330222
Assertion : In electrolysis, the quantity of electricity needed for depositing 1 mole of silver is different from that required for 1 mole of copper. Reason : The atomic weight of silver and copper are different.
1 Both Assertion and Reason are correct and Reason is the correct explanation of the Assertion.
2 Both Assertion and Reason are correct but Reason is not the correct explanation of the Assertion.
3 Assertion is correct but Reason is incorrect.
4 Assertion is incorrect but Reason is correct.
Explanation:
As per Faraday's law of electrolysis, the mass of the substances deposited are proportional to their respective equivalent weights. Equivalent weight of \(\mathrm{Ag}=108\) and equivalent weight of \(\mathrm{Cu}=63.5 / 2=31.75\). Atomic weights of \(\mathrm{Ag}\) and \(\mathrm{Cu}\) are different, but the correct reason for the assertion is the difference in their equivalent weights, not molecular weights. So, option (2) is correct.
CHXII03:ELECTROCHEMISTRY
330223
The quantity of electricity needed to separately electrolyse \(1 \mathrm{M}\) solution of \(\mathrm{ZnSO}_{4}, \mathrm{AlCl}_{3}\), and \(\mathrm{AgNO}_{3}\) completely is in the ratio of:
1 \(2: 3: 1\)
2 \(2: 1: 1\)
3 \(2: 1: 3\)
4 \(2: 2: 1\)
Explanation:
1 mole \(\mathrm{Zn}^{+2}\) needs \(2 \mathrm{~F}, 1\) mole \(\mathrm{Al}^{+3}\) needs \(3 \mathrm{~F}\) and 1 mole \(\mathrm{Ag}^{+}\)needs \(1 \mathrm{~F}\) Thus; The ratio of the amounts of electricity required is \(2: 3: 1\).
CHXII03:ELECTROCHEMISTRY
330224
Two faraday of electricity is passed through a solution of \({\rm{CuS}}{{\rm{O}}_{\rm{4}}}\). The mass of copper deposited at the cathode is (at. mass of Cu = 63.5 amu)
330225
What is the time required (in seconds) for depositing all the silver present in \(125 \mathrm{~mL}\) of \({\text{1M AgN}}{{\text{O}}_{\text{3}}}\) solution by passing a current of 241.25 A \([1 {\text{F}} = 96500\,\,{\text{C}}]?\)
1 \({\text{10}}\,\,{\text{sec}}\)
2 \({\text{50}}\,\,{\text{sec}}\)
3 \({\text{1000}}\,\,{\text{sec}}\)
4 \({\text{100}}\,\,{\text{sec}}\)
Explanation:
\(1000 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution contains \(108 \mathrm{~g}\) \(\mathrm{Ag}\) \(\therefore 125 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution will contain \(=\dfrac{108 \times 125}{1000}=13.5 \mathrm{~g} \mathrm{Ag}\) \(\mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}\) \((1 \mathrm{~F}=96500) 108 \mathrm{~g}\) \(108 \mathrm{~g}\) of \(\mathrm{Ag}\) is deposited by \(96500 \mathrm{C}\) \(\therefore 13.5\) of \(\mathrm{Ag}\) is deposited by \(=\dfrac{96500}{108} \times 13.5=\) 12062.5 C \(\mathrm{Q}=\mathrm{It}\) \(\mathrm{t}=\dfrac{\mathrm{Q}}{\mathrm{I}}=\dfrac{12062.5}{241.25}=50 \mathrm{sec}\)
CHXII03:ELECTROCHEMISTRY
330226
How many coulombs of electricity are required for the oxidation of one mole of water to dioxygen?
\({{\rm{H}}_{\rm{2}}}{\rm{O}} \to {\rm{2}}{{\rm{H}}^{\rm{ + }}}{\rm{ + }}\frac{{\rm{1}}}{{\rm{2}}}{{\rm{O}}_{\rm{2}}}{\rm{ + 2}}{{\rm{e}}^{\rm{ - }}}\) 2F of electricity required. Hence coulombs of electricity required for the oxidation of one mole of water to dioxygen \({\rm{ = 2 \times 96500 = 193000}}\,\,{\rm{C = 1}}{\rm{.93 \times 1}}{{\rm{0}}^{\rm{5}}}\,{\rm{C}}\).
330222
Assertion : In electrolysis, the quantity of electricity needed for depositing 1 mole of silver is different from that required for 1 mole of copper. Reason : The atomic weight of silver and copper are different.
1 Both Assertion and Reason are correct and Reason is the correct explanation of the Assertion.
2 Both Assertion and Reason are correct but Reason is not the correct explanation of the Assertion.
3 Assertion is correct but Reason is incorrect.
4 Assertion is incorrect but Reason is correct.
Explanation:
As per Faraday's law of electrolysis, the mass of the substances deposited are proportional to their respective equivalent weights. Equivalent weight of \(\mathrm{Ag}=108\) and equivalent weight of \(\mathrm{Cu}=63.5 / 2=31.75\). Atomic weights of \(\mathrm{Ag}\) and \(\mathrm{Cu}\) are different, but the correct reason for the assertion is the difference in their equivalent weights, not molecular weights. So, option (2) is correct.
CHXII03:ELECTROCHEMISTRY
330223
The quantity of electricity needed to separately electrolyse \(1 \mathrm{M}\) solution of \(\mathrm{ZnSO}_{4}, \mathrm{AlCl}_{3}\), and \(\mathrm{AgNO}_{3}\) completely is in the ratio of:
1 \(2: 3: 1\)
2 \(2: 1: 1\)
3 \(2: 1: 3\)
4 \(2: 2: 1\)
Explanation:
1 mole \(\mathrm{Zn}^{+2}\) needs \(2 \mathrm{~F}, 1\) mole \(\mathrm{Al}^{+3}\) needs \(3 \mathrm{~F}\) and 1 mole \(\mathrm{Ag}^{+}\)needs \(1 \mathrm{~F}\) Thus; The ratio of the amounts of electricity required is \(2: 3: 1\).
CHXII03:ELECTROCHEMISTRY
330224
Two faraday of electricity is passed through a solution of \({\rm{CuS}}{{\rm{O}}_{\rm{4}}}\). The mass of copper deposited at the cathode is (at. mass of Cu = 63.5 amu)
330225
What is the time required (in seconds) for depositing all the silver present in \(125 \mathrm{~mL}\) of \({\text{1M AgN}}{{\text{O}}_{\text{3}}}\) solution by passing a current of 241.25 A \([1 {\text{F}} = 96500\,\,{\text{C}}]?\)
1 \({\text{10}}\,\,{\text{sec}}\)
2 \({\text{50}}\,\,{\text{sec}}\)
3 \({\text{1000}}\,\,{\text{sec}}\)
4 \({\text{100}}\,\,{\text{sec}}\)
Explanation:
\(1000 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution contains \(108 \mathrm{~g}\) \(\mathrm{Ag}\) \(\therefore 125 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution will contain \(=\dfrac{108 \times 125}{1000}=13.5 \mathrm{~g} \mathrm{Ag}\) \(\mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}\) \((1 \mathrm{~F}=96500) 108 \mathrm{~g}\) \(108 \mathrm{~g}\) of \(\mathrm{Ag}\) is deposited by \(96500 \mathrm{C}\) \(\therefore 13.5\) of \(\mathrm{Ag}\) is deposited by \(=\dfrac{96500}{108} \times 13.5=\) 12062.5 C \(\mathrm{Q}=\mathrm{It}\) \(\mathrm{t}=\dfrac{\mathrm{Q}}{\mathrm{I}}=\dfrac{12062.5}{241.25}=50 \mathrm{sec}\)
CHXII03:ELECTROCHEMISTRY
330226
How many coulombs of electricity are required for the oxidation of one mole of water to dioxygen?
\({{\rm{H}}_{\rm{2}}}{\rm{O}} \to {\rm{2}}{{\rm{H}}^{\rm{ + }}}{\rm{ + }}\frac{{\rm{1}}}{{\rm{2}}}{{\rm{O}}_{\rm{2}}}{\rm{ + 2}}{{\rm{e}}^{\rm{ - }}}\) 2F of electricity required. Hence coulombs of electricity required for the oxidation of one mole of water to dioxygen \({\rm{ = 2 \times 96500 = 193000}}\,\,{\rm{C = 1}}{\rm{.93 \times 1}}{{\rm{0}}^{\rm{5}}}\,{\rm{C}}\).
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CHXII03:ELECTROCHEMISTRY
330222
Assertion : In electrolysis, the quantity of electricity needed for depositing 1 mole of silver is different from that required for 1 mole of copper. Reason : The atomic weight of silver and copper are different.
1 Both Assertion and Reason are correct and Reason is the correct explanation of the Assertion.
2 Both Assertion and Reason are correct but Reason is not the correct explanation of the Assertion.
3 Assertion is correct but Reason is incorrect.
4 Assertion is incorrect but Reason is correct.
Explanation:
As per Faraday's law of electrolysis, the mass of the substances deposited are proportional to their respective equivalent weights. Equivalent weight of \(\mathrm{Ag}=108\) and equivalent weight of \(\mathrm{Cu}=63.5 / 2=31.75\). Atomic weights of \(\mathrm{Ag}\) and \(\mathrm{Cu}\) are different, but the correct reason for the assertion is the difference in their equivalent weights, not molecular weights. So, option (2) is correct.
CHXII03:ELECTROCHEMISTRY
330223
The quantity of electricity needed to separately electrolyse \(1 \mathrm{M}\) solution of \(\mathrm{ZnSO}_{4}, \mathrm{AlCl}_{3}\), and \(\mathrm{AgNO}_{3}\) completely is in the ratio of:
1 \(2: 3: 1\)
2 \(2: 1: 1\)
3 \(2: 1: 3\)
4 \(2: 2: 1\)
Explanation:
1 mole \(\mathrm{Zn}^{+2}\) needs \(2 \mathrm{~F}, 1\) mole \(\mathrm{Al}^{+3}\) needs \(3 \mathrm{~F}\) and 1 mole \(\mathrm{Ag}^{+}\)needs \(1 \mathrm{~F}\) Thus; The ratio of the amounts of electricity required is \(2: 3: 1\).
CHXII03:ELECTROCHEMISTRY
330224
Two faraday of electricity is passed through a solution of \({\rm{CuS}}{{\rm{O}}_{\rm{4}}}\). The mass of copper deposited at the cathode is (at. mass of Cu = 63.5 amu)
330225
What is the time required (in seconds) for depositing all the silver present in \(125 \mathrm{~mL}\) of \({\text{1M AgN}}{{\text{O}}_{\text{3}}}\) solution by passing a current of 241.25 A \([1 {\text{F}} = 96500\,\,{\text{C}}]?\)
1 \({\text{10}}\,\,{\text{sec}}\)
2 \({\text{50}}\,\,{\text{sec}}\)
3 \({\text{1000}}\,\,{\text{sec}}\)
4 \({\text{100}}\,\,{\text{sec}}\)
Explanation:
\(1000 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution contains \(108 \mathrm{~g}\) \(\mathrm{Ag}\) \(\therefore 125 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution will contain \(=\dfrac{108 \times 125}{1000}=13.5 \mathrm{~g} \mathrm{Ag}\) \(\mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}\) \((1 \mathrm{~F}=96500) 108 \mathrm{~g}\) \(108 \mathrm{~g}\) of \(\mathrm{Ag}\) is deposited by \(96500 \mathrm{C}\) \(\therefore 13.5\) of \(\mathrm{Ag}\) is deposited by \(=\dfrac{96500}{108} \times 13.5=\) 12062.5 C \(\mathrm{Q}=\mathrm{It}\) \(\mathrm{t}=\dfrac{\mathrm{Q}}{\mathrm{I}}=\dfrac{12062.5}{241.25}=50 \mathrm{sec}\)
CHXII03:ELECTROCHEMISTRY
330226
How many coulombs of electricity are required for the oxidation of one mole of water to dioxygen?
\({{\rm{H}}_{\rm{2}}}{\rm{O}} \to {\rm{2}}{{\rm{H}}^{\rm{ + }}}{\rm{ + }}\frac{{\rm{1}}}{{\rm{2}}}{{\rm{O}}_{\rm{2}}}{\rm{ + 2}}{{\rm{e}}^{\rm{ - }}}\) 2F of electricity required. Hence coulombs of electricity required for the oxidation of one mole of water to dioxygen \({\rm{ = 2 \times 96500 = 193000}}\,\,{\rm{C = 1}}{\rm{.93 \times 1}}{{\rm{0}}^{\rm{5}}}\,{\rm{C}}\).
330222
Assertion : In electrolysis, the quantity of electricity needed for depositing 1 mole of silver is different from that required for 1 mole of copper. Reason : The atomic weight of silver and copper are different.
1 Both Assertion and Reason are correct and Reason is the correct explanation of the Assertion.
2 Both Assertion and Reason are correct but Reason is not the correct explanation of the Assertion.
3 Assertion is correct but Reason is incorrect.
4 Assertion is incorrect but Reason is correct.
Explanation:
As per Faraday's law of electrolysis, the mass of the substances deposited are proportional to their respective equivalent weights. Equivalent weight of \(\mathrm{Ag}=108\) and equivalent weight of \(\mathrm{Cu}=63.5 / 2=31.75\). Atomic weights of \(\mathrm{Ag}\) and \(\mathrm{Cu}\) are different, but the correct reason for the assertion is the difference in their equivalent weights, not molecular weights. So, option (2) is correct.
CHXII03:ELECTROCHEMISTRY
330223
The quantity of electricity needed to separately electrolyse \(1 \mathrm{M}\) solution of \(\mathrm{ZnSO}_{4}, \mathrm{AlCl}_{3}\), and \(\mathrm{AgNO}_{3}\) completely is in the ratio of:
1 \(2: 3: 1\)
2 \(2: 1: 1\)
3 \(2: 1: 3\)
4 \(2: 2: 1\)
Explanation:
1 mole \(\mathrm{Zn}^{+2}\) needs \(2 \mathrm{~F}, 1\) mole \(\mathrm{Al}^{+3}\) needs \(3 \mathrm{~F}\) and 1 mole \(\mathrm{Ag}^{+}\)needs \(1 \mathrm{~F}\) Thus; The ratio of the amounts of electricity required is \(2: 3: 1\).
CHXII03:ELECTROCHEMISTRY
330224
Two faraday of electricity is passed through a solution of \({\rm{CuS}}{{\rm{O}}_{\rm{4}}}\). The mass of copper deposited at the cathode is (at. mass of Cu = 63.5 amu)
330225
What is the time required (in seconds) for depositing all the silver present in \(125 \mathrm{~mL}\) of \({\text{1M AgN}}{{\text{O}}_{\text{3}}}\) solution by passing a current of 241.25 A \([1 {\text{F}} = 96500\,\,{\text{C}}]?\)
1 \({\text{10}}\,\,{\text{sec}}\)
2 \({\text{50}}\,\,{\text{sec}}\)
3 \({\text{1000}}\,\,{\text{sec}}\)
4 \({\text{100}}\,\,{\text{sec}}\)
Explanation:
\(1000 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution contains \(108 \mathrm{~g}\) \(\mathrm{Ag}\) \(\therefore 125 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution will contain \(=\dfrac{108 \times 125}{1000}=13.5 \mathrm{~g} \mathrm{Ag}\) \(\mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}\) \((1 \mathrm{~F}=96500) 108 \mathrm{~g}\) \(108 \mathrm{~g}\) of \(\mathrm{Ag}\) is deposited by \(96500 \mathrm{C}\) \(\therefore 13.5\) of \(\mathrm{Ag}\) is deposited by \(=\dfrac{96500}{108} \times 13.5=\) 12062.5 C \(\mathrm{Q}=\mathrm{It}\) \(\mathrm{t}=\dfrac{\mathrm{Q}}{\mathrm{I}}=\dfrac{12062.5}{241.25}=50 \mathrm{sec}\)
CHXII03:ELECTROCHEMISTRY
330226
How many coulombs of electricity are required for the oxidation of one mole of water to dioxygen?
\({{\rm{H}}_{\rm{2}}}{\rm{O}} \to {\rm{2}}{{\rm{H}}^{\rm{ + }}}{\rm{ + }}\frac{{\rm{1}}}{{\rm{2}}}{{\rm{O}}_{\rm{2}}}{\rm{ + 2}}{{\rm{e}}^{\rm{ - }}}\) 2F of electricity required. Hence coulombs of electricity required for the oxidation of one mole of water to dioxygen \({\rm{ = 2 \times 96500 = 193000}}\,\,{\rm{C = 1}}{\rm{.93 \times 1}}{{\rm{0}}^{\rm{5}}}\,{\rm{C}}\).
330222
Assertion : In electrolysis, the quantity of electricity needed for depositing 1 mole of silver is different from that required for 1 mole of copper. Reason : The atomic weight of silver and copper are different.
1 Both Assertion and Reason are correct and Reason is the correct explanation of the Assertion.
2 Both Assertion and Reason are correct but Reason is not the correct explanation of the Assertion.
3 Assertion is correct but Reason is incorrect.
4 Assertion is incorrect but Reason is correct.
Explanation:
As per Faraday's law of electrolysis, the mass of the substances deposited are proportional to their respective equivalent weights. Equivalent weight of \(\mathrm{Ag}=108\) and equivalent weight of \(\mathrm{Cu}=63.5 / 2=31.75\). Atomic weights of \(\mathrm{Ag}\) and \(\mathrm{Cu}\) are different, but the correct reason for the assertion is the difference in their equivalent weights, not molecular weights. So, option (2) is correct.
CHXII03:ELECTROCHEMISTRY
330223
The quantity of electricity needed to separately electrolyse \(1 \mathrm{M}\) solution of \(\mathrm{ZnSO}_{4}, \mathrm{AlCl}_{3}\), and \(\mathrm{AgNO}_{3}\) completely is in the ratio of:
1 \(2: 3: 1\)
2 \(2: 1: 1\)
3 \(2: 1: 3\)
4 \(2: 2: 1\)
Explanation:
1 mole \(\mathrm{Zn}^{+2}\) needs \(2 \mathrm{~F}, 1\) mole \(\mathrm{Al}^{+3}\) needs \(3 \mathrm{~F}\) and 1 mole \(\mathrm{Ag}^{+}\)needs \(1 \mathrm{~F}\) Thus; The ratio of the amounts of electricity required is \(2: 3: 1\).
CHXII03:ELECTROCHEMISTRY
330224
Two faraday of electricity is passed through a solution of \({\rm{CuS}}{{\rm{O}}_{\rm{4}}}\). The mass of copper deposited at the cathode is (at. mass of Cu = 63.5 amu)
330225
What is the time required (in seconds) for depositing all the silver present in \(125 \mathrm{~mL}\) of \({\text{1M AgN}}{{\text{O}}_{\text{3}}}\) solution by passing a current of 241.25 A \([1 {\text{F}} = 96500\,\,{\text{C}}]?\)
1 \({\text{10}}\,\,{\text{sec}}\)
2 \({\text{50}}\,\,{\text{sec}}\)
3 \({\text{1000}}\,\,{\text{sec}}\)
4 \({\text{100}}\,\,{\text{sec}}\)
Explanation:
\(1000 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution contains \(108 \mathrm{~g}\) \(\mathrm{Ag}\) \(\therefore 125 \mathrm{~mL}\) of \(\mathrm{AgNO}_{3}\) solution will contain \(=\dfrac{108 \times 125}{1000}=13.5 \mathrm{~g} \mathrm{Ag}\) \(\mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}\) \((1 \mathrm{~F}=96500) 108 \mathrm{~g}\) \(108 \mathrm{~g}\) of \(\mathrm{Ag}\) is deposited by \(96500 \mathrm{C}\) \(\therefore 13.5\) of \(\mathrm{Ag}\) is deposited by \(=\dfrac{96500}{108} \times 13.5=\) 12062.5 C \(\mathrm{Q}=\mathrm{It}\) \(\mathrm{t}=\dfrac{\mathrm{Q}}{\mathrm{I}}=\dfrac{12062.5}{241.25}=50 \mathrm{sec}\)
CHXII03:ELECTROCHEMISTRY
330226
How many coulombs of electricity are required for the oxidation of one mole of water to dioxygen?
\({{\rm{H}}_{\rm{2}}}{\rm{O}} \to {\rm{2}}{{\rm{H}}^{\rm{ + }}}{\rm{ + }}\frac{{\rm{1}}}{{\rm{2}}}{{\rm{O}}_{\rm{2}}}{\rm{ + 2}}{{\rm{e}}^{\rm{ - }}}\) 2F of electricity required. Hence coulombs of electricity required for the oxidation of one mole of water to dioxygen \({\rm{ = 2 \times 96500 = 193000}}\,\,{\rm{C = 1}}{\rm{.93 \times 1}}{{\rm{0}}^{\rm{5}}}\,{\rm{C}}\).