Question 1

In a common car battery, six identical cells each carry out the following reaction: Pb + PbO₂ + 2HSO₄^- + 2H⁺ → 2PbSO₄ + 2H₂O Suppose that to start a car on a cold morning, 138 amperes is drawn for 12.0 seconds from such a cell. How many grams of Pb are consumed? (The atomic mass of Pb is 207.19 g/mol.)

Accepted Answer

First, we need to find the amount of charge that passes through the cell using the equation Q = It, where Q is charge (coulombs), I is current (amperes), and t is time (seconds). 
Q = (138 A)(12.0 s) = 1656 C 
Next, we use the balanced equation to find the moles of Pb consumed per cell: 
1 mol Pb : 2 mol e- (from the half-reaction) 
2 mol e- : 2HSO4- 
1 mol Pb : 2HSO4- 
1 mol Pb : 2(98.08 g/mol) = 196.16 g/mol 
2 mol HSO4- : 1 mol Pb 
1656 C × 1 mol e-/96500 C (Faraday constant) × 1 mol HSO4- / 2 mol e- × 1 mol Pb / 2 mol HSO4- × 196.16 g/mol Pb = 1.78 g Pb

Question 2

Calculate E at 25°C for this cell, given the following data: Ag⁺ + e^- → Ag(s) E° = 0.80 V Ni²⁺ + 2e^- → Ni(s) E° = -0.23 V K_sp for AgCl = 1.6 × 10^-10

Accepted Answer

The answer of Calculate E at 25°C for this cell,...

Question 3

Fe²⁺ + 2e^- → Fe(s) E° = -0.440 V 2H⁺ + 2e^- → H₂(g) E° = 0.000 V In a galvanic cell, the iron compartment contains an iron electrode, and [Fe²⁺] = 1.00 × 10^-3 M. The hydrogen compartment contains a platinum electrode (P_H2 = 1.00 atm) and a weak acid HA at an initial concentration of 1.00 M. If the observed cell potential is 0.333 V at 25°C, calculate K_a for the weak acid HA at 25°C.

Accepted Answer

The answer of Fe²⁺ + 2e^- → Fe(s) E° =...

Question 4

An antique automobile bumper is to be chrome plated. The bumper, which is dipped into an acidic CrS1U1B12S1U1B0OS1U1B17S1U1B0S1U1P12-S1S1P0 solution, serves as a cathode of an electrolytic cell. The atomic mass of Cr is 51.996; 1 faraday = 96,485 coulombs. -If oxidation of H₂O occurs at the anode, how many moles of oxygen gas will evolve for every 3.30 × 10² g of Cr(s) deposited?

Accepted Answer

The answer of An antique automobile bumper is to be...

Question 5

A solution of MnO₄^2- is electrolytically reduced to Mn³⁺. A current of 8.46 amp is passed through the solution for 15.0 minutes. What is the number of moles of Mn³⁺ produced in this process? (1 faraday = 96,485 coulombs)

Accepted Answer

The charge passed is calculated as current (8.46 A) multiplied by time (15.0 min converted to seconds, i.e., 900 s), which equals 7614 C. The reduction of MnO₄²⁻ to Mn³⁺ involves a 3-electron change. Therefore, the moles of Mn³⁺ produced is 7614 C divided by (3 × 96485 C/mol), which equals approximately 0.0263 mol.

Question 6

If an electrolysis plant operates its electrolytic cells at a total current of 1.0 × 10⁶ amp, how long will it take to produce one metric ton (one million grams) of Mg(s) from seawater containing Mg²⁺? (1 faraday = 96,485 coulombs)

Accepted Answer

To calculate the time needed to produce one metric ton of Mg(s), we need to use the stoichiometry of the reaction and Faraday's law.

The balanced equation for the electrolysis of seawater containing MgS1U1P12+S1S1P0 is:

2MgS1U1P12+(aq) + 2H2O(l) → 2Mg(s) + O2(g) + 2H2(g)

According to the equation, for every 2 moles of Mg produced, the amount of charge needed is given by:

2 mol Mg × (2 F/mol Mg) × (96,485 C/F) = 3.88 × 10S5U1P1S1S1P20 C

Therefore, to produce one ton (1000 kg) of Mg, we need:

1000 kg Mg × (1 mol Mg/24.31 g Mg) × (3.88 × 10S5U1P1S1S1P20 C/2 mol Mg) = 7.98 × 10S7U1P1S1S1P20 C

Now, we can use Faraday's law to find the time needed:

Q = I × t

where Q is the charge in coulombs, I is the current in amperes, and t is the time in seconds.

Solving for t:

t = Q/I = (7.98 × 10S7U1P1S1S1P20 C) / (1.0 × 10S1U1P16S1S1P0 A) = 2.2 × 10S3U1P4S1S1P0 s

Converting to hours:

t = 2.2 × 10S3U1P4S1S1P0 s × (1 h/3600 s) ≈ 0.61 h ≈ 36.6 min

Therefore, the best answer choice is C, 2.2 h.

Question 7

How many seconds would it take to deposit 21.40 g of Ag (atomic mass = 107.87) from a solution of AgNO₃ using a current of 10.00 amp?

Accepted Answer

The answer of How many seconds would it take to...

Question 8

If a constant current of 5.2 amperes is passed through a cell containing Cr³⁺ for 2.1 hour, how many grams of Cr will plate out onto the cathode? (The atomic mass of Cr is 51.996 g/mol.)

Accepted Answer

The answer of If a constant current of 5.2 amperes...

Question 9

What quantity of charge is required to reduce 32.6 g of CrCl₃ to chromium metal? (1 faraday = 96,485 coulombs)

Accepted Answer

The answer of What quantity of charge is required to...

Question 10

Use the following data to calculate the K_sp value at 25°C for PbSO₄(s).

Accepted Answer

The answer of Use the following data to calculate the...

Question 11

An electrolytic cell process involves plating Zr(s) from a solution containing Zr⁴⁺. If 5.80 amp is run through this mixture for 1.86 h, what mass of Zr is plated?

Accepted Answer

To calculate the mass of Zr plated, use the formula: mass = (It/nF) * M, where I = current (5.80 A), t = time (1.86 h = 6696 s), n = number of electrons (4 for Zr^4+), F = Faraday's constant (96485 C/mol), and M = molar mass of Zr (91.22 g/mol). The calculation gives approximately 9.18 g.

Question 12

Electrolysis of a molten salt with the formula MCl, using a current of 3.86 amp for 16.2 min, deposits 1.52 g of metal. Identify the metal. (1 faraday = 96,485 coulombs)

Accepted Answer

The answer of Electrolysis of a molten salt with the...

Question 13

Why is aluminum protected from corrosion? (Note: The standard reduction potential for Al³⁺ is -1.66 V.)

Accepted Answer

Aluminum forms a protective oxide coating on its surface, which prevents it from further corrosion. This oxide coating is formed due to the reaction of aluminum with oxygen in the air. The oxide layer is very thin and adherent, preventing any further reaction between aluminum and its surroundings. Therefore, option C is the correct choice. Option B is partially correct, but the standard reduction potential alone does not explain the formation of the oxide layer. Option A is incorrect since aluminum does react with oxygen to form the protective oxide layer. Option D is incorrect since only option C is correct. Option E is also incorrect since aluminum is indeed protected from corrosion.

Question 14

AgS1U1P1+S1S1P0 + eS1U1P1-S1S1P0 ? Ag(s) E° = 0.80 V
CuS1U1P12+S1S1P0 + 2eS1U1P1-S1S1P0 ? Cu(s) E° = 0.34 V
In a galvanic cell, the silver compartment contains a silver electrode and excess AgCl(s) (KS1U1B1spS1U1B0 = 1.6 × 10S1U1P1-10S1S1P0). The copper compartment contains a copper electrode, and [CuS1U1P12+S1S1P0] = 2.0 M.

A. Calculate the potential for this cell at 25°C.
B. Assuming 1.0 L of 2.0 M CuS1U1P12+S1S1P0 in the copper compartment, calculate how many moles of NHS1U1B13S1U1B0 would have to be added to establish the cell potential at 0.52 V at 25°C (assume no volume change on addition of NHS1U1B13S1U1B0).
CuS1U1P12+S1S1P0 + 4NHS1U1B13S1U1B0 Cu(NHS1U1B13S1U1B0)S1U1B14S1U1B0S1U112+S1S1P0 KS1U1B1fS1U1B0 = 1.0 ×10S1U1P113S1S1P0

Accepted Answer

The answer of AgS1U1P1+S1S1P0 + eS1U1P1-S1S1P0 ? Ag(s) E° =...

Question 15

Copper is electroplated from an aqueous CuSO₄ solution. A constant current of 4.70 amp is applied by an external power supply. How long will it take to deposit 3.76 × 10² g of Cu? The atomic mass of copper is 63.546 g/mol.

Accepted Answer

The answer of Copper is electroplated from an aqueous CuSO₄...

Question 16

Calculate the solubility product of silver iodide at 25°C, given the following data:

Accepted Answer

The answer of Calculate the solubility product of silver iodide...

Question 17

An antique automobile bumper is to be chrome plated. The bumper, which is dipped into an acidic CrS1U1B12S1U1B0OS1U1B17S1U1B0S1U1P12-S1S1P0 solution, serves as a cathode of an electrolytic cell. The atomic mass of Cr is 51.996; 1 faraday = 96,485 coulombs. -If the current is 18.8 amperes, how long will it take to deposit 1.66 × 10² g of Cr(s) onto the bumper?

Accepted Answer

To calculate the time required to deposit 1.66 × 10^12 g of Cr onto the bumper, we use Faraday's laws of electrolysis. First, we need to find the number of moles of Cr to be deposited:$$ \text{Moles of Cr} = \frac{\text{Mass}}{\text{Molar Mass}} = \frac{1.66 \times 10^{12} \text{ g}}{51.996 \text{ g/mol}} $$Next, we calculate the total charge needed to deposit this amount of Cr, knowing that 1 mole of Cr requires 1 Faraday (96,485 C) to be deposited:$$ \text{Total Charge} = \text{Moles of Cr} \times 96,485 \text{ C/mol} $$Finally, we find the time required using the current (I) and the total charge (Q) with the formula $ t = \frac{Q}{I} $, where $ t $ is in seconds. To convert seconds to hours, we divide by 3600.Given the current is 18.8 amperes, and using the steps above, the calculation will show that the time required is approximately 27.3 hours.

Question 18

Gold (atomic mass = 197 g/mol) is plated from a solution of chlorauric acid, HAuCl₄; it deposits on the cathode. Calculate the time it takes to deposit 0.55 g of gold, passing a current of 0.12 amperes. (1 faraday = 96,485 coulombs)

Accepted Answer

The first step is to calculate the number of moles of gold that will be deposited:

Atomic mass of gold = 197 g/mol
Moles of gold = Mass of gold / Atomic mass of gold
Moles of gold = 0.55 g / 197 g/mol
Moles of gold = 0.00279 mol

Next, we can use Faraday's Law of Electrolysis to calculate the amount of time (t) required to deposit this amount of gold:

Q = nF
where
Q = total charge passed in coulombs
n = number of moles of electrons transferred
F = Faraday's constant = 96,485 coulombs/mol

In this case, each gold ion (Au3+) receives 3 electrons to become neutral gold atoms, so:

n = 0.00279 mol x 3 = 0.00837 mol e-
Q = 0.00837 mol x 96,485 coulombs/mol = 808 coulombs

Finally, we can use the definition of current (I = Q/t) to solve for time:

t = Q / I
t = 808 coulombs / 0.12 A
t = 6733.3 seconds

Converting to hours:

t = 6733.3 seconds x (1 min / 60 seconds) x (1 hour / 60 min)
t = 1.87 hours, or about 1.9 hours

Therefore, the best choice is C.

Question 19

Nickel is electroplated from a NiSO₄ solution. A constant current of 5.72 amp is applied by an external power supply. How long will it take to deposit 2.08 × 10² g of Ni? (The atomic mass of Ni is 58.69 g/mol.)

Accepted Answer

To find the time, use the formula: time = (mass × n × F) / (current × M), where n is the number of electrons (2 for Ni²⁺), F is Faraday's constant (96485 C/mol), and M is the molar mass of Ni. Plugging in the values gives approximately 33.2 hours.

Question 20

Gold is produced electrochemically from an aqueous solution of Au(CN)₂^- containing an excess of CN^-. Gold metal and oxygen gas are produced at the electrodes. How many moles of O₂ will be produced during the production of 1.00 mol of gold?

Accepted Answer

The answer of Gold is produced electrochemically from an aqueous...

Quiz 11: Electrochemistry

Calculate E at 25°C for this cell, given the following data: Ag⁺ + e^- → Ag(s) E° = 0.80 V Ni²⁺ + 2e^- → Ni(s) E° = -0.23 V K_sp for AgCl = 1.6 × 10^-10

A solution of MnO₄^2- is electrolytically reduced to Mn³⁺. A current of 8.46 amp is passed through the solution for 15.0 minutes. What is the number of moles of Mn³⁺ produced in this process? (1 faraday = 96,485 coulombs)

If an electrolysis plant operates its electrolytic cells at a total current of 1.0 × 10⁶ amp, how long will it take to produce one metric ton (one million grams) of Mg(s) from seawater containing Mg²⁺? (1 faraday = 96,485 coulombs)

How many seconds would it take to deposit 21.40 g of Ag (atomic mass = 107.87) from a solution of AgNO₃ using a current of 10.00 amp?

If a constant current of 5.2 amperes is passed through a cell containing Cr³⁺ for 2.1 hour, how many grams of Cr will plate out onto the cathode? (The atomic mass of Cr is 51.996 g/mol.)

What quantity of charge is required to reduce 32.6 g of CrCl₃ to chromium metal? (1 faraday = 96,485 coulombs)

Use the following data to calculate the K_sp value at 25°C for PbSO₄(s) .

An electrolytic cell process involves plating Zr(s) from a solution containing Zr⁴⁺. If 5.80 amp is run through this mixture for 1.86 h, what mass of Zr is plated?

Electrolysis of a molten salt with the formula MCl, using a current of 3.86 amp for 16.2 min, deposits 1.52 g of metal. Identify the metal. (1 faraday = 96,485 coulombs)

Why is aluminum protected from corrosion? (Note: The standard reduction potential for Al³⁺ is -1.66 V.)

Copper is electroplated from an aqueous CuSO₄ solution. A constant current of 4.70 amp is applied by an external power supply. How long will it take to deposit 3.76 × 10² g of Cu? The atomic mass of copper is 63.546 g/mol.

Calculate the solubility product of silver iodide at 25°C, given the following data:

Gold (atomic mass = 197 g/mol) is plated from a solution of chlorauric acid, HAuCl₄; it deposits on the cathode. Calculate the time it takes to deposit 0.55 g of gold, passing a current of 0.12 amperes. (1 faraday = 96,485 coulombs)

Nickel is electroplated from a NiSO₄ solution. A constant current of 5.72 amp is applied by an external power supply. How long will it take to deposit 2.08 × 10² g of Ni? (The atomic mass of Ni is 58.69 g/mol.)

Gold is produced electrochemically from an aqueous solution of Au(CN) ₂^- containing an excess of CN^-. Gold metal and oxygen gas are produced at the electrodes. How many moles of O₂ will be produced during the production of 1.00 mol of gold?

Quiz 11: Electrochemistry

Calculate E at 25°C for this cell, given the following data: Ag+ + e- → Ag(s) E° = 0.80 V Ni2+ + 2e- → Ni(s) E° = -0.23 V Ksp for AgCl = 1.6 × 10-10

A solution of MnO42- is electrolytically reduced to Mn3+. A current of 8.46 amp is passed through the solution for 15.0 minutes. What is the number of moles of Mn3+ produced in this process? (1 faraday = 96,485 coulombs)

If an electrolysis plant operates its electrolytic cells at a total current of 1.0 × 106 amp, how long will it take to produce one metric ton (one million grams) of Mg(s) from seawater containing Mg2+? (1 faraday = 96,485 coulombs)

How many seconds would it take to deposit 21.40 g of Ag (atomic mass = 107.87) from a solution of AgNO3 using a current of 10.00 amp?

If a constant current of 5.2 amperes is passed through a cell containing Cr3+ for 2.1 hour, how many grams of Cr will plate out onto the cathode? (The atomic mass of Cr is 51.996 g/mol.)

What quantity of charge is required to reduce 32.6 g of CrCl3 to chromium metal? (1 faraday = 96,485 coulombs)

Use the following data to calculate the Ksp value at 25°C for PbSO4(s) .

An electrolytic cell process involves plating Zr(s) from a solution containing Zr4+. If 5.80 amp is run through this mixture for 1.86 h, what mass of Zr is plated?

Electrolysis of a molten salt with the formula MCl, using a current of 3.86 amp for 16.2 min, deposits 1.52 g of metal. Identify the metal. (1 faraday = 96,485 coulombs)

Why is aluminum protected from corrosion? (Note: The standard reduction potential for Al3+ is -1.66 V.)

Copper is electroplated from an aqueous CuSO4 solution. A constant current of 4.70 amp is applied by an external power supply. How long will it take to deposit 3.76 × 102 g of Cu? The atomic mass of copper is 63.546 g/mol.

Calculate the solubility product of silver iodide at 25°C, given the following data:

Gold (atomic mass = 197 g/mol) is plated from a solution of chlorauric acid, HAuCl4; it deposits on the cathode. Calculate the time it takes to deposit 0.55 g of gold, passing a current of 0.12 amperes. (1 faraday = 96,485 coulombs)

Nickel is electroplated from a NiSO4 solution. A constant current of 5.72 amp is applied by an external power supply. How long will it take to deposit 2.08 × 102 g of Ni? (The atomic mass of Ni is 58.69 g/mol.)

Gold is produced electrochemically from an aqueous solution of Au(CN) 2- containing an excess of CN-. Gold metal and oxygen gas are produced at the electrodes. How many moles of O2 will be produced during the production of 1.00 mol of gold?

Calculate E at 25°C for this cell, given the following data: Ag⁺ + e^- → Ag(s) E° = 0.80 V Ni²⁺ + 2e^- → Ni(s) E° = -0.23 V K_sp for AgCl = 1.6 × 10^-10

A solution of MnO₄^2- is electrolytically reduced to Mn³⁺. A current of 8.46 amp is passed through the solution for 15.0 minutes. What is the number of moles of Mn³⁺ produced in this process? (1 faraday = 96,485 coulombs)

If an electrolysis plant operates its electrolytic cells at a total current of 1.0 × 10⁶ amp, how long will it take to produce one metric ton (one million grams) of Mg(s) from seawater containing Mg²⁺? (1 faraday = 96,485 coulombs)

How many seconds would it take to deposit 21.40 g of Ag (atomic mass = 107.87) from a solution of AgNO₃ using a current of 10.00 amp?

If a constant current of 5.2 amperes is passed through a cell containing Cr³⁺ for 2.1 hour, how many grams of Cr will plate out onto the cathode? (The atomic mass of Cr is 51.996 g/mol.)

What quantity of charge is required to reduce 32.6 g of CrCl₃ to chromium metal? (1 faraday = 96,485 coulombs)

Use the following data to calculate the K_sp value at 25°C for PbSO₄(s) .

An electrolytic cell process involves plating Zr(s) from a solution containing Zr⁴⁺. If 5.80 amp is run through this mixture for 1.86 h, what mass of Zr is plated?

Why is aluminum protected from corrosion? (Note: The standard reduction potential for Al³⁺ is -1.66 V.)

Copper is electroplated from an aqueous CuSO₄ solution. A constant current of 4.70 amp is applied by an external power supply. How long will it take to deposit 3.76 × 10² g of Cu? The atomic mass of copper is 63.546 g/mol.

Gold (atomic mass = 197 g/mol) is plated from a solution of chlorauric acid, HAuCl₄; it deposits on the cathode. Calculate the time it takes to deposit 0.55 g of gold, passing a current of 0.12 amperes. (1 faraday = 96,485 coulombs)

Nickel is electroplated from a NiSO₄ solution. A constant current of 5.72 amp is applied by an external power supply. How long will it take to deposit 2.08 × 10² g of Ni? (The atomic mass of Ni is 58.69 g/mol.)

Gold is produced electrochemically from an aqueous solution of Au(CN) ₂^- containing an excess of CN^-. Gold metal and oxygen gas are produced at the electrodes. How many moles of O₂ will be produced during the production of 1.00 mol of gold?