0 hr
7. Two platinum electrodes were immersed in a solution of CuSO4 and electric current was passed through the solution. After some time, it was found that colour of CuSO4 disappeared with the evolution of gas at the electrode. The colourless solution contains
(A) platinum sulphate (B) copper sulphate (C) copper hydroxide (D) sulphuric acid 8. The amount Faradays required to liberate 1 mole of an element indicates
(A) weight of element (B) conductance of electrolyte (C) charge on the ion of that element (D) chemical equivalent
9. An ion is a reduced to the element when it absorbs 6×1020 electrons. The number of equivalents of the
ion is
(A) 0.10 (B) 0.01 (C) 0.001 (D) 0.0001
10. A solution containing one mole per litre each of Cu(NO3)2, Mg(NO3)2, AgNO3, Hg(NO3)2 is being electrolysed using inert electrodes. The values of standard electrode potential in volts are Ag+ | Ag =
0.80V, Hg2+ | Hg = 0.79V, Cu2+ | Cu = 0.34V, Mg2+ | Mg = – 2.37V. With increasing voltage the
sequence of deposit of metals on the cathode will be
(A) Ag, Hg, Cu, Mg (B) Mg, Cu, Hg, Ag (C) Ag, Mg, Cu (D) Cu, Hg, Ag 11. An ammeter and a copper voltameter are connected in series in an electric circuit through which a
constant direct current flows. The ammeter shows 0.525 ampere. If 0.635 gm of Cu is deposited in one hour, what is the percentage error of ammeter?
12. In the lead storage battery the anode reaction is Pb(s) + HSO4¯ + H2O → PbSO4(s) + H3O+ + 2e. How many gm of Pb will be used upto deliver 1 amp for 100 hrs? (Pb = 208)
(A) 776 gm (B) 388 gm (C) 194 gm (D) 0.1 gm
13. Element A (at. wt. 112) and element B (at. wt. = 27) form chlorides. Solutions of these chlorides are electrolysed seperately and it is found that when the same quantity of electricity is passed, 5.6 gm of A was deposited while only 0.9gm of B was deposited. The valency of B is 3.The valency of A is
(A) 1 (B) 2 (C) 3 (D) 4
14. During the electrolysis of an aqueous salt solution, the PH in the space near one of the electrode was increased and the other one was decreased. The salt solution was
(A) NaCl (very dilute) (B) ZnCl2 (C) NaCl (Conc.) (D) Cu(NO3)2 15. A hydrogen electrode placed in a buffer solution of CH3COONa and CH3COOH in the ratio’s x : y
and y : x has electrode potential values E1 and E2 volts, respectively at 25 0
C. The PKa value of acetic acid is
(A) (E1 + E2)/ 0.118 (B) (E2 – E1)/ 0.118 (C) -(E1 + E2)/ 0.118 (D) (E1 – E2)/ 0.118 16. 100 ml of a buffer of 1M- NH3 and 1M- NH4
+
are placed in two voltaic cells separately. A current of 1.5 amp is passed through both cells for 20 min. If electrolysis of water takes place only and the electrode reactions are: Right: 2H2O + O2 + 4e → 4 OH
and Left: 2H2O → 4H +
+ O2 + 4e, then the PH of the
(A) right electrode will increase (B) left electrode will increase (C) both electrode will increase (D) both electrode will decrease
17. The same current was passed successively through solution of zinc-ammonium sulphate and nickel-ammonium sulphate rendered alkaline with ammonia. The weight of zinc and nickel deposited in a certain time were found to be 22.89 gm and 20.55 gm, respectively. Given that the chemical equivalent weight of zinc is 32.7, what is the chemical equivalent weight of nickel?
(A) 58.71 (B) 29.36 (C) 14.39 (D) 36.42
18. The preparation of LiOH by the electrolysis of a 35% solution of LiCl using a platinum anode led to a current efficiency of 95%. What weight of LiOH was formed by the passage of 2.68A for 1 hr?
(A) 2.28 gm (B) 2.40 gm (C) 0.66 gm (D) 2.53 gm
19. Assuming that copper contains only iron, silver and gold as impurities. After passage of 140 ampere for 482.5 seconds, the mass of anode decreased by 22.26 gm and the cathode increased by 22.011gm. The percentage of iron and copper in the original sample are respectively
(A) 0.85%, 98.88% (B) 0.01%, 98.88% (C) 0.1%, 98.88% (D) 98.88%, 0.01% 20. After electrolysis of an aqueous sodium chloride solution, it was found that the solution is being
neutralised by 60 ml N – HCl solution. During the same period of electrolysis, 3.18 gm of copper was deposited in a copper voltameter in series. What is the percentage of the theoretical yield of sodium hydroxide obtained?
P
ATHRILA
R
ASTA
(DPP–VI)1. 100 ml of 0.6N-CuSO4 solution is electrolysed between two platinum electrodes till the concentration in the residual liquid is 0.1N, when a steady current of 5.0 amp is used. How long should the current be passed to get the above change?
2. An object whose surface area is 78.6 cm2 is to be plated with an even layer of gold 8.0 × 10–4 cm thick. The density of gold is 19.3 gm/ml. The object is placed in a solution of Au(NO3)3 and a current of 2.75 A is applied. Find the time required for the electroplating to be completed, assuming that the layer of gold builds up evenly. (Au = 197)
3. The following galvanic cell
Zn | Zn (NO3)2 (aq) || Cu (NO3)2 (aq) | Cu (100 ml, 1M) (100 ml, 1M)
was operated as an electrolytic cell as Cu as the anode and Zn as the cathode. A current of 0.48 ampere was passed for 10 hours and then the cell was allowed to function as galvanic cell. What would be the e.m.f. of the cell at 25ºC? Assume that the only electrode reactions occuring were those involving Cu | Cu2+ and Zn | Zn2+. Eº
Cu 2+
|Cu = + 0.34V, EºZn 2+
|Zn = – 0.76V.
4. An alloy of lead (valency = 2)-thallium (valency = 1) containing 70% Pb and 30% Tl by weight can be electroplated onto a cathode from a perchloric acid solution. How many hours would be required to deposit 5.0 gm of this alloy at a current of 1.10 amp? (Pb = 208, Tl = 204)
5. Iridium was plated from a solution containing IrCl6y for 2.0 hour with a current of 0.075 amp. The Ir
deposited on the cathode weighed 0.359 gm (a) what is the oxidation number of Ir in IrCl6y? (b) what
is the charge, y, on this ion? (Ir = 192)
6. The electrolysis of cold sodium chloride solution produces sodium hypochlorite by reacting NaOH and Cl2 throughly. How long will a cell operate to produce 1000 litre of 5.25% (by wt.) solution of NaClO if the cell current is 2.50 × 103 ampere? Assume that the density of solution is 1.0 gm/ml.
7. At the Nangal fertilizer plant in Punjab, hydrogen is produced by the elctrolysis of water. The hydrogen is used for the production of ammonia and nitric acid (by the oxidation of ammonia). If the average production of ammonium nitrate is 5000 Kg/day, estimate the daily consumption of electricity per day.
8. 50 ml of 0.1M-CuSO4 solution is electrolysed for 12 min. at a current of 0.06 amp. If Cu is produced at one electrode and oxygen at the other, what will be the PH of the final solution? For HSO4
-, Ka =
1.3×10–2.
9. To reduce nitrobezene to aniline, 20 gm of C6H5NO2, 30 ml of an alcohol, 250 ml of water, 11 gm of HCl and 1gm of SnCl2.2H2O were placed in the cathode space. After passing current at a rate of 26.5 amp-hour through the lead cathode electrolytic cell, 12.76 gm of aniline was produced. Determine the current yield.
10. A test for complete removal of Cu2+ ions from a solution of Cu2+ is to add NH
3(aq). A blue colour signifies the formation of complex [Cu(NH3)4]2+ having Kf = 1.1×1013 and thus confirms the presence of Cu2+ in solution. 250 ml of 0.1M-CuSO
4 is electrolysed by passing a current of 3.512 ampere for 1368 second. After passage of this charge, sufficient quantity of NH3 is added to electrolysed solution maintaining [NH3] = 0.10 M. If [Cu(NH3)4]2+ is detectable upto its concentration as low as 1×10–5M, would a blue colour be shown by the electrolysed solution on addition of NH3.
11. Perdisulphuric acid, H2S2O8 can be prepared by electrolytic oxidation of H2SO4 as 2 H2SO4→ H2S2O8 + 2H+ + 2e. Oxygen and hydrogen are by products. In such an electrolysis, 0.72L of H
2 and 2.35L of O2 were generated at STP. What is the weight of H2S2O8 formed?
12. The electrode reactions for charging of a lead storage battery are: PbSO4 + 2e → Pb + SO42–
PbSO4 + 2H2O → PbO2 + SO42– + 4H+ + 2e
The electrolyte in the battery is an aqueous solution of sulphuric acid. Before charging, the specific gravity of the liquid was found to be 1.11 (15.7% H2SO4 by wt.). After charging for 100 hours, the specific gravity of the liquid was found to be 1.28 (36.9% H2SO4 by wt). If the battery contained 2 litres of the liquid, calculate the average current used for charging the battery.
13. By passing a certain amount of charge through NaCl solution, 9.2 litre of chlorine were liberated at STP. When the same amount of charge is passed through a nitrate solution of metal M, 7.467 gm of the metal was deposited. If the specific heat of metal is 0.216 Cal/ºC-gm, what is the formula of metal
nitrate? 14.
Two litre solution of a buffer mixture containing 1.0M-NaH2PO4 and 1.0M-Na2HPO4 is placed in two compartments (one litre in each) of an electrolytic cell. The platinum electrodes are inserted in each compartment and 1.25 ampere current is passed for 212 min. Assuming electrolysis of water only at each compartment, what will be PH in each compartment after passage of
above charge ? PKa for H
2PO4– = 2.15.
15. H2O2 can be prepared by successive reactions: 2 NH4HSO4 → H2 + (NH4)2 S2O8
(NH4)2 S2O8 + 2H2O → 2 NH4HSO4 + H2O2
The first reaction is an electrolytic reaction and second is steam distillation. What amount of current would have to be used in first reaction to produce enough intermediate to yield 100 gm pure H2O2 per hour. Assume current efficiency 50%.
CONDUCTANCE
H
ASEEN
S
HURUAT
(DPP–VII)1. Equivalence conductance at infinite dilution of NH4Cl, NaOH and NaCl are 129.8, 217.4 and 108.9 ohm-1, respectively. If the equivalent conductance of 0.01N solution of NH
4OH is 9.33ohm
-1cm2eq-1, then the degree of dissociation of NH4OH at this temperature is
(A) 2.0 % (B) 0.039 % (C) 3.9 % (D) 4.7 %
2. Which of the following solution have highest resistance?
(A) 1N – NaCl (B) 0.1N – NaCl (C) 2N – NaCl (D) 0.05 N - NaCl 3. Which pure substance will not conduct electricity?
(A) Molten NaCl (B) Molten KOH (C) Liquefied HCl (D) Liquid Hg 4. The resistance of a solution A is 50 ohms and that of solution B is 100 ohms, both solutions being
taken in the same conductivity cell. If equal volumes of solution A and B are mixed, what will be the resistance of the mixture using the same cell? Assume that there is no increase in the degree of dissociation of A and B on mixing.
(A) 150 ohm (B) 75 ohm (C) 33.33 ohm (D) 66.67 ohm
5. In a conductivity cell, the two platinum electrodes, each of area 10cm2 are fixed 1.5 cm apart. The cell
contained 0.05N solution of a salt. If the two electrodes are just half dipped into the solution which has a resistance of 50 ohms, what is the equivalent conductance of the salt solution in ohm-1cm2eq-1.
(A) 120 (B) 60 (C) 240 (D) 6000
6. At 18ºC, the mobilities of NH4+ and ClO 4
– ions are 6.6 × 10–4 and 5.4 × 10–4 cm2volt–1sec–1 at infinite
dilution. What is the equivalent conductance of ammonium chromate solution in ohm-1cm2eq-1. (A) 833.3 (B) 115.8 (C) 8.64 x 10-3 (D) 1.24 x 10-8 7. At 25ºC, the equivalent conductance at infinite dilution for HCl solution is 425 ohm-1 cm2 eqv-1, while
its specific conductance is 3.825 ohm-1 cm-1. If the apparent degree of dissociation is 90%, the
normality of the solution is
(A) 0.90 N (B) 1.0 N (C) 1.1 N (D) 1.2 N
8. The degree of dissociation of acetic acid in an aqueous solution of the acid is practically unaffected by (A) adding a piece of sodium chloride
(B) adding a drop of concentrated hydrochloric acid (C) diluting with water
(D) raising the temperature
9. The difference between a decinormal solution of HCl and a decinormal solution of acetic acid is that (A) one of them conducts electricity and the other does not
(B) one of them is corrosive to the skin and the other is not
(C) one of them contains undissociated molecules of the acid and the other does not (D) one of them decompose sodium carbonate and the other does not.
10. The resistance of 1N solution of acetic acid is 250 ohm, when measured in a cell of cell constant 1.15 cm–1. The equivalent conductivity in ohm–1 cm equivalent–1 of 1N acetic acid is
P
ATHRILA
R
ASTA
(DPP–VIII)1. The specific conductance at 298K of AgCl solution in water was determined to be 1.826 x 10-6 ohm-1 cm-1. The ionic conductances at infinite dilution, of Ag+ and Cl¯ are 61.92 and 76.34 ohm-1 cm2 eq-1 respectively. What is the solubility of AgCl in water?
(A) 9.2 x 10−8 g/L (B) 1.32 x 10−5 g/L (C) 1.9 x10−3 g/L (D) 1.76 x 10−10 g/L 2. A cell whose resistance, when filled with 0.1M-KCl is 200 ohm, is measured to be 6400 ohm, when
filled with 0.003M-NaCl solution. All measurements were done at 298K. What is the equivalent conductance of the NaCl solution, in ohm-1cm2eq-1? Equivalent conductance of 0.1M-KCl is 120 ohm-1cm2eq-1.
(A) 12.5 (B) 250 (C) 125 (D) 25
3. A big irregular shaped vessel contained water, the specific conductance of which was 2.56 ×10–5 mho cm–1. 500 gm of NaCl was then added to the water and the specific conductance after the addition of NaCl, was found to be 3.10 × 10–5 mho cm–1. Find the capacity of the vessel if it is fulfilled with water. ΛºNaCl = 149.9.
4. The equivalent conductance of 0.10N solution of MgCl2 is 97.1 mho cm2 eq–1 at 25ºC. A cell with electrodes that are 1.50 cm2 in surface area and 0.50 cm apart is filled with 0.10N-MgCl
2 solution. How much current will flow when the potential difference between the electrodes is 5 volts?
5. For H+ and Na+, the values of symbol λº are 349.8 and 50.11. Calculate the mobilities of these ions
and their velocities if they are in a cell in which the electrodes are 5 cm apart and to which a potential of 2 volt is applied.
6. The equivalent conductance of an infinitely dilute solution of NH4Cl is 150 and the ionic conductance of OH– and Cl– ions are 198 and 76, respectively. What will be the equivalent conductance of the
solution of NH4OH at infinite dilution? If the equivalent conductance of a 0.01N–solution of NH4OH is 9.6, what will be its degree of dissociation?
7. Calculate the ionic product of water at 25ºC from the following data: specific conductance of water = 5.8 × 10–8 mho cm–1. λº H + = 350.0 and λºOH - = 198.0 mho cm2.
8. Calculate Ka of acetic acid if its 0.05N – solution has equivalent conductance of 7.36 mho cm2 eq–1 at
25ºC. Λº
3
CH COOH = 390.7.
9. A dilute solution of KCl was placed between two platinum electrodes 10cm apart, across which a potential of 6 volts was applied. How far would the K+ ion move in 2 hours at 25ºC? Ionic
conductance of K+ ion at infinite dilution at 25ºC is 73.52 mho cm2 mol-1 .
10. The specific conductance of a saturated solution of AgCl at 25ºC after substracting the specific conductance of water is 2.28 × 10–6 mho cm–1. Find the solubility product of AgCl at 25ºC. Λº
AgCl = 138.3 mho cm2.
C
HEMISTRY
A
RENA
(LEVEL #I)OBJECTIVE QUESTIONS HAVING ONLY ONE CORRECT OPTION
1. By how much would the oxidising power of MnO4¯ /Mn2+ couple change if the H+ ions concentration is decreased 100 times?
(A) increases by 189 mV (B) decreases by 189 mV (C) will increase by 19 mV (D) will decrease by 19 mV 2. The standard reduction potential for the reaction Ag+ + e → Ag and Ag(NH
3)2
+ + e → Ag + 2NH 3 are + 0.7991V and +0.373V, respectively. From these values and the Nearnst equation, what should be Kf for the Ag(NH3)2+ ion?
(A) 6.02 × 10-8 (B) 1.66 × 107 (C) 7.3 × 1019 (D) 1.37 × 10-20 3. The overall formation constant for the reaction of 6 mole of CN– with cobalt (II) is 1 × 1019. What is
the formation constant for the reaction of 6 moles of CN– with cobalt (III)? Given that Co(CN)6 + e → Co (CN)6 4-; Eº = – 0.83 V Co3+ + e → Co2+; Eº = + 1.82 V (A) 8.23 × 1063 (B) 8.23 × 1025 (C) 1.22 × 10 -26 (D) 1.22 × 10-64 4. Four colourless salt solutions are placed in separate test tubes and a strip of copper is placed in each.
Which solution finally turns blue?
(A) AgNO3 (B) Pb(NO3) 2 (C) Zn(NO3) 2 (D) Cd(NO3) 2 5. A student made the following observations in the laboratory:
(I) Clean copper metal did not react with 1M-Pb(NO3)2 solution
(II) Clean lead metal dissoles in 1M-AgNO3 solution and crystals of Ag metal appeared (III) Clean silver metal did not react with 1M-Cu(NO3)2 solution.
The order of decreasing reducing character of the three metals is
(A) Cu, Pb, Ag (B) Cu, Ag, Pb (C) Pb, Cu, Ag (D) Pb, Ag, Cu 6. Electrode potential will be more for hydrogen electrode with PH (at the same temperature)
(A) 4 (B) 3 (C) 2 (D) 5
7. If hydrochloric acid present in a hydrogen electrode coupled with a saturated calomel electrode is titrated against ten times stronger solution of NaOH, the variance in electrode potential is
(A) electrode potential changes slowly first and more and more rapidly as the end point approaches (B) electrode potential changes fast initially and very little change is observed as the end point
approaches
(C) electrode potential does not change at all during titration (D) none of the above
8. The solution of CuSO4 in which copper rod is immersed is diluted to 10 times, the reduction electrode potential
(A) Increases by 0.030 V (B) Decreases by 0.030 V (C) Increases by 0.059 V (D) Decreases by 0.059 V
9. K, Ca and Li metals may arranged in the decreasing order of their standard electrode potential as (A) K, Ca, Li (B) Li, K, Ca (C) Li, Ca, K (D) Ca, K, Li
10. A 0.1M solution of copper sulphate was electrolysed using copper electrodes, when the concentrations of copper ions were found to be 0.12M and 0.06M respectively. The back emf developed at 27ºC is
(A) 20 mV (B) 5.2 mV (C) 8.8 mV (D) 10 mV
11. The cell reaction for the given cell is spontaneous if Pt, Cl2 (P1 atm) | Cl
| Cl2 (P2 atm), Pt (A) P1 > P2 (B) P1 < P2 (C) P1 = P2 (D) P1 = 1 atm 12. What is the equilibrium constant of the reaction: 2Fe2+ + Au+⇌2 Fe+ + Au3+. Given Eº
Au+|Au = 1.68 V,
EºAu3+|Au = 1.50 V, EºFe3+|Fe2+= 0.77 V.
(A) 5 × 1021 (B) 2 × 10-22 (C) 3 × 10-18 (D) 1.5 × 10-11 13. What is the emf of the cell Pt, (1atm) | CH3COOH (0.1M) || (0.01M) NH4OH | H2 (1atm), Pt. Ka for
CH3COOH = 1.8 × 10−5 and K
b for NH4OH = 1.8 × 10− 5.
(A) 0.458 V (B) -0.458 V (C) -0.089 V (D) +0.89 V
14. The e.m.f. of a Daniel cell at 298K is E1, when the concentration of ZnSO4 is 1.0M and that of CuSO4 is 0.01M, the e.m.f. changed to E2. What is the relationship between E1 and E2?
(A) E1 = E2 (B) E2 = 0 (C) E1 > E2 (D) E1 < E2 15. When a lead storage battery is discharged
(A) SO2 is evolved (B) lead sulphate is consumed
(C) lead is formed (D) sulphuric acid is consumed
16. For a reaction A(s) + 2B+ → A2+ + B(s); Kc has been found to be 1012. The EMF of the cell is
(A) 0.354 V (B) 0.708 V (C) 0.534 V (D) 0.453 V
17. Which of the following metals when coupled will give maximum e.m.f. for a voltaic cell?
(A) Fe and Cu (B) Pb and Au (C) Cu and Au (D) Ca and Cu
18. F2 gas can not be obtained by the electrolysis of any aqueous fluoride salt because (A) F2 is the strongest oxidizing agent (B) F2 eaisly combines with water (C) F2 readily combines with the electrodes (D) All of these
19. Copper can be deposited from acidified copper sulphate and alkaline cuprous cyanide. If the same current is passed for a definite time
(A) the amount of copper deposited from acidic copper sulphate will be higher (B) the amount of copper deposited from alkaline cuprous cyanide will be higher (C) the same amount of copper will be deposited
(D) copper will not deposite in either case
20. The same current is passed through acidulated water and stannous chloride solution. What volume of dry detonating gas at NTP is evolved from water, when 1.20 gm of tin is deposited from the other solution? (Sn = 120)
(A) 448 ml (B) 336 ml (C) 224 ml (D) 672 ml
21. When molten ICI3 is electrolysed using platinum electrodes (A) I2 is evolved at cathode and Cl2 at anode
(B) Cl2 is evolved at cathode and I2 at anode
(C) I2 is evolved at cathode and both I2 and Cl2 at anode (D) electrolysis does not take place
22. When electric current is passed through a cell having an electrolyte, positive ions move towards the cathode and negative ions towards the anode. If the cathode is pulled out of the solution, then the (A) positive and negative ions will move towards the anode
(B) positive ions will start moving towards the anode, the negative ions will stop moving
(C) negative ions will continue to move towards the anode, the positive ions will start moving randomly
(D) positive and negative ions will start moving randomly 23. When 10-6M-HCl is electrolysed
(A) O2 is produced at the anode (B) H2 is produced at the anode
(C) Cl2 is produced at the anode (D) Cl2 and O2 are produecd at the anode 24. A constant current flowed for 2 hours through a potassuim iodide solution, oxidising the iodide ion to
iodine. At the end of the experiment, the iodine was titrated with 25 ml of 0.08M – Na2S2O3 solution. What was the average rate of current flow in ampere?
(A) 0.268 amp (B) 2.68 amp (C) 0.536 amp (D) 0.134 amp 25. Anthracene, C14H10, can be oxidised anodically to anthraquinone, C14H8O2. What weight of
anthraquinone can be produced by the passage of a current of 1 amp of 60 minutes if the current efficiency is 90 %?
(A) 6.98 gm (B) 1.16 gm (C) 41.90 gm (D) 0.99 gm
26. Most of the copper used to make wire has been electrically refined by depositing it from copper salts solution (divalent) on to a cathode. What is the cost of electrical energy required per kg of copper if the cost of electricity is Rs 4.00 per KWh and the cell operates at 0.33 volt? The electrochemical equivalent of copper is 0.00033 gm/coul.
(A) Rs. 11.11 (B) Rs. 5.55 (C) Rs. 2.22 (D) Rs. 1.11 27. Electrolysis of an acetate solution produces ethane according to the reaction:
2CH3COO– → C2H6(g) + 2CO2 (g) + 2e
What total volume of ethane and CO2 would be produced at STP if a current of 0.5 amp is passed through the solution for 482.5 min? Assume current efficiency 80%.
(A) 1.344 L (B) 2.016 L (C) 4.032 L (D) 1.792 L
28. To perform an analysis of a mixture of metal ions by electrodeposition, the second metal to be deposited must not being plating out until the concentration ratio of the second to the first is about 106. What must be the minimum difference in standard potential of the two metals which form
dipositive ions in order for such an analysis to be feasible?
(A) 0.177 V (B) 0.354 V (C) 0.708 V (D) 0.005 V
29. Three faradays of electricity are passed through molten Al2O3, aqueous solutions of CuSO4 and molten NaCl solutions. The amounts of Al, Cu and Na deposited at the cathodes will be in the ratio of moles indicated
(A) 1 : 2 : 3 (B) 1.5 : 2 : 3 (C) 1 : 1.5 : 3 (D) 1 : 3 : 2 30. Electrolysis of a solution of HSO4− ions produces S2O8
2−
. Assuming 75% current efficiency, what current should be employed to achieve a production rate of 1 mole of S2O82− per hour?
MULTIPLE CHOICE QUESTIONS MAY HAVE ONE OR MORE THAN ONE CORRECT OPTION.
31. Which one of the following statement is/are incorrect regarding an electrochemical cell? (A) the electrode on which oxidation takes place is called anode
(B) anode is the negative pole
(C) the direction of the current is same as that of the direction of flow of electrons (D) the flow of current is partly due to flow of electrons and partly due to flow of ions
32. Which of the following statement(s) differentiate between electrochemical cell and electrolytic cell? (A) spontaneous or non-spontaneous nature of the chemical process
(B) chemical reactions occuring at the electrodes (C) positive and negative nature of anode (D) dependence on Faraday’s law
33. Pick up the false statement(s):
(A) Galvanic cell reactions are always redox reactions
(B) In a galvanic cell made of cobalt and cadmium electrodes, cobalt electrode acts as anode (C) Standard potential increases with increasing concentration of the electrolyte
(D) Calomel electrode is a reference electrode having 0.00 volt potential
34. Consider the cell: Ag (s), AgCl(s) | KCl (0.1M) | Hg2Cl2 (s), Hg(l). The cell potential (A) Increases on increasing concentration of Cl– ions.
(B) Decreases on decreasing concentration of Cl– ions.
(C) is independent of concentration of Cl– ions
(D) is independent of amounts of AgCl and Hg2Cl2
COMPREHENSION BASED QUESTIONS HAVING ONLY ONE CORRECT OPTION
COMPREHENSION -I
The cell potential for the unbalanced chemical reaction: Hg2 2+ + NO3 + H3O+ →→→ Hg→ 2+ + HNO 2 + H2O
is measured under standard conditions in
the electrochemical cell shown in the accompanying diagram. 35. In which dish, the solution is acidic?
(A) Dish A (B) Dish B (C) Both (D) None
36. What is the equilibrium constant for the reaction?
(A) 1.97 (B) 4.76 (C) 2.18 (D) 1.40
37. How many moles of electrons pass through the circuit when 0.60 mole of Hg2+ and 0.30 mole of HNO2 are produced in the cell that contains 0.50 mole of Hg2
2+
and 0.40 mole of NO3
at the beginning of the reaction?
(A) 0.30 (B) 0.60 (C) 0.15 (D) 1.20
38. How long will it take to produce 0.10 mole of HNO2 by this reaction if a current of 10 A passes through the cell?
COMPREHENSION -II
A fuel cell is the device to convert the energy of a fuel into electrical energy without the use of heat engine, where the fuel is burnt directly. Such conversions are possible because the combustion reactions are essentially redox reactions and highly exothermic as well as highly exergonic. Electrical energy can be obtained indefinitely from a fuel cell as along as the outside supply of fuel is maintained. In hydrogen – oxygen fuel cell, the following reaction take place: Anode reaction: 2H2 (g)+ 4OH
–
(aq) →→→→ 4H2O(l) + 4e –
Cathode reaction: O2(g) + 2H2O(l) + 4e –→→→→
4OH–(aq) Overall reaction: 2H2(g) + O2(g) →→→ 2H→ 2O(l)
The overall all reaction has a value of ∆∆∆∆H°°°° = –285.8 kJ mol–1 and ∆∆∆∆G°°°° = –237.39 kJ mol–1 at 25°°°°C.
39. What is the standard EMF of the cell?
(A) 0.615 V (B) 1.23 V (C) 2.46 V (D) 0.74 V
40. Which of the following expressions gives ∆G° for the reaction in the fuel cell?
(A) 4 × 96500 × 1.23 J (B) – 4 × 96500 × 1.23 J
(C) –8.314 × 298 ln 1.23 J (D) 8.314 × 298 × ln 286 J
41. How many litres of gaseous H2, when combined with excess O2 in the fuel cell at 25°C and 1.00 atm, are needed to produce 23.739 KJ of useful work under ideal conditions?
(A) 2.44L (B) 4.88L (C) 1.22L (D) 2.24L
42. Suppose the concentration of hydroxide ion in the cell is doubled. The cell voltage will be
(A) reduced by half (B) increased by a factor of 2
(C) increased by a factor of 4 (D) unchanged 43. What is the approximate value of ∆S° for the fuel cell reaction at 25°C?
(A) –0.164 JK–1 (B) –164 JK–1 (C) 164 JK–1 (D) 0.164 JK–1 44. The theoretical efficiency of the fuel cell is given by
(A) 83.06 % (B) 100 %
(C) 67.53 % (D) 97.88 %
COMPREHENSION -III
Suppose that the S.H.E. was arbitrarily assigned a value of 1.00 volt for 2H+ + 2e →→→→ H
2(g). What
would this do to the observed voltage under standard condition for each of the following: (EºZn 2+ |Zn = – 0.76 volt, EºCu 2+ |Cu = + 0.34 volt) 45. Zn – H2 cell (A) +0.34 V (B) -0.76 V (C) +1.76 V (D) -0.34 V 46. Cu – H2 Cell (A) -0.34 V (B) +1.76 V (C) -1.76 V (D) +0.34 V 47. Zn – Cu Cell (A) +1.10 V (B) -1.10 V (C) 0.0 V (D) Indeterminate
COMPREHENSION -IV
A current of 15.0 A is employed to plate nickel in a NiSO4 solution. Both Ni and H2 are formed
at the cathode. The current efficiency with respect to formation of Ni is 60%. The density of nickel = 8.9 gm/ml. (Ni = 58.7)
48. How much of nickel is plated on the cathode per hour?
(A) 16.43 gm (B) 32.85 gm (C) 19.7 gm (D) 9.85 gm
49. What is the thickness of the plating if the cathode consists of a sheet of metal 4.0 cm2 which is to be coated on both sides?
(A) 13.8 mm (B) 27.6 mm (C) 6.9 mm (D) 23.0 mm
50. What volume of H2 at STP is formed per hour?
