redox reactions - Faradays laws and calculation in electrochemistry- exercises

1) A spoon was placed in a silver plating cell for 40 minutes with a current of 30.3A. Calculate the mass of silver deposited on the spoon.

Step 1 Calculate the charge passed through the cell during the 15 minute period.

Q = I X t
Q = 30.3 X (40 X 60)s = 72,720C
(time must be expressed in seconds)

Step 2 Calculate the mole of electrons passed through the circuit.

mole = 72,720C/96,500C/mol = 0.754 mole.

Step 3 Calculate the mass of silver

According to the equation above 0.754 mole of electrons will produce 0.754 mole of silver.
The mass of silver is therefore 0.754 X 107.9 = 81.31grams.

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A copper plating cell operates for 20 minutes with a steady current of 25.0A. What mass of copper is deposited?

Step 1 Calculate the charge passed through the cell during the 20 minute period.

Q = I X t
Q = 25.0 X (20 X 60)s = 30,000C
(time must be expressed in seconds)

Step 2 Calculate the mole of electrons passed through the circuit.

mole = 30,000C/96,500C/mol = 0.311 mole.

Step 3 Calculate the mass of copper

According to the equation above 0.311 mole of electrons will produce 0.311/2 mole of copper
The mass of copper is therefore 0.155 X 63.54 = 9.88grams.

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An iron window frame is placed in an aluminium plating cell for 53 minutes with a steady current of 34.5A. What mass of aluminium is deposited?

Step 1 Calculate the charge passed through the cell during the 53 minute period.

Q = I X t
Q = 34.5 X (53 X 60)s = 109,710C
(time must be expressed in seconds)

Step 2 Calculate the mole of electrons passed through the circuit.

mole = 109,710C/96,500C/mol = 1.14 mole.

Step 3 Calculate the mass of aluminium

According to the equation above 1.14 mole of electrons will produce 0.1.14/3 mole of aluminium.
The mass of aluminium is therefore 0.379 X 27 = 10.23grams.

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60g of silver must be evenly deposited on a serving tray. The iron serving tray is placed in a silver plating cell with a steady current of 45.5A. How long should the tray be left in the cell for?

Step 1 Calculate the mole of silver deposited

mole =m/M = 60/107 = 0.561mole

Step 2 Calculate the mole of electrons passed through the circuit.
According to the equation below one mole of silver needs one mole of eelctrons to deposit.

mole = 0.561mol of silver= 0.561 mol of electrons

Step 3. Calculate the charge of 0.561 mol of electrons

0.561 X 96,500 = 54,136.5

Step 4 Calculate the time needed.

Q =I X t
54,136.5 = 45.5 X t
54,136.5/45.5 = t = 1,190s = 19.8 minutes

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An extruded iron pipe is placed in a copper plating cell. A mass of 300 grams of copper must be deposited evenly along its length. The pipe is left in the cell for 1 hour, to what current should the operator set the dial?

Step 1 Calculate the mole of copper

300 / 63.5 = 4.72 mole

Step 2 Calculate the mole of electrons needed.

According to the equation above we need twice as many mole of electrons as mole of copper.
mole of electrons = 4.72 X 2 = 9.45mole

Step 3 Calculate the charge of 9.45mole of electrons

Q = mole of electrons X 96,500 = 9.45 X 96,500 = 911,811C

Step 4 Calculate the current needed.

911,811 = I X 60 X 60

I =911,811/3600 = 253.2A

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An iron serving tray has a surface area of 350 square centimetres. Silver is to be deposited to a depth of 4mm in an silver plating cell. If the current delivered is a steady 34.6A how long should the operator keep the tray in the cell for?
Density of silver =

Step 1 Calculate the mass of silver needed.

Density X volume = mass
The volume of silver deposited is 350 X 0.4 =140 cubic cm.
10.54 X 140 = 1475.6 grams

Step 2 Calculate the mole of silver

1475.6 /107.9 =13.68mole of silver

Step 3 Calculate the mole of electrons needed

13.68 mole of electrons is needed according to the equation above.

Step 4 Calculate the charge on the 13.68 mole of electrons

13.68 X 96.500 = 13,196,978C

Step 5 Calculate the time needed.

Q =I X t

13,196,978 = 34.6 X t

t = 13,196,978/34.6 = 38141s = 10.6 hours

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A teapot with a surface area of 0.05 square metres is placed in a silver plating cell for 34 minutes with a steady current of 1.2A. If the silver is deposited evenly how deep is the layer deposited on the surface of the teapot?
Density of silver =

Step 1 Calculate the charge delivered.

Q =I X t
Q = 1.2 X 34 X 60= 2448C

Step 2 Calculate the mole of electrons

2448/96500=0.0254mole

Step 3 Calculate the mole of silver needed

0.0245 mole of silver is needed according to the equation above.

Step 4 Calculate the mass silver

0.0245 X 107.9 = 2.64g

Step 5 Calculate the volume of silver that is required

volume =mass / density

volume = 2.64/10.54 =0.25 cubic cm

Step 6 Calculate the depth

Volume = area x depth
depth(cm) = 0.25 / 500 sq cm =0.0005cm

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An unknown metal which forms ions with a charge of 3+ () was deposited on a small key ring in an electrolytic cell. A steady current of 4.6A was used for 34 minutes and a mass of 1.815 grams was deposited on the key ring. Calculate the relative atomic mass of the element (X) and attempt to identify it.

Step 1 Calculate the charge delivered

Q = I X t
4.6 X 34 X 60 = 9384C

Step 2 Calculate the mole of electrons delivered

9384 /96500 =0.097mole

Step 3 Calculate the mole of metal "X" deposited according to the reaction below.

0.097 / 3 = 0.0324 mole of "X"

Step 4 Calculate the relative atomic mass of "X"

mole = mass/ rel.atomic mass

mole = 1.815 / M

M =1.815/0.0324

M =56

It is most likely iron

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A metal with a cation of charge 2+ was deposited in an electrolytic cell using a steady current of 200A for 55 minutes. If 136.79 grams was deposited identify the metal.

Step 1 Calculate the charge delivered

Q = I X t
200 X 55 X 60 = 660,000C

Step 2 Calculate the mole of electrons delivered

660,000 /96500 =6.84mole

Step 3 Calculate the mole of metal "X" deposited according to the reaction below.

6.84 / 2 = 3.42 mole of "X"

Step 4 Calculate the relative atomic mass of "X"

mole = mass/ rel.atomic mass

mole = 136.79 / M

M =136.79/3.42

M = 40.0

It is most likely calcium

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Using Faraday in electrochemical calculations.

Exercises

1) A spoon was placed in a silver plating cell for 40.0 minutes with a current of 30.3A. Calculate the mass of silver deposited on the spoon.
Solution

2) A copper plating cell operates for 20.0 minutes with a steady current of 25.0A. What mass of copper is deposited?
Solution

3) An iron window frame is placed in an aluminium plating cell for 53.0 minutes with a steady current of 34.5A. What mass of aluminium is deposited?
Solution

4) 60.0g of silver must be evenly deposited on a serving tray. The iron serving tray is placed in a silver plating cell with a steady current of 45.5A. How long should the tray be left in the cell for?
Solution

5) An extruded iron pipe is placed in a copper plating cell. A mass of 300.00 grams of copper must be deposited evenly along its length. The pipe is left in the cell for 1.00 hour, to what current should the operator set the dial?
Solution

6) An iron serving tray has a surface area of 350.0 square centimetres. Silver is to be deposited to a depth of 4.00 mm in an silver plating cell. If the current delivered is a steady 34.6A how long should the operator keep the tray in the cell for?
Density of silver = 10.54g/cm³
Solution

7) A teapot with a surface area of 0.0500 square metres is placed in a silver plating cell for 34.0 minutes with a steady current of 1.20A. If the silver is deposited evenly how deep is the layer deposited on the surface of the teapot?
Density of silver =10.54g/cm³
Solution

8) An unknown metal which forms ions with a charge of 3+ (X⁺³) was deposited on a small key ring in an electrolytic cell. A steady current of 4.60A was used for 34.0 minutes and a mass of 1.815 grams was deposited on the key ring. Calculate the relative atomic mass of the element (X) and attempt to identify it.
Solution

9) A metal with a cation of charge 2+ was deposited in an electrolytic cell using a steady current of 200.0A for 55.0 minutes. If 136.79 grams was deposited identify the metal.
Solution

Using Faraday in electrochemical calculations. Exercises
1) A spoon was placed in a silver plating cell for 40.0 minutes with a current of 30.3A. Calculate the mass of silver deposited on the spoon. Solution 2) A copper plating cell operates for 20.0 minutes with a steady current of 25.0A. What mass of copper is deposited? Solution 3) An iron window frame is placed in an aluminium plating cell for 53.0 minutes with a steady current of 34.5A. What mass of aluminium is deposited? Solution 4) 60.0g of silver must be evenly deposited on a serving tray. The iron serving tray is placed in a silver plating cell with a steady current of 45.5A. How long should the tray be left in the cell for? Solution 5) An extruded iron pipe is placed in a copper plating cell. A mass of 300.00 grams of copper must be deposited evenly along its length. The pipe is left in the cell for 1.00 hour, to what current should the operator set the dial? Solution 6) An iron serving tray has a surface area of 350.0 square centimetres. Silver is to be deposited to a depth of 4.00 mm in an silver plating cell. If the current delivered is a steady 34.6A how long should the operator keep the tray in the cell for? Density of silver = 10.54g/cm³ Solution 7) A teapot with a surface area of 0.0500 square metres is placed in a silver plating cell for 34.0 minutes with a steady current of 1.20A. If the silver is deposited evenly how deep is the layer deposited on the surface of the teapot? Density of silver =10.54g/cm³ Solution 8) An unknown metal which forms ions with a charge of 3+ (X⁺³) was deposited on a small key ring in an electrolytic cell. A steady current of 4.60A was used for 34.0 minutes and a mass of 1.815 grams was deposited on the key ring. Calculate the relative atomic mass of the element (X) and attempt to identify it. Solution 9) A metal with a cation of charge 2+ was deposited in an electrolytic cell using a steady current of 200.0A for 55.0 minutes. If 136.79 grams was deposited identify the metal. Solution
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