Electrolysis uses electrical energy to drive a chemical reaction that would not happen on its own. It is important in metal refining, electroplating, production of elements, and rechargeable battery chemistry. In an electrolytic cell, a DC power supply pushes electrons through an external circuit while ions move through the electrolyte.
Faraday's laws let us predict how much substance is produced from the current and time.
Understanding Chemistry: Electrolysis and Faraday's Laws
At each electrode, particles can gain or lose electrons only if they reach the surface. Positive ions move toward the cathode, where they receive electrons. Negative ions move toward the anode, where they give up electrons.
The electrode material can take part in the reaction or simply provide a conducting surface. A graphite electrode is often chosen because it is fairly unreactive.
A copper anode can dissolve during copper purification, replacing copper ions that are removed at the cathode. This links what happens at the two electrodes through the movement of charge.
The electrolyte makes a major difference. In a molten ionic compound, the only ions present come from that compound. Molten lead bromide, for example, can form lead at one electrode and bromine at the other.
An aqueous solution is less simple because water can react too. At the cathode, a metal ion may be deposited, or hydrogen gas may form from water. At the anode, a negative ion may be discharged, or oxygen gas may form from water or hydroxide ions.
The outcome depends on the type of ion, its concentration, and the electrode material. Students should learn electrode reactions rather than assuming every ion in solution becomes an element.
Faraday's laws connect a visible result to electron transfer. A larger current sends more electrons each second, so it can produce more material in a fixed time. Running the same current for longer has the same effect.
The key step is finding how many electrons each particle needs. Silver ions need one electron to make silver metal. Copper ions with a two positive charge need two electrons.
Aluminium ions with a three positive charge need three electrons. For the same charge passed, ions needing fewer electrons produce more moles of metal. Molar mass then determines the mass measured on a balance.
Units matter throughout. Current must be in amperes and time must be in seconds before calculating charge in coulombs.
Real electrolysis is not always perfectly efficient. Some current can cause unwanted reactions, heat the solution, or form gases instead of the intended product. A calculated mass is therefore often greater than the mass actually collected.
This difference is described by current efficiency. In electroplating, poor control can make a coating rough, thin, or uneven. Clean surfaces, suitable ion concentration, temperature, and current density help give a better deposit.
These ideas appear in chrome plated objects, copper coated circuit boards, aluminium extraction, and purification of metals. When solving problems, identify the electrode products first, write the electron change in words, then connect charge, moles of electrons, moles of product, and mass in that order.
Key Facts
- Charge passed: Q = It, where Q is in coulombs, I is in amperes, and t is in seconds.
- Faraday constant: F = 96485 C mol^-1, the charge carried by 1 mole of electrons.
- Moles of electrons transferred: n(e-) = Q/F.
- For copper plating: Cu^2+ + 2e- -> Cu(s), so 2 moles of electrons deposit 1 mole of copper.
- Mass deposited: m = (ItM)/(zF), where M is molar mass and z is electrons per ion.
- In an electrolytic cell, oxidation occurs at the positive anode and reduction occurs at the negative cathode.
Vocabulary
- Electrolysis
- Electrolysis is the use of electrical energy to force a nonspontaneous chemical reaction to occur.
- Electrolyte
- An electrolyte is a solution or molten substance containing mobile ions that carry charge through the cell.
- Anode
- The anode is the electrode where oxidation occurs, and in an electrolytic cell it is connected to the positive terminal.
- Cathode
- The cathode is the electrode where reduction occurs, and in an electrolytic cell it is connected to the negative terminal.
- Faraday constant
- The Faraday constant is the charge of one mole of electrons, approximately 96485 coulombs per mole.
Common Mistakes to Avoid
- Using minutes directly in Q = It is wrong because time must be in seconds when current is in amperes.
- Forgetting the electron coefficient z gives the wrong mass because ions with different charges require different numbers of electrons per atom or ion deposited.
- Calling the cathode positive in electrolysis is wrong because in an electrolytic cell the cathode is connected to the negative terminal even though reduction still occurs there.
- Assuming all current always plates the desired metal can be wrong because side reactions, impurities, or gas formation can reduce current efficiency.
Practice Questions
- 1 A current of 2.00 A flows through a CuSO4(aq) electrolytic cell for 30.0 min. How many grams of Cu are plated at the cathode? Use M(Cu) = 63.55 g/mol and z = 2.
- 2 How long, in minutes, must a 5.00 A current run to deposit 1.50 g of Ag from Ag+ solution? Use M(Ag) = 107.87 g/mol and z = 1.
- 3 In a copper electroplating cell with a copper anode and a metal object as the cathode, explain why the copper anode loses mass while the cathode gains a copper coating.