Chemical Cells and Electrolysis
AdvancedA chemical cell splits a spontaneous redox reaction onto two electrodes, converting chemical energy to electrical energy as electrons flow from the negative to the positive electrode through the external circuit. Electrolysis uses electrical energy to drive non-spontaneous reactions: reduction at the cathode, oxidation at the anode, with products set by the preferential discharge of ions.
- Chemical cell: a spontaneous redox reaction split onto two electrodes, chemical energy to electrical energy, electrons flowing from the negative electrode through the external circuit to the positive.
- The salt bridge lets ions migrate, keeps neutrality and completes the internal circuit; electrons never use it.
- Electrolysis: electrical energy to chemical energy; reduction at the cathode (negative terminal), oxidation at the anode (positive terminal).
文體小知識：科學知識說明文——解釋化學電池如何把化學能變成電能、電解如何用電流驅動非自發反應，並比較兩極反應與產物的判斷方法。
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1Insert a zinc strip and a copper strip into a lemon and connect them to a small bulb — the bulb lights up. This is the fun of a fruit cell. The principle of a chemical cell is to split a spontaneous redox reaction onto two separate electrodes: oxidation, the loss of electrons, happens at one electrode, and reduction, the gain of electrons, at the other. Electrons flow through the external circuit as a current, and chemical energy is converted into electrical energy.
2Electrons travel only through the external circuit, never across the solution. Inside the cell, ions carry the current. A simple chemical cell has two metals of different reactivity dipped in an electrolyte: the more reactive metal acts as the negative electrode, losing electrons and being oxidised, while the less reactive metal is the positive electrode, where species in solution gain electrons and are reduced. Reactivity is compared using the electrochemical series — the higher a metal stands, the more readily it loses electrons.
3The more precise metal–metal ion half-cell arrangement joins two half-cells with a salt bridge or porous pot. The salt bridge serves two purposes: it lets ions migrate between the half-cells to keep the solutions electrically neutral, and it completes the internal circuit so the current can continue. Without it, charge builds up rapidly, the reaction stops, and the current ceases.
4Electrolysis is the reverse process: electrical energy drives redox reactions that cannot occur spontaneously, converting electrical energy into chemical energy. An electrolytic cell consists of a power supply, electrodes and an electrolyte. The electrode joined to the negative terminal is the cathode, where reduction occurs; the one joined to the positive terminal is the anode, where oxidation occurs. When molten lead(II) bromide is electrolysed, lead ions gain electrons at the cathode to form molten lead, while bromide ions lose electrons at the anode to give reddish-brown bromine vapour.
5When aqueous solutions are electrolysed, several kinds of ions may compete to discharge, and the products depend on the preferential discharge of ions, governed by the electrochemical series and ion concentrations. Applications are everywhere in industry: electroplating coats objects with a protective, attractive metal layer, and extracting reactive metals such as aluminium relies on electrolysis of molten ore; the dry cell, the lead-acid accumulator and the hydrogen-oxygen fuel cell are chemical cells living in daily life.
1Summary 重點摘要
A chemical cell splits a spontaneous redox reaction onto two electrodes, converting chemical energy to electrical energy as electrons flow from the negative to the positive electrode through the external circuit. Electrolysis uses electrical energy to drive non-spontaneous reactions: reduction at the cathode, oxidation at the anode, with products set by the preferential discharge of ions.
2Vocabulary cards Vocabulary
A device converting the chemical energy of a spontaneous redox reaction into electrical energy.
Using electrical energy to drive non-spontaneous redox reactions, converting electrical energy into chemical energy.
A conductor in contact with an electrolyte or solution, at which oxidation or reduction occurs.
In an electrolytic cell, the electrode joined to the negative terminal, where reduction occurs.
In an electrolytic cell, the electrode joined to the positive terminal, where oxidation occurs.
A substance that conducts electricity when molten or dissolved in water, by movement of ions.
A bridge joining two half-cells, allowing ions to migrate so neutrality is kept and the internal circuit is completed.
An electrode with its electrolyte, carrying out just the oxidation or the reduction half of the reaction.
The order of metals (and their ions) by reactivity or tendency to lose or gain electrons, used to compare electrode reactions.
Using electrolysis to coat an object with a layer of metal; the object is made the cathode.
3Step by step Steps
- Step 1: A chemical cell splits the reaction in twoThe central idea of a chemical cell is to run the two halves of a spontaneous redox reaction separately: oxidation loses electrons at the negative electrode, reduction gains them at the positive electrode, and the electrons flowing through the external circuit form the current. For example, dip a zinc strip and a copper strip into lemon juice, an electrolyte, and a bulb lights up — the zinc-losing-electrons half and the solution's gaining-electrons half proceed on different metals. This design converts chemical energy directly into electrical energy instead of letting it escape as heat — the scientific basis of the battery industry.
- Step 2: Electron flow in the zinc-copper cellIn a zinc-copper cell, first compare reactivity with the electrochemical series: zinc stands above copper and loses electrons more easily, so the zinc strip is the negative electrode and the copper strip the positive. The negative electrode reaction is Zn → Zn²⁺ + 2e⁻ (oxidation); at the positive electrode, species in solution gain electrons and are reduced. For example, lemon juice contains hydrogen ions, which gain electrons on the copper surface and give off hydrogen gas. Electrons flow from zinc through the wire to copper, and the current flows the opposite way. Writing 'electrons flow from the negative electrode through the external circuit to the positive' plus both electrode reactions is the standard way to score.
- Step 3: The roles of the salt bridge and half-cellsA metal–metal ion half-cell dips a metal into a solution of its own ions; two half-cells are joined by a salt bridge or a porous pot. The salt bridge lets ions migrate: when one side becomes charged because cations are consumed or anions accumulate, ions from the bridge move in to restore neutrality. For example, remove the salt bridge and the bulb soon goes out, because charge accumulates and the reaction halts. The bridge is not for electrons — electrons only use the external circuit; ions use the internal one, and only with both routes working does the current continue.
- Step 4: Electrolysis runs the reaction backwardsElectrolysis uses electrical energy to drive reactions that are not spontaneous, converting electrical energy into chemical energy. In an electrolytic cell, the electrode joined to the negative terminal is the cathode, where reduction happens; the one joined to the positive terminal is the anode, where oxidation happens. Mnemonic: the cathode gains electrons, the anode loses them. For example, dip two carbon rods into molten lead(II) bromide and connect a d.c. supply: lead ions move to the cathode and gain electrons to become lead metal, while bromide ions move to the anode, lose electrons and give bromine vapour. To name an electrode, first check which supply terminal it joins, then check whether the reaction gains or loses electrons — cross-checking the two prevents errors.
- Step 5: Electrolysing molten compoundsA molten ionic compound has just one type of cation and one type of anion, so the products are easiest to predict. Take molten PbBr₂: at the cathode Pb²⁺ + 2e⁻ → Pb, giving silvery-grey molten lead; at the anode 2Br⁻ → Br₂ + 2e⁻, releasing reddish-brown bromine vapour. For example, industrially extracting the reactive metal aluminium works exactly this way — molten aluminium oxide is electrolysed, and aluminium ions become metal at the cathode; aluminium is far too reactive for carbon reduction, so only electrolysis can win it from its ore. Full-mark answers state both half equations and the observations.
- Step 6: Aqueous solutions and preferential dischargeIn aqueous solutions the water itself contributes hydrogen ions and hydroxide ions, which may compete with the solute ions to discharge; the products follow the preferential discharge of ions. General rules: at the cathode, ions of less reactive metals discharge first; if only ions of very reactive metals such as potassium, sodium, calcium, magnesium and aluminium are present, hydrogen ions discharge and hydrogen gas forms. At the anode, halide ions discharge before hydroxide ions, giving halogens; if only nitrate or sulfate ions are present, hydroxide ions discharge and oxygen forms. For example, electrolysis of dilute copper(II) sulfate gives copper at the cathode and oxygen at the anode, while the blue colour fades as copper ions are used up.
- Step 7: Electroplating and the cells around youElectroplating applies electrolysis directly: the object to be plated is the cathode, the plating metal the anode, and the electrolyte contains ions of the plating metal. For example, in copper plating the object is joined to the cathode and a copper sheet to the anode; at the cathode Cu²⁺ + 2e⁻ → Cu deposits a smooth copper layer, while the copper anode dissolves to replenish the ions — the coating is decorative and rust-resistant. Everyday cells are also this lesson's heroes: the dry cell for portability, the lead-acid accumulator in cars, and the hydrogen-oxygen fuel cell, which turns the chemical energy of hydrogen and oxygen directly into electricity with only water as the product — an important direction for clean energy.
4Diagrams Diagrams
5Worked example Example
A science-fair project makes a fruit cell by dipping zinc and copper strips into a lemon, wired to a meter; separately, carbon rods electrolyse molten PbBr₂. State the direction of electron flow and the electrode reactions of the cell, name the products at the two electrodes of the electrolysis, and say which terminal an iron spoon should join to be copper-plated.
- Compare the electrochemical series: zinc stands above copper and loses electrons more readily, so zinc is the negative electrode (Zn → Zn²⁺ + 2e⁻) and copper the positive; electrons flow from zinc through the wire to copper.
- Hydrogen ions from the lemon juice are reduced on the copper surface, giving off hydrogen gas; the internal circuit is carried by moving ions.
- For molten PbBr₂: cathode Pb²⁺ + 2e⁻ → Pb, giving silvery-grey molten lead; anode 2Br⁻ → Br₂ + 2e⁻, giving reddish-brown bromine vapour.
- To copper-plate the spoon, it must be the cathode so that Cu²⁺ gains electrons and deposits copper on its surface; a copper sheet at the anode keeps the ions topped up.
- Answer: In the zinc-copper cell, zinc is the negative electrode and electrons flow from zinc to copper through the external circuit, where hydrogen ions are reduced. Electrolysing molten PbBr₂ gives molten lead at the cathode and bromine vapour at the anode. For copper plating, the spoon joins the cathode and the copper sheet the anode.
6Common mistakes Mistakes
7Quick check Quick check
Q1What energy conversions take place in a chemical cell and in an electrolytic cell?
Q2In a zinc-copper cell, towards which metal do electrons flow through the external circuit, and which metal is the negative electrode?
Q3What are the two functions of a salt bridge in a chemical cell?
Q4What products form at the cathode and anode when molten PbBr₂ is electrolysed?
Q5To copper-plate an iron spoon, which terminal should the spoon join, and what ions must the electrolyte contain?
8Recap Recap
- Chemical cell: a spontaneous redox reaction split onto two electrodes, chemical energy to electrical energy, electrons flowing from the negative electrode through the external circuit to the positive.
- The salt bridge lets ions migrate, keeps neutrality and completes the internal circuit; electrons never use it.
- Electrolysis: electrical energy to chemical energy; reduction at the cathode (negative terminal), oxidation at the anode (positive terminal).
- In aqueous electrolysis the products follow preferential discharge of ions; reactive metals must be extracted by electrolysing molten ores.
- In electroplating the object is the cathode, the plating metal the anode, and the electrolyte carries ions of the plating metal.