Electrolysis Revision Notes

Ionic compounds consist of positive and negative ions held together by strong electrostatic forces in a lattice structure. In the solid state, ions are fixed in position and cannot move freely, so ionic compounds do not conduct electricity.

To conduct electricity and undergo electrolysis, ionic compounds must be either molten (melted by heating) or dissolved in water (aqueous solution). In these states, the ionic lattice breaks down, allowing ions to move freely and carry charge through the liquid.

For example, molten sodium chloride (NaCl) contains free Na⁺ and Cl⁻ ions able to migrate under an electric field.

• Ionic compounds contain positive and negative ions.
• Solid ionic compounds do not conduct electricity due to fixed ions.
• Molten or aqueous ionic compounds conduct electricity because ions are free to move.

Electrolysis involves passing an electric current through a molten or aqueous ionic compound. The electric current is supplied by an external power source connected to two electrodes immersed in the electrolyte.

Positive ions (cations) migrate towards the cathode (negative electrode) where they gain electrons (reduction). Negative ions (anions) migrate towards the anode (positive electrode) where they lose electrons (oxidation).

This movement of ions and electron transfer at electrodes causes the ionic compound to break down into its elements or simpler substances.

• Electric current drives ion movement in molten or aqueous ionic compounds.
• Cations move to the cathode and gain electrons (reduction).
• Anions move to the anode and lose electrons (oxidation).

The electric current provides the energy needed to overcome the strong ionic bonds in the compound, causing it to decompose.

Electrons flow from the power supply into the cathode and out of the anode, completing the circuit.

The flow of electrons at the electrodes causes ions to gain or lose electrons, resulting in chemical changes that separate the compound into elements or new substances.

• Electric current supplies energy for ionic bond breakdown.
• Electron flow occurs through the external circuit and electrodes.
• Electron transfer at electrodes causes ion discharge and compound decomposition.

In molten sodium chloride, NaCl dissociates into Na⁺ and Cl⁻ ions.

At the cathode, Na⁺ ions gain electrons to form sodium metal: Na⁺ + e⁻ → Na (reduction).

At the anode, Cl⁻ ions lose electrons to form chlorine gas: 2Cl⁻ → Cl₂ + 2e⁻ (oxidation).

• Molten NaCl contains free Na⁺ and Cl⁻ ions.
• Cathode reaction: Na⁺ + e⁻ → Na (metal formed).
• Anode reaction: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas formed).

During electrolysis of molten ionic compounds, metal is deposited at the cathode and non-metal gases may be released at the anode.

The electrolyte must be molten or aqueous to allow ion mobility.

Electrodes are usually inert (e.g., graphite or platinum) to avoid unwanted reactions with the products or electrolyte.

• Metal forms at cathode; non-metal gases may form at anode.
• Electrolyte must be molten or dissolved for ion movement.
• Inert electrodes prevent side reactions during electrolysis.

An electrolytic cell is a device that uses electrical energy to drive a non-spontaneous chemical reaction called electrolysis. It consists of three fundamental components: two electrodes and an electrolyte. Understanding these components is essential to grasp how electrolysis occurs and how substances are decomposed or deposited during the process.

• Electrolysis requires an external power source to force the reaction.
• The cell contains two electrodes: anode and cathode.
• The electrolyte is the ionic medium that allows charge flow.

The anode is the electrode connected to the positive terminal of the external power supply. It carries a positive charge during electrolysis. Because it is positive, it attracts negatively charged ions (anions) from the electrolyte.

At the anode, oxidation occurs — this means electrons are lost by ions or atoms. These electrons flow out of the anode into the external circuit.

For example, in the electrolysis of molten lead(II) bromide (PbBr₂), bromide ions (Br⁻) migrate to the anode, where each loses an electron to form bromine atoms, which then pair to form bromine molecules (Br₂). The half-equation at the anode is: 2Br⁻ → Br₂ + 2e⁻.

• Anode is connected to the positive terminal of the power supply.
• Anode attracts anions (negative ions).
• Oxidation occurs at the anode (loss of electrons).
• Example half-equation: 2Br⁻ → Br₂ + 2e⁻.

The cathode is the electrode connected to the negative terminal of the power supply and carries a negative charge during electrolysis. It attracts positively charged ions (cations) from the electrolyte.

At the cathode, reduction occurs — electrons are gained by ions or atoms. Electrons flow from the external circuit into the cathode and then to the ions in the electrolyte.

For instance, in molten lead(II) bromide, lead ions (Pb²⁺) move to the cathode, where each ion gains two electrons to form lead atoms. The half-equation at the cathode is: Pb²⁺ + 2e⁻ → Pb.

• Cathode is connected to the negative terminal of the power supply.
• Cathode attracts cations (positive ions).
• Reduction occurs at the cathode (gain of electrons).
• Example half-equation: Pb²⁺ + 2e⁻ → Pb.

The electrolyte is the ionic compound in molten or aqueous form that undergoes electrolysis. It contains free ions that can move and carry charge between the electrodes.

In molten electrolytes, the ionic compound is melted to allow ions to move freely. In aqueous electrolytes, the ionic compound is dissolved in water, dissociating into ions.

The electrolyte completes the electrical circuit by allowing ions to migrate to the electrodes: cations to the cathode and anions to the anode. This ion movement maintains electrical neutrality in the solution during electrolysis.

• Electrolyte contains mobile ions to conduct electricity.
• Can be molten ionic compound or aqueous ionic solution.
• Allows ion migration to electrodes during electrolysis.
• Maintains charge balance in the cell.

During electrolysis, electrons flow from the power supply into the cathode and out of the anode back to the power supply. Ions in the electrolyte move to the electrodes to balance charge transfer.

Cations (positive ions) migrate towards the cathode to gain electrons (reduction). Anions (negative ions) migrate towards the anode to lose electrons (oxidation).

This coordinated movement of electrons and ions enables the decomposition or formation of substances at the electrodes.

• Electrons flow: power supply → cathode → electrolyte → anode → power supply.
• Cations move to cathode; anions move to anode.
• Oxidation at anode; reduction at cathode.
• Electrolyte ion movement maintains electrical neutrality.

Lead(II) bromide is an ionic compound composed of Pb²⁺ and Br⁻ ions. When molten, it conducts electricity because the ions are free to move.

During electrolysis, the molten PbBr₂ decomposes into lead metal and bromine gas.

At the cathode (negative electrode), Pb²⁺ ions gain electrons (reduction) to form lead atoms, which deposit as a silvery liquid or solid metal: Pb²⁺(l) + 2e⁻ → Pb(l). Lead has a high melting point but molten lead can be observed if the temperature is sufficiently high; otherwise, solid lead deposits on the cathode surface.

• At the cathode: Pb²⁺ ions are reduced to lead metal.
• At the anode: Br⁻ ions are oxidised to bromine gas.
• Electrodes are inert (platinum or graphite) and do not react.
• The electrolyte must be molten because solid PbBr₂ does not conduct electricity.

At the cathode, lead metal forms and may be seen as a silvery deposit.

At the anode, bromide ions lose electrons (oxidation) to form bromine molecules, which are released as brown fumes or gas bubbles: 2Br⁻(l) → Br₂(g) + 2e⁻.

The overall reaction is: PbBr₂(l) → Pb(l) + Br₂(g).

• Brown bromine vapour is visible at the anode.
• Lead metal collects at the cathode.
• No other products form because the compound is molten and no water is present.

In concentrated aqueous sodium chloride solution, the electrolyte contains Na⁺, Cl⁻, H⁺, and OH⁻ ions due to water ionisation.

In electrolysis with inert electrodes, the products depend on the relative ease of discharge of ions.

At the cathode, hydrogen ions (H⁺) are preferentially reduced over sodium ions because hydrogen ions require less energy to gain electrons: 2H⁺(aq) + 2e⁻ → H₂(g). Hydrogen gas bubbles form at the cathode.

• At the cathode: hydrogen gas is produced.
• At the anode: chloride ions are oxidised to chlorine gas.
• Sodium ions and hydroxide ions remain in solution, increasing the solution's alkalinity.
• Chlorine gas appears as greenish-yellow bubbles at the anode.

At the anode, chloride ions lose electrons to form chlorine gas: 2Cl⁻(aq) → Cl₂(g) + 2e⁻.

Chlorine gas is visible as greenish-yellow bubbles with a pungent smell.

At the cathode, hydrogen gas bubbles form due to reduction of H⁺ ions from water: 2H⁺(aq) + 2e⁻ → H₂(g).

• The solution becomes alkaline due to leftover OH⁻ ions.
• No sodium metal forms because it is more difficult to reduce than hydrogen ions in aqueous solution.
• The overall reaction is: 2NaCl(aq) + 2H₂O(l) → H₂(g) + Cl₂(g) + 2NaOH(aq).

Dilute sulfuric acid ionises completely into H⁺ and SO₄²⁻ ions in aqueous solution.

During electrolysis with inert electrodes, the products are hydrogen gas at the cathode and oxygen gas at the anode.

At the cathode, hydrogen ions gain electrons to form hydrogen gas: 2H⁺(aq) + 2e⁻ → H₂(g).

• At the anode, hydroxide ions from water are oxidised to oxygen gas: 4OH⁻(aq) → O₂(g) + 2H₂O(l) + 4e⁻.
• Oxygen gas bubbles appear at the anode, often with a faint smell.
• No sulfur-containing gases form because sulfate ions are not discharged under these conditions.
• The overall reaction is: 2H₂O(l) → 2H₂(g) + O₂(g).

Electrolysis is the process of breaking down an ionic compound by passing an electric current through its molten or aqueous state. The apparatus contains two electrodes: the cathode, which is the negative electrode attracting positive ions (cations), and the anode, the positive electrode attracting negative ions (anions).

At the cathode, reduction occurs as cations gain electrons. At the anode, oxidation occurs as anions lose electrons. The products formed depend on the ions present in the electrolyte.

• Cathode: negative electrode; site of reduction (gain of electrons).
• Anode: positive electrode; site of oxidation (loss of electrons).

At the cathode, cations are attracted and reduced. If the electrolyte contains metal ions less reactive than hydrogen, the metal is deposited at the cathode as a solid.

If the metal is more reactive than hydrogen or if hydrogen ions (H⁺) are present (for example, in aqueous solutions), hydrogen gas is produced instead.

This is because hydrogen ions are easier to reduce than highly reactive metal ions, so hydrogen gas forms preferentially in aqueous solutions containing such metals.

• Metal ions (e.g., Cu²⁺) gain electrons to form solid metal: Cu²⁺ + 2e⁻ → Cu (s).
• Hydrogen ions (H⁺) gain electrons to form hydrogen gas: 2H⁺ + 2e⁻ → H₂ (g).
• Hydrogen gas bubbles can be observed at the cathode if hydrogen is produced.

At the anode, anions are attracted and oxidised. Non-metal ions lose electrons to form non-metal elements.

For example, halide ions (Cl⁻, Br⁻, I⁻) lose electrons to form halogen gases (Cl₂, Br₂, I₂).

If the electrolyte contains no halide ions, hydroxide ions (OH⁻) from water are oxidised to form oxygen gas and water. Hydrogen gas is never produced at the anode during electrolysis.

• Halide ions oxidation: 2Cl⁻ → Cl₂ (g) + 2e⁻.
• Hydroxide ions oxidation: 4OH⁻ → O₂ (g) + 2H₂O + 4e⁻.
• Oxygen gas bubbles or halogen gas colour can be observed at the anode.

Copper(II) sulfate solution contains Cu²⁺ and SO₄²⁻ ions, plus H⁺ and OH⁻ ions from water.

At the cathode, Cu²⁺ ions are reduced to copper metal because copper is less reactive than hydrogen: Cu²⁺ + 2e⁻ → Cu (s).

At the anode, sulfate ions (SO₄²⁻) are not discharged; instead, hydroxide ions are oxidised to oxygen gas: 4OH⁻ → O₂ (g) + 2H₂O + 4e⁻.

• Cathode reaction: Cu²⁺ + 2e⁻ → Cu (s) (copper deposits on cathode).
• Anode reaction: 4OH⁻ → O₂ (g) + 2H₂O + 4e⁻ (oxygen gas bubbles at anode).
• Observation: copper metal forms on cathode; oxygen bubbles form at anode.

The products formed at each electrode depend on the ions present and their relative ease of discharge.

At the cathode, metals or hydrogen gas form depending on metal reactivity and presence of H⁺ ions.

At the anode, non-metal elements (halogens or oxygen) form, never hydrogen gas.

• Cathode: metal deposition or hydrogen gas evolution.
• Anode: halogen gas or oxygen gas evolution.
• Observations include metal plating, gas bubbles, and colour changes depending on products.

Electrolysis is the process of using an electric current to break down an ionic compound into its elements. When a binary ionic compound (composed of a metal and a non-metal) is melted to form a molten liquid, the ions are free to move and conduct electricity. Passing an electric current through the molten compound causes the compound to decompose into its constituent elements at the electrodes.

In molten state, the compound contains only its ions, without water or other solvents, so the products formed depend solely on the ions present in the compound. This makes predicting the products straightforward compared to electrolysis of aqueous solutions.

• Molten binary compounds contain free metal cations and non-metal anions.
• Electric current causes ions to move to electrodes where redox reactions occur.
• Electrolysis decomposes the compound into elemental metal and non-metal.

In electrolysis, the electrodes are connected to an external power supply. The cathode is the negative electrode and attracts cations (positive ions), while the anode is the positive electrode and attracts anions (negative ions).

For a molten binary compound: the metal ions (cations) move to the cathode, where they gain electrons (reduction) to form the metal element. The non-metal ions (anions) move to the anode, where they lose electrons (oxidation) to form the non-metal element.

This movement and electron transfer cause the compound to break down into its elements.

• Cathode (negative): attracts metal cations, site of reduction.
• Anode (positive): attracts non-metal anions, site of oxidation.
• Electrons flow from anode to cathode through the external circuit.

At the cathode, metal ions gain electrons to form metal atoms. For example, if the metal ion is M^n+, the half-equation is:

M^{n+} + n e^- → M (metal) (reduction) At the anode, non-metal ions lose electrons to form non-metal atoms or molecules. For example, if the non-metal ion is X^-, the half-equation is:

2 X^- → X_2 + 2 e^- (oxidation) These half-equations must be balanced for both mass and charge.

• Cathode half-equation: M^{n+} + n e^- → M
• Anode half-equation: 2 X^- → X_2 + 2 e^-
• Electrons are conserved by flowing through the external circuit.

Lead(II) bromide (PbBr_2) is a binary ionic compound composed of Pb^{2+} and Br^- ions. When molten, it conducts electricity and undergoes electrolysis.

At the cathode, Pb^{2+} ions gain two electrons to form lead metal: Pb^{2+} + 2 e^- → Pb (s)

At the anode, bromide ions lose electrons to form bromine gas: 2 Br^- → Br_2 (g) + 2 e^- Observations include the formation of molten lead metal at the cathode and brown bromine vapour bubbles at the anode.

• Cathode: Pb^{2+} + 2 e^- → Pb (metal formed)
• Anode: 2 Br^- → Br_2 (brown gas formed)
• Observation: molten lead collects at cathode; bromine gas bubbles at anode.

To predict products during electrolysis of molten binary compounds: identify the metal cation and non-metal anion, determine their charges, and write half-equations for reduction at the cathode and oxidation at the anode.

Remember that metals form at the cathode by reduction, and non-metals form at the anode by oxidation. Observations such as gas evolution or metal deposition confirm the products.

Balanced half-equations and correct identification of ions are essential for full marks in exams.

• Metal cations reduced at cathode to metal atoms.
• Non-metal anions oxidised at anode to non-metal elements.
• Write balanced ionic half-equations for electrode reactions.
• Use observations to support predictions.

Electroplating is used primarily to improve the appearance of metal objects by depositing a thin, shiny layer of metal onto their surface. This layer can be of a more attractive or valuable metal, such as silver or gold, enhancing the object's aesthetic appeal.

Besides appearance, electroplating increases the resistance of the metal object to corrosion. The plated metal acts as a protective barrier, preventing oxygen and moisture from reaching the underlying metal and thus reducing rust or tarnish formation.

• Enhances visual appeal by coating with a shiny metal layer.
• Protects the base metal from corrosion by forming a barrier.
• Can also improve surface hardness and wear resistance.

Electroplating involves an electrolytic cell where the object to be plated acts as the cathode (negative electrode). The metal to be deposited is usually the anode (positive electrode) or is supplied as ions in the electrolyte solution.

When an electric current passes through the electrolyte, metal ions in solution are reduced at the cathode and deposit as a solid metal layer on the object.

At the anode, metal atoms lose electrons (oxidation) and dissolve into the electrolyte as metal ions, maintaining the concentration of metal ions in solution.

• Object to be plated is the cathode (negative electrode).
• Anode is made of the plating metal or inert electrode with metal ions in solution.
• Electric current causes metal ions to gain electrons and deposit on the cathode.
• Anode metal dissolves to replenish metal ions in the electrolyte.

Gold, silver, and nickel are commonly used metals for electroplating due to their desirable properties.

Gold plating provides a corrosion-resistant, attractive finish often used in jewellery and electronics.

Silver plating offers a bright, reflective surface and good conductivity, used in cutlery and electrical contacts, but it can tarnish over time without protection from corrosion products like sulphides. Nickel plating is widely used for its hardness, corrosion resistance, and ability to provide a smooth finish on objects such as car parts and household fixtures.

• Gold: corrosion-resistant, decorative, used in jewellery and electronics.
• Silver: bright, conductive, used in cutlery and electrical contacts.
• Nickel: hard, corrosion-resistant, used for protective coatings and decorative finishes.

At the cathode, metal ions in solution gain electrons and deposit as solid metal:

Example: deposition of silver from silver nitrate solution:

Ag⁺(aq) + e⁻ → Ag(s) (reduction at cathode) At the anode, metal atoms lose electrons and enter solution as ions: Ag(s) → Ag⁺(aq) + e⁻ (oxidation at anode) This maintains the concentration of metal ions in the electrolyte, allowing continuous plating without depletion of ions.

• Cathode reaction: metal ion + electron → metal (solid)
• Anode reaction: metal (solid) → metal ion + electron
• Overall process deposits metal onto the object while replenishing ions in solution.

Electroplating requires careful control of variables such as current, time, and electrolyte concentration to ensure an even, adherent metal coating.

The thickness of the plating depends on the amount of charge passed, which can be calculated using Faraday's laws of electrolysis.

Electroplating is widely used in jewellery, automotive parts, electronics, and cutlery to improve appearance and durability.

• Control current and time to regulate plating thickness.
• Use appropriate electrolyte containing metal ions of plating metal.
• Applications include decorative finishes, corrosion protection, and improved wear resistance.

Electroplating involves an electrolytic cell consisting of three main components: the cathode, the anode, and the electrolyte solution.

The object to be electroplated acts as the cathode (negative electrode).

The anode (positive electrode) is usually made of the metal that will be deposited onto the cathode. This metal dissolves into the electrolyte to replenish metal ions in solution during electroplating, maintaining a constant concentration of metal ions in the electrolyte solution. Alternatively, an inert anode such as platinum can be used, but then metal ions must be supplied by the electrolyte itself.

• Cathode: object to be coated, connected to the negative terminal.
• Anode: metal to be deposited, connected to the positive terminal.
• Electrolyte: aqueous solution containing metal ions of the plating metal.

When the electric current passes through the electrolyte, metal ions in solution migrate towards the cathode where they gain electrons (are reduced) and deposit as solid metal on the object.

At the anode, metal atoms lose electrons (are oxidised) and enter the solution as metal ions, maintaining the electrolyte concentration.

This continuous transfer of metal ions from anode to cathode allows a uniform metal coating to form on the cathode surface.

• Cathode reaction (reduction): Mⁿ⁺(aq) + n e⁻ → M(s)
• Anode reaction (oxidation): M(s) → Mⁿ⁺(aq) + n e⁻
• Overall: metal atoms move from anode to cathode via ionic form in solution.

Consider electroplating a metal object with copper using copper electrodes and copper(II) sulfate solution as the electrolyte.

At the cathode (object): Cu²⁺(aq) + 2 e⁻ → Cu(s) (copper ions gain electrons and deposit as copper metal).

At the anode (copper electrode): Cu(s) → Cu²⁺(aq) + 2 e⁻ (copper metal dissolves into solution as ions).

• The electrolyte remains copper(II) sulfate solution throughout the process.
• The mass of the cathode increases as copper is deposited.
• The mass of the anode decreases as copper dissolves.

The electrolyte must contain a sufficient concentration of metal ions to allow continuous plating.

The current must be controlled to ensure a smooth, even coating without roughness or burning.

The temperature of the electrolyte can affect the plating rate and quality; typically, moderate temperatures are used to optimise deposition and reduce defects.

• Observation: the cathode gains a shiny metal coating.
• Observation: the anode gradually dissolves, decreasing in mass.
• Variables to control: current, electrolyte concentration, temperature, time.

Electroplating is widely used to improve the appearance of objects by providing a shiny, decorative metal layer such as silver or gold plating on jewellery.

It also enhances corrosion resistance, for example, chromium plating on car parts prevents rusting.

Electroplating can improve wear resistance and reduce friction on mechanical parts, extending their service life.

• Common plating metals: silver, gold, chromium, copper, nickel.
• Industries: electronics, automotive, jewellery, aerospace.
• Benefits: aesthetic appeal, corrosion protection, increased durability.

During electrolysis, an external power supply drives the flow of electrons through the external circuit connecting two electrodes immersed in an electrolyte. Electrons flow from the negative terminal of the power supply to the cathode (negative electrode), where reduction occurs, and from the anode (positive electrode) back to the positive terminal of the power supply.

This electron flow is essential to maintain the redox reactions at the electrodes. The direction of electron flow is always from the anode to the cathode through the external circuit, opposite to the direction of conventional current.

The external circuit completes the electrical circuit, allowing continuous electron transfer and sustaining the electrolysis process.

• Electrons flow from anode to cathode in the external circuit.
• External circuit connects power supply to electrodes.
• Electron flow sustains redox reactions at electrodes.

At the electrodes, oxidation and reduction reactions occur, involving electron transfer between the electrode and ions in the electrolyte.

Oxidation occurs at the anode, where species lose electrons. These electrons enter the external circuit and flow towards the cathode.

Reduction occurs at the cathode, where species gain electrons arriving from the external circuit. This gain of electrons reduces the ions to atoms or molecules, often forming metals or gases depending on the electrolyte composition and electrode material.

• Anode: oxidation (electron loss) occurs; electrons released to external circuit.
• Cathode: reduction (electron gain) occurs; electrons accepted from external circuit.
• Electrode reactions involve electron transfer between ions and electrodes.

The electrolyte contains positive ions (cations) and negative ions (anions) that move to maintain electrical neutrality during electrolysis.

Cations migrate towards the cathode to gain electrons (reduction), while anions migrate towards the anode to lose electrons (oxidation).

This ion movement balances the charge as electrons move through the external circuit, ensuring continuous current flow and completion of the electrical circuit within the electrolyte.

• Cations move towards cathode to gain electrons.
• Anions move towards anode to lose electrons.
• Ion migration maintains charge balance in electrolyte.

Electrode reactions can be represented by half-equations showing electron transfer.

For example, in the electrolysis of molten lead(II) bromide (PbBr₂):

At the cathode (reduction): Pb²⁺(l) + 2e⁻ → Pb(l) (lead ions gain electrons to form lead metal) At the anode (oxidation): 2Br⁻(l) → Br₂(g) + 2e⁻ (bromide ions lose electrons to form bromine gas) These half-equations demonstrate electron gain and loss at electrodes and the corresponding ion movement in the electrolyte.

• Cathode half-equation shows reduction (electron gain).
• Anode half-equation shows oxidation (electron loss).
• Half-equations balance electrons to show charge transfer.

Consider the electrolysis of molten PbBr₂. The electrolyte contains Pb²⁺ and Br⁻ ions.

At the cathode, Pb²⁺ ions gain two electrons to form lead metal: Pb²⁺ + 2e⁻ → Pb

At the anode, bromide ions lose electrons to form bromine gas: 2Br⁻ → Br₂ + 2e⁻ Electrons released at the anode flow through the external circuit to the cathode, where they are used to reduce Pb²⁺ ions. Meanwhile, Pb²⁺ ions move towards the cathode, and Br⁻ ions move towards the anode to maintain charge balance in the molten electrolyte.

• Pb²⁺ ions reduced at cathode by electron gain.
• Br⁻ ions oxidised at anode by electron loss.
• Electron flow in external circuit connects oxidation and reduction.

Electrolysis is the process of using an electric current to drive a non-spontaneous chemical reaction. When copper(II) sulfate (CuSO₄) is dissolved in water, it dissociates into copper ions (Cu²⁺) and sulfate ions (SO₄²⁻). The solution also contains water molecules, which can participate in the electrolysis reactions.

During electrolysis, two electrodes are immersed in the aqueous copper(II) sulfate solution and connected to a power supply. The cathode is the negative electrode where reduction occurs, and the anode is the positive electrode where oxidation occurs.

The products formed at each electrode depend on the electrode material used: inert carbon/graphite electrodes or copper electrodes. This affects the reactions and observations during electrolysis.

• Copper(II) sulfate solution contains Cu²⁺ and SO₄²⁻ ions.
• Electrolysis involves oxidation at the anode and reduction at the cathode.
• Electrode material influences the products formed.

When inert carbon or graphite electrodes are used, they do not react chemically with the electrolyte or products. The ions in solution compete to be discharged at the electrodes.

At the cathode (negative electrode), copper ions (Cu²⁺) are reduced to copper metal because copper ions have a higher reduction potential than hydrogen ions from water. The copper metal deposits as a reddish-brown solid on the cathode.

At the anode (positive electrode), water molecules are oxidised to oxygen gas and hydrogen ions. Oxygen gas bubbles form at the anode and can be observed as effervescence. The sulfate ions (SO₄²⁻) are not discharged because they are less easily oxidised than water molecules under these conditions.

• Cathode reaction: Cu²⁺ + 2e⁻ → Cu (solid, reddish-brown deposit)
• Anode reaction: 2H₂O → O₂ (gas bubbles) + 4H⁺ + 4e⁻
• Observation: Copper deposits on cathode; oxygen bubbles at anode.

The half-equations show the electron transfer at each electrode during electrolysis with inert electrodes:

At the cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)

At the anode (oxidation): 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻

• Electrons flow from anode to cathode through the external circuit.
• Copper ions gain electrons and deposit as copper metal.
• Water molecules lose electrons to form oxygen gas and hydrogen ions.

When copper electrodes are used, the electrode material participates in the reactions.

At the cathode (negative electrode), copper ions (Cu²⁺) in solution are reduced to copper metal and deposit on the cathode, similar to the inert electrode case.

At the anode (positive electrode), copper metal atoms from the copper electrode lose electrons (are oxidised) to form copper ions (Cu²⁺) which enter the solution. This causes the anode to gradually dissolve and lose mass over time.

• Cathode reaction: Cu²⁺ + 2e⁻ → Cu (deposition)
• Anode reaction: Cu (solid) → Cu²⁺ + 2e⁻ (dissolution)
• Observation: Cathode gains copper deposit; anode gradually dissolves.

The half-equations for electrolysis with copper electrodes are:

At the cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)

At the anode (oxidation): Cu(s) → Cu²⁺(aq) + 2e⁻ (oxidation of copper electrode)

• The copper ions released at the anode replace those removed at the cathode.
• The concentration of copper ions in solution remains approximately constant.
• The mass of the cathode increases, and the mass of the anode decreases.

Using inert electrodes, copper deposits on the cathode and oxygen gas bubbles form at the anode. The anode does not change mass.

Using copper electrodes, copper deposits on the cathode and the anode dissolves, releasing copper ions into solution. The solution concentration remains stable.

In practical electroplating or purification of copper, copper electrodes are preferred to maintain copper ion concentration and avoid oxygen evolution, which can cause impurities or corrosion.

• Inert electrodes: copper deposited, oxygen gas evolved, no electrode mass change.
• Copper electrodes: copper deposited, anode dissolves, solution concentration stable.
• Electrode mass changes can be measured to confirm reactions.
• Safety: oxygen gas is flammable; ensure good ventilation during electrolysis.

Electrolysis is the process of using an electric current to drive a non-spontaneous chemical reaction. When an aqueous solution of a halide compound (such as sodium chloride, NaCl) undergoes electrolysis, ions in the solution move towards electrodes where oxidation or reduction occurs.

Halide compounds contain halide ions (Cl⁻, Br⁻, I⁻) which can be discharged at the anode during electrolysis. The products formed depend on the concentration of the solution and the relative ease of discharge of ions present.

• Halide ions: Cl⁻, Br⁻, I⁻
• Electrolysis involves oxidation at the anode and reduction at the cathode
• Products depend on ion discharge priority and solution concentration

In aqueous solution, halide compounds dissociate into metal cations (e.g., Na⁺) and halide anions (e.g., Cl⁻). Water also ionises slightly to H⁺ and OH⁻ ions. During electrolysis, all these ions compete to be discharged at the electrodes.

At the cathode (negative electrode), cations and H⁺ ions can be reduced. At the anode (positive electrode), anions and OH⁻ ions can be oxidised. The relative discharge depends on electrode potentials and ion concentrations.

• Cations: metal ions (e.g., Na⁺), H⁺ from water
• Anions: halide ions (Cl⁻, Br⁻, I⁻), OH⁻ from water
• Competition between ions for discharge at electrodes

At the cathode, positive ions gain electrons (reduction). In dilute aqueous halide solutions, hydrogen ions (H⁺) from water are preferentially reduced over metal ions because hydrogen ions have a higher discharge priority.

The half-equation for hydrogen ion reduction is: 2H⁺(aq) + 2e⁻ → H₂(g)

In concentrated solutions containing metal ions of low reactivity (e.g., Cu²⁺), metal ions may be reduced instead of hydrogen ions, producing the metal at the cathode. However, for alkali metals like sodium, hydrogen is always produced because the metal ions are too reactive to be discharged from aqueous solution.

• Dilute solution: hydrogen gas produced at cathode
• Concentrated solution with less reactive metal ions: metal deposited
• Reduction half-equation for hydrogen: 2H⁺ + 2e⁻ → H₂

At the anode, negative ions lose electrons (oxidation). Halide ions (Cl⁻, Br⁻, I⁻) are oxidised to their respective halogen gases (Cl₂, Br₂, I₂).

The half-equations for halide oxidation are: 2Cl⁻ → Cl₂ + 2e⁻, 2Br⁻ → Br₂ + 2e⁻, 2I⁻ → I₂ + 2e⁻

If halide ions are present, they are preferentially discharged over hydroxide ions (OH⁻) because halide ions are easier to oxidise. Therefore, halogen gas is produced at the anode in both dilute and concentrated solutions of halide compounds.

• Halide ions oxidised to halogen gas at anode
• Halogen gases: Cl₂ (greenish-yellow), Br₂ (brown), I₂ (violet vapour)
• Oxidation half-equations: 2X⁻ → X₂ + 2e⁻ (X = Cl, Br, I)

The products formed during electrolysis of aqueous halide solutions can be summarised as follows:

At the cathode: hydrogen gas (H₂) is produced in dilute solutions; in concentrated solutions, metal deposition may occur if the metal ion is less reactive than hydrogen.

At the anode: halogen gas (Cl₂, Br₂, or I₂) is produced regardless of concentration, with characteristic colours and smells.

• Cathode: H₂ gas in dilute solutions; metal in concentrated (if less reactive)
• Anode: halogen gas (Cl₂, Br₂, I₂) always produced
• Observations: bubbles of gas at electrodes; halogen colour and smell at anode

In concentrated aqueous sodium bromide (NaBr), the ions present are Na⁺, Br⁻, H⁺, and OH⁻.

At the cathode, hydrogen ions are reduced because sodium ions are too reactive to discharge: 2H⁺ + 2e⁻ → H₂(g)

At the anode, bromide ions are oxidised to bromine gas: 2Br⁻ → Br₂(g) + 2e⁻ (brown vapour observed)

• Cathode reaction: 2H⁺ + 2e⁻ → H₂(g) (bubbles of hydrogen)
• Anode reaction: 2Br⁻ → Br₂(g) + 2e⁻ (brown bromine vapour)
• Overall: electrolysis produces hydrogen gas and bromine gas

Electrolysis involves passing an electric current through an ionic compound that is molten or dissolved in water, causing chemical changes at the electrodes. The electrodes are the anode (positive electrode) and the cathode (negative electrode). At the anode, oxidation occurs, meaning species lose electrons. At the cathode, reduction occurs, meaning species gain electrons.

Oxidation and reduction always occur simultaneously; electrons lost at the anode travel through the external circuit to the cathode, where they are gained. Ionic half-equations represent these processes separately, showing the species involved, electrons lost or gained, and the balanced charges and atoms.

• Anode: site of oxidation (electron loss).
• Cathode: site of reduction (electron gain).
• Electrons flow from anode to cathode through the external circuit.

To construct ionic half-equations, first identify the species undergoing oxidation or reduction at each electrode. Then, write the chemical species on the reactant side and the products formed on the product side. Balance all atoms except hydrogen and oxygen first, then balance oxygen atoms by adding H₂O molecules, and hydrogen atoms by adding H⁺ ions (in acidic solution). Finally, balance the charges by adding electrons (e⁻).

At the anode, electrons appear on the product side (since they are lost). At the cathode, electrons appear on the reactant side (since they are gained).

The half-equations must be balanced in terms of both atoms and charge.

• Identify species oxidised at anode and reduced at cathode.
• Balance atoms other than H and O first.
• Balance oxygen with H₂O, hydrogen with H⁺ (acidic conditions).
• Balance charge with electrons (e⁻).
• Electrons on product side for oxidation; on reactant side for reduction.

1. Electrolysis of molten lead(II) bromide (PbBr₂):

At the cathode (reduction): Pb²⁺ ions gain electrons to form lead atoms.

Half-equation: Pb²⁺(l) + 2e⁻ → Pb(l) (reduction) At the anode (oxidation): bromide ions lose electrons to form bromine molecules, which are diatomic gases released as brown vapour. Half-equation: 2Br⁻(l) → Br₂(g) + 2e⁻ (oxidation) 2. Electrolysis of aqueous copper(II) sulfate with inert electrodes: At the cathode: Cu²⁺ ions gain electrons to form copper metal. Half-equation: Cu²⁺(aq) + 2e⁻ → Cu(s) At the anode: water molecules are oxidised to oxygen gas and H⁺ ions, as sulfate ions are not discharged. Half-equation: 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻ These examples illustrate the principle that oxidation involves electron loss at the anode and reduction involves electron gain at the cathode, with ionic species changing oxidation states accordingly.

• Pb²⁺ + 2e⁻ → Pb (cathode reduction)
• 2Br⁻ → Br₂ + 2e⁻ (anode oxidation)
• Cu²⁺ + 2e⁻ → Cu (cathode reduction)
• 2H₂O → O₂ + 4H⁺ + 4e⁻ (anode oxidation)

During electrolysis of molten sodium chloride (NaCl), the compound dissociates into Na⁺ and Cl⁻ ions. At the cathode, sodium ions gain electrons to form sodium metal. At the anode, chloride ions lose electrons to form chlorine gas.

Step 1: Identify species at cathode and anode: Cathode: Na⁺ ions reduced to Na(s) Anode: Cl⁻ ions oxidised to Cl₂(g) Step 2: Write half-equations: Cathode (reduction): Na⁺ + e⁻ → Na Anode (oxidation): 2Cl⁻ → Cl₂ + 2e⁻ Step 3: Check atom and charge balance: Cathode: 1 Na atom and charge balanced (1+ on left, neutral on right with 1 e⁻ gained). Anode: 2 Cl atoms on left and right, charges balanced by 2 electrons lost.

This example demonstrates the construction of ionic half-equations by identifying species, electron transfer, and balancing atoms and charges.

• Na⁺ + e⁻ → Na (cathode reduction)
• 2Cl⁻ → Cl₂ + 2e⁻ (anode oxidation)

Do not confuse oxidation and reduction: oxidation is loss of electrons (anode), reduction is gain of electrons (cathode).

Always include electrons explicitly in half-equations to balance charge.

Balance atoms carefully; do not omit water or H⁺ ions if balancing in aqueous acidic conditions (though for electrolysis of molten salts, water is absent). Use correct state symbols where relevant (e.g., (s), (l), (g), (aq)) to clarify phases and exam clarity. Remember that electrons are on the product side for oxidation and on the reactant side for reduction half-equations. Practice writing half-equations for common electrolysis reactions to build confidence and accuracy.

• Oxidation = electron loss at anode; reduction = electron gain at cathode.
• Always balance charge with electrons.
• Include state symbols for clarity.
• Electrons on product side for oxidation; reactant side for reduction.
• Practice common examples for exam readiness.