Electrolysis — Revision Notes

⚡ 30-Second Revision

Electrolysis: — Electrical energy $\rightarrow$ Chemical energy (non-spontaneous).
Electrolytic Cell: — Anode (+ve, oxidation), Cathode (-ve, reduction).
Faraday's 1st Law: — $m = ZIt = \frac{MIt}{nF}$.
Faraday's 2nd Law: — $\frac{m_1}{m_2} = \frac{E_1}{E_2}$ (for series connection).
Faraday's Constant (F): — $96485,\text{C/mol}$ (approx $96500,\text{C/mol}$).
Equivalent Weight (E): — $M/n$.
Product Prediction: — Compare reduction potentials at cathode, oxidation potentials at anode. Consider concentration, overpotential, and electrode nature.
Water Reactions: — Cathode: $2H_2O + 2e^- \rightarrow H_2 + 2OH^-$. Anode: $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$.

2-Minute Revision

Electrolysis is the process of using electrical energy to drive non-spontaneous chemical reactions in an electrolytic cell. This cell consists of an electrolyte (molten ionic compound or aqueous solution) and two electrodes connected to a DC power source.

Oxidation always occurs at the anode (positive electrode), and reduction at the cathode (negative electrode). Faraday's First Law states that the mass of substance deposited ($m$) is directly proportional to the current ($I$) and time ($t$), given by $m = \frac{MIt}{nF}$, where $M$ is molar mass, $n$ is the number of electrons, and $F$ is Faraday's constant ($96500,\text{C/mol}$).

Faraday's Second Law applies when multiple cells are in series; the masses deposited are proportional to their equivalent weights ($E = M/n$). Product prediction in aqueous solutions is crucial: compare standard electrode potentials of all species present (ions from salt and water).

Remember that water can be reduced to $H_2$ at the cathode or oxidized to $O_2$ at the anode. Overpotential for gas evolution and the concentration of ions significantly influence the actual products.

Active electrodes can also participate in the reaction, especially at the anode.

5-Minute Revision

Electrolysis is the process of converting electrical energy into chemical energy by forcing a non-spontaneous redox reaction. This occurs in an electrolytic cell, comprising an electrolyte (molten or aqueous) and two electrodes connected to a DC power supply. The anode is the positive electrode where oxidation occurs, while the cathode is the negative electrode where reduction takes place.

Faraday's Laws are quantitative:

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First Law: — The mass ($m$) of a substance deposited or liberated is directly proportional to the quantity of electricity ($Q$) passed. $Q = It$. So, $m = ZIt$, where $Z$ is the electrochemical equivalent ($Z = M/nF$). The most common form is $m = \frac{MIt}{nF}$, where $M$ is molar mass, $I$ is current in Amperes, $t$ is time in seconds, $n$ is the number of electrons involved in the electrode reaction, and $F$ is Faraday's constant ($96500,\text{C/mol}$).
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Second Law: — When the same quantity of electricity is passed through different electrolytes connected in series, the masses of substances deposited are directly proportional to their equivalent weights ($E = M/n$). Thus, $\frac{m_1}{m_2} = \frac{E_1}{E_2}$.

Product Prediction (Aqueous Solutions): This is a key conceptual area. Consider all species present ($H_2O$, cations, anions).

At Cathode (Reduction): — Compare reduction potentials. The species with the higher (less negative) reduction potential will be reduced. For highly reactive metal ions ($Na^+, K^+, Ca^{2+}$), water is preferentially reduced ($2H_2O + 2e^- \rightarrow H_2 + 2OH^-$). For less reactive metals ($Cu^{2+}, Ag^+$), the metal ion is reduced.
At Anode (Oxidation): — Compare oxidation potentials. The species with the lower (less positive) oxidation potential will be oxidized. However, overpotential for gas evolution (especially $O_2$) and concentration effects are critical. For example, in concentrated $NaCl$, $Cl^-$ oxidizes to $Cl_2$ ($2Cl^- \rightarrow Cl_2 + 2e^-$) despite water having a slightly lower standard oxidation potential ($2H_2O \rightarrow O_2 + 4H^+ + 4e^-$). In dilute $NaCl$, $O_2$ is preferred. If the anode is active (e.g., $Cu$ anode in $CuSO_4$), the anode itself may oxidize ($Cu \rightarrow Cu^{2+} + 2e^-$).

Example: Electrolysis of molten $NaCl$: Cathode: $Na^+ + e^- \rightarrow Na$. Anode: $2Cl^- \rightarrow Cl_2 + 2e^-$. Example: Electrolysis of aqueous $CuSO_4$ with inert electrodes: Cathode: $Cu^{2+} + 2e^- \rightarrow Cu$. Anode: $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$.

Prelims Revision Notes

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Definition: — Electrolysis is the process of using electrical energy to drive non-spontaneous chemical reactions. It's an electrical to chemical energy conversion.
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Electrolytic Cell Components:

* Electrolyte: Molten ionic compound or aqueous solution containing free ions. * Electrodes: Conductors where redox reactions occur. * DC Power Source: Provides the electrical energy.

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Electrode Roles & Polarity:

* Anode: Positive electrode, site of oxidation (loss of electrons). * Cathode: Negative electrode, site of reduction (gain of electrons).

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Faraday's First Law:

* Mass ($m$) of substance deposited/liberated is directly proportional to the quantity of electricity ($Q$). * $Q = I \times t$ (Current in Amperes, time in seconds). * Formula: $m = ZIt = \frac{MIt}{nF}$. * $M$: Molar mass (g/mol). * $n$: Number of electrons involved in the electrode reaction (valency factor). * $F$: Faraday's constant ($96485,\text{C/mol}$ or $96500,\text{C/mol}$). One Faraday = charge of one mole of electrons.

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Faraday's Second Law:

* When the same quantity of electricity is passed through different electrolytes in series, the masses of substances deposited are proportional to their equivalent weights ($E$). * $E = M/n$. * Formula: $\frac{m_1}{m_2} = \frac{E_1}{E_2}$.

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Product Prediction in Aqueous Solutions (Key Factors):

* **Standard Electrode Potentials ($E^circ$):** Compare $E^circ$ values for all possible reactions. * Cathode (Reduction): Species with higher (less negative) $E^circ_{red}$ gets reduced. For $Na^+, K^+, Ca^{2+}$, water is reduced ($2H_2O + 2e^- \rightarrow H_2 + 2OH^-$).

For $Cu^{2+}, Ag^+$, metal ion is reduced. * Anode (Oxidation): Species with lower (less positive) $E^circ_{ox}$ (or higher $E^circ_{red}$) gets oxidized. Common oxidations: $2Cl^- \rightarrow Cl_2 + 2e^-$, $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$.

* Overpotential: Extra voltage required for gas evolution. Significant for $O_2$ and $H_2$. Can alter product prediction (e.g., $Cl_2$ over $O_2$ from concentrated $NaCl$). * Concentration: Higher concentration of an ion can favor its discharge (e.

g., concentrated $NaCl$ vs. dilute $NaCl$). * Nature of Electrodes: * Inert (Pt, Graphite): Do not participate in reaction. * Active (Cu, Ag): Can be oxidized at the anode if their oxidation potential is favorable (e.

g., $Cu$ anode in $CuSO_4$ solution, $Cu \rightarrow Cu^{2+} + 2e^-$).

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Important Reactions:

* Molten $NaCl$: Cathode: $Na(s)$, Anode: $Cl_2(g)$. * Aqueous concentrated $NaCl$ (inert electrodes): Cathode: $H_2(g)$, Anode: $Cl_2(g)$. * Aqueous dilute $NaCl$ (inert electrodes): Cathode: $H_2(g)$, Anode: $O_2(g)$. * Aqueous $CuSO_4$ (inert electrodes): Cathode: $Cu(s)$, Anode: $O_2(g)$. * Aqueous $CuSO_4$ (Cu anode): Cathode: $Cu(s)$, Anode: $Cu^{2+}(aq)$ (anode dissolves).

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Stoichiometry: — Relate moles of electrons to moles of product (e.g., 4 Faradays for 1 mole of $O_2$, 2 Faradays for 1 mole of $H_2$).

Vyyuha Quick Recall

An Ox, Red Cat: Anode = Oxidation; Reduction = Cathode. For electrolytic cells, remember 'PANIC': Positive Anode, Negative In Cathode.