Hey Tara! Welcome to Topic 4 - Electrochemistry. This is where chemistry meets electricity. You will learn how we use electrical energy to break apart compounds (electrolysis), how we use it industrially to extract aluminium and purify copper, and how chemical reactions can actually generate electricity in fuel cells and batteries. Electrochemistry is everywhere - from the phone battery you are charging right now to the aluminium foil in your kitchen. Let us dive in!
4.1 Electrolysis
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What is Electrolysis?
Electrolysis is the decomposition (breaking down) of an ionic compound, when molten or dissolved in water, using an electric current. The word itself gives you a clue: electro = electricity, lysis = splitting apart.
The substance being decomposed is called the electrolyte. For electrolysis to work, the electrolyte must contain free-moving ions - this is why it must be either molten (melted) or dissolved in water (aqueous). In the solid state, the ions are locked in a fixed lattice and cannot move, so electrolysis cannot happen.
What Do You Need for Electrolysis?
Three things are required for electrolysis to work:
- An electrolyte - a molten or dissolved ionic compound that contains free-moving ions. Examples: molten lead(II) bromide, copper sulfate solution, dilute sulfuric acid.
- Two electrodes - these are rods or plates (usually made of an inert material like carbon/graphite or platinum) that are dipped into the electrolyte. They carry the electric current into and out of the liquid.
- The cathode is the negative electrode (connected to the negative terminal of the battery).
- The anode is the positive electrode (connected to the positive terminal of the battery).
- A direct current (d.c.) supply - a battery or power pack that provides a one-way flow of electrons. We use d.c. (not a.c.) because we need a fixed positive and negative electrode.
Memory Trick
How to remember which electrode is which? Think PANIC: Positive = ANode, In Cathode. Or remember: aNode = Negative? NO! The anode is Positive (A-P). The cathode attracts Cations (positive ions) because it is negative.
Movement of Ions During Electrolysis
When the d.c. supply is switched on, something amazing happens - the ions in the electrolyte start to move:
- Cations (positive ions, e.g. Na⁺, Cu²⁺, Pb²⁺, H⁺) move towards the cathode (negative electrode). Opposite charges attract!
- Anions (negative ions, e.g. Cl⁻, Br⁻, O²⁻, OH⁻, SO₄²⁻) move towards the anode (positive electrode).
This movement of ions is what carries the electric current through the liquid. In the external wires, the current is carried by electrons. Inside the electrolyte, it is carried by the moving ions.
Battery / d.c. supply
(-) ___________ (+)
| | | |
| |___________| |
| |
wire | | wire
(e- flow down) (e- flow up)
| |
______|____ ______|____
| CATHODE | | ANODE |
| (-) | | (+) |
|___________| |___________|
\ /
\ /
--------\-----------/--------
| Electrolyte (molten or |
| dissolved ionic compound) |
| |
| Cations ------> Cathode |
| (+ ions) (-) |
| |
| Anions -------> Anode |
| (- ions) (+) |
|____________________________|
Diagram: Electrolysis setup showing ion movement. Cations move to the cathode, anions move to the anode.
Products of Electrolysis
When the ions reach the electrodes, they undergo chemical reactions:
- At the cathode (-): metals are deposited, or hydrogen gas is produced. The cations gain electrons (they are reduced).
- At the anode (+): non-metals are produced (e.g. oxygen, chlorine, bromine). The anions lose electrons (they are oxidised).
Exam Tip
A very common exam question asks: "State what happens at the cathode during electrolysis." The answer must always mention that cations (positive ions) move to the cathode and gain electrons to form atoms. At the anode, anions (negative ions) lose electrons.
Example: Electrolysis of Molten Lead(II) Bromide
This is the key example you must know for your exam. Let us walk through it carefully.
Lead(II) bromide (PbBr₂) is an ionic compound. In the solid state, its ions are fixed in a lattice and cannot move, so solid PbBr₂ does not conduct electricity. But when we melt it (it melts at 373°C), the ions become free to move.
The ions present in molten PbBr₂ are:
- Pb²⁺ ions (lead cations - positive)
- Br⁻ ions (bromide anions - negative)
When d.c. is passed through molten PbBr₂:
- At the cathode (-): Pb²⁺ ions move to the cathode, gain 2 electrons, and are deposited as silvery-grey molten lead metal at the bottom of the container.
- At the anode (+): Br⁻ ions move to the anode, lose electrons, and are released as brown/orange bromine vapour (toxic fumes).
PbBr₂(l) → Pb(l) + Br₂(g)
Overall equation: molten lead(II) bromide decomposes into lead and bromine
Supplement
Half-Equations at the Electrodes
Half-equations show exactly what happens at each electrode. They show the electrons being gained or lost.
At the cathode (-):
Pb²⁺ + 2e⁻ → Pb
Lead ions gain 2 electrons to form lead atoms. This is reduction (gain of electrons).
At the anode (+):
2Br⁻ → Br₂ + 2e⁻
Bromide ions lose electrons to form bromine molecules. This is oxidation (loss of electrons).
Memory Trick - OIL RIG
Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons). At the anode = oxidation (both start with vowels!). At the cathode = reduction (both start with consonants!).
Electrolysis as Oxidation and Reduction
Electrolysis always involves both oxidation and reduction happening simultaneously:
- Reduction at the cathode: Cations gain electrons. Gaining electrons = reduction. Example: Cu²⁺ + 2e⁻ → Cu
- Oxidation at the anode: Anions lose electrons. Losing electrons = oxidation. Example: 2Cl⁻ → Cl₂ + 2e⁻
This makes electrolysis a redox reaction - reduction and oxidation happen at the same time, but at different electrodes.
Relating Products to the Electrolyte and Electrodes
The products of electrolysis depend on:
- The electrolyte used - which ions are present in the solution or melt? The products come from these ions.
- The electrodes used - if the electrodes are inert (unreactive, e.g. carbon or platinum), they do not take part in the reaction. But if a reactive electrode is used (e.g. copper anode), the electrode itself can dissolve and become part of the reaction.
For molten electrolytes, it is straightforward - the only ions present are from the compound itself. For example, molten NaCl contains only Na⁺ and Cl⁻, so you get sodium at the cathode and chlorine at the anode.
For aqueous solutions, it is more complex because water (H₂O) adds H⁺ and OH⁻ ions to the mix. This means there are competing ions at each electrode, and we need rules to predict which ion is discharged. We will cover this in Section 4.2.
Worked Examples
Worked Example 1
Describe what happens during the electrolysis of molten lead(II) bromide using carbon electrodes.
Show Solution
Step 1: Identify the ions
PbBr₂ contains Pb²⁺ ions and Br⁻ ions. When molten, these ions are free to move.
Step 2: Cathode (negative electrode)
Pb²⁺ ions (cations) are attracted to the cathode. They gain 2 electrons and are reduced: Pb²⁺ + 2e⁻ → Pb. A silvery-grey layer of lead metal forms.
Step 3: Anode (positive electrode)
Br⁻ ions (anions) are attracted to the anode. They lose electrons and are oxidised: 2Br⁻ → Br₂ + 2e⁻. Brown/orange bromine vapour is observed.
At the cathode: silvery-grey lead metal is deposited. At the anode: brown/orange bromine gas is released.
Worked Example 2
Write the half-equations for the electrolysis of molten sodium chloride (NaCl). State the product at each electrode.
Show Solution
Step 1: Identify the ions
NaCl contains Na⁺ and Cl⁻ ions. When molten, both are free to move.
Step 2: Cathode half-equation
Na⁺ + e⁻ → Na (sodium metal is deposited - a silvery reactive metal). This is reduction.
Step 3: Anode half-equation
2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas is produced - a pale green/yellow gas with a pungent smell). This is oxidation.
Cathode: Na⁺ + e⁻ → Na (sodium metal). Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas).
Worked Example 3
Explain why solid lead(II) bromide does not conduct electricity, but molten lead(II) bromide does.
Show Solution
Step 1: Solid state
In solid PbBr₂, the ions (Pb²⁺ and Br⁻) are held in a fixed, regular lattice structure by strong electrostatic forces. They cannot move from their positions, so they cannot carry an electric current.
Step 2: Molten state
When PbBr₂ is heated above its melting point (373°C), the ionic lattice breaks down. The ions become free to move throughout the liquid. These free-moving ions can carry the electric current to the electrodes.
Step 3: Key point
Electrolysis requires free-moving ions. Only molten or dissolved ionic compounds have free-moving ions. In the solid state, ions are fixed in position and cannot move.
Solid PbBr₂ has ions fixed in a lattice - they cannot move. Molten PbBr₂ has free-moving ions that can carry the electric current to the electrodes.
Exam Tips for 4.1
1. Always state that ions must be free to move for electrolysis to work. Do not just say "ions are present" - solid ionic compounds have ions too, but they cannot conduct because the ions are fixed.
2. When writing half-equations, make sure the charges balance and the atoms balance. For example: 2Br⁻ → Br₂ + 2e⁻ has 2 bromines on each side and charge -2 on each side.
3. Remember that d.c. (direct current) is needed, not a.c. With a.c., the electrodes keep swapping between positive and negative, so you cannot get a consistent product at each electrode.
4. If asked "why is lead(II) bromide used molten rather than dissolved in water?", the answer is: if dissolved in water, H⁺ and OH⁻ ions from water would compete with Pb²⁺ and Br⁻ ions, changing the products. Molten PbBr₂ gives pure lead and bromine.
Apply It: Real-World Chemistry
Electrolysis is not just a textbook concept - it powers some of the world's most important industries. See how the science you just learned shows up in real life.
1
A student at Bangalore International School sets up an electrolysis experiment using solid lead(II) bromide connected to a battery and bulb circuit. The bulb does not light up. She then heats the lead(II) bromide with a Bunsen burner until it melts, and the bulb lights up.
Explain why the bulb only lights up when the lead(II) bromide is molten.
Explanation
In solid lead(II) bromide, the Pb²⁺ and Br⁻ ions are held in a fixed lattice. They cannot move, so no current flows and the bulb stays off. When the compound melts, the lattice breaks down and the ions become free to move. The Pb²⁺ ions move towards the cathode and Br⁻ ions move towards the anode, carrying the current. This completes the circuit and the bulb lights up.
Key Takeaway
Electrolysis requires free-moving ions. Ionic solids have ions but they are locked in place. Melting or dissolving sets them free.
2
The Dow Chemical Company in Freeport, Texas, historically extracted bromine from seawater using electrolysis. They would concentrate the bromide ions from the sea, then pass a current through the solution.
What product forms at the anode, and what type of reaction (oxidation or reduction) occurs there?
Explanation
At the anode (positive electrode), the Br⁻ ions are attracted and lose electrons: 2Br⁻ → Br₂ + 2e⁻. Bromine (Br₂) is the product - a reddish-brown, pungent liquid. This is oxidation because the bromide ions are losing electrons.
Key Takeaway
Oxidation always happens at the anode. Anions lose electrons to become atoms or molecules. OIL RIG: Oxidation Is Loss.
3
A chemistry teacher in London demonstrates electrolysis using molten zinc chloride (ZnCl₂) with graphite electrodes. She asks her class to predict what they will observe at each electrode.
What forms at the cathode and anode? Write the half-equations.
At the cathode
Zn²⁺ ions gain electrons: Zn²⁺ + 2e⁻ → Zn. Zinc metal is deposited as a grey metallic coating on the electrode.
At the anode
Cl⁻ ions lose electrons: 2Cl⁻ → Cl₂ + 2e⁻. Chlorine gas is produced - a pale green/yellow gas with a pungent, bleach-like smell.
Key Takeaway
For any molten binary ionic compound, the metal always forms at the cathode and the non-metal at the anode. The half-equations must balance for both atoms and charge.
4
In a metals recycling plant in Jamshedpur, India, engineers need to extract pure lead from lead(II) bromide waste. They decide to use electrolysis rather than chemical reduction.
Why must they use direct current (d.c.) rather than alternating current (a.c.) for this process?
Explanation
Direct current flows in one direction only, which means the cathode always stays negative and the anode always stays positive. This ensures Pb²⁺ ions consistently move to the cathode to form lead metal. With alternating current, the polarity of the electrodes would keep switching (50 times per second in India's 50 Hz supply), so ions would keep changing direction. No consistent product would form at either electrode.
Key Takeaway
D.C. provides a fixed positive and negative electrode. A.C. would reverse the ion flow 50-60 times per second, preventing any useful product from building up.
5
A student connects a battery to two graphite rods dipped in pure distilled water. She expects electrolysis to occur, but observes no bubbles and no current flow on the ammeter.
Why does pure water not undergo electrolysis? What could she add to make it work?
Explanation
Pure water is a very poor conductor of electricity because it contains extremely few ions. Water molecules (H₂O) are covalent and mostly un-ionised. For electrolysis to work, the electrolyte must contain free-moving ions in sufficient concentration. She could add a small amount of dilute sulfuric acid (H₂SO₄) or sodium hydroxide (NaOH) to provide ions, making the solution conductive without changing the products (which would be hydrogen at the cathode and oxygen at the anode).
Key Takeaway
Electrolysis needs free-moving ions. Pure covalent substances like water have almost no ions and cannot conduct. Adding an acid, alkali, or dissolved salt provides the ions needed.
Practice Questions: 4.1 Electrolysis
Test your understanding - 20 questions
Your Score
0 / 20
QUESTION 1 OF 20
What is the definition of electrolysis?
Electrolysis is the decomposition (breaking down) of an ionic compound, when molten or dissolved in water, using an electric current. 'Electro' means electricity, 'lysis' means splitting.
QUESTION 2 OF 20
Which of the following is NOT required for electrolysis?
Electrolysis requires a d.c. (direct current) supply, NOT an a.c. supply. D.C. provides a fixed positive and negative electrode. A.C. would reverse the polarity constantly, preventing consistent products from forming.
QUESTION 3 OF 20
During electrolysis, cations move towards the:
Cations are positive ions. They are attracted to the cathode (negative electrode) because opposite charges attract. Remember: CATions go to the CAThode.
QUESTION 4 OF 20
Why does solid lead(II) bromide not conduct electricity?
Solid PbBr₂ does contain ions (Pb²⁺ and Br⁻), but they are held in a fixed lattice by strong electrostatic forces. The ions cannot move, so they cannot carry the current. When melted, the lattice breaks and ions become free to move.
QUESTION 5 OF 20
What is observed at the anode during electrolysis of molten lead(II) bromide?
At the anode, Br⁻ ions lose electrons and form bromine (Br₂), which appears as brown/orange fumes. Lead metal forms at the cathode, not the anode.
QUESTION 6 OF 20
The cathode is connected to which terminal of the battery?
The cathode is the negative electrode, connected to the negative terminal of the battery. The anode is the positive electrode, connected to the positive terminal.
QUESTION 7 OF 20
What is the correct half-equation for the reaction at the cathode during electrolysis of molten PbBr₂?
At the cathode, lead ions gain electrons (reduction): Pb²⁺ + 2e⁻ → Pb. Option A shows the reverse (oxidation). Option C is the anode half-equation.
QUESTION 8 OF 20
Oxidation occurs at which electrode during electrolysis?
Oxidation (loss of electrons) occurs at the anode. Reduction (gain of electrons) occurs at the cathode. Remember: AN OX (anode = oxidation) and RED CAT (reduction = cathode).
QUESTION 9 OF 20
What type of substance always forms at the cathode during electrolysis?
Metals or hydrogen always form at the cathode. This is because metal ions and hydrogen ions are cations (positive), so they are attracted to the negative cathode where they gain electrons.
QUESTION 10 OF 20
An electrolyte is best described as:
An electrolyte is specifically a molten or dissolved ionic compound. It conducts because it contains free-moving ions. Mercury (liquid metal) conducts but is not an electrolyte. Sugar solution does not conduct because sugar is covalent.
QUESTION 11 OF 20
During electrolysis of molten NaCl, what is the product at the cathode?
Na⁺ ions are cations that move to the cathode where they gain an electron: Na⁺ + e⁻ → Na. Sodium metal is deposited. Chlorine forms at the anode from Cl⁻ ions.
QUESTION 12 OF 20
Which of the following substances can act as an electrolyte?
Copper sulfate solution contains free-moving Cu²⁺ and SO₄²⁻ ions (plus H⁺ and OH⁻ from water). Solid NaCl has ions but they cannot move. Pure water has very few ions. Hexane is a covalent liquid with no ions at all.
QUESTION 13 OF 20
What carries the electric current through the electrolyte?
In the electrolyte, current is carried by moving ions (cations move one way, anions the other). In the external wires, current is carried by electrons. Do not confuse these - ions carry the current inside the liquid, electrons carry it in the wires.
QUESTION 14 OF 20
The half-equation 2Br⁻ → Br₂ + 2e⁻ represents:
The bromide ions are losing electrons (shown by e⁻ on the product side). Loss of electrons = oxidation. Oxidation occurs at the anode. OIL RIG!
QUESTION 15 OF 20
Inert electrodes are used in electrolysis because they:
Inert electrodes (like graphite or platinum) do not react with the electrolyte or the products formed. This means the products observed come only from the electrolyte, not from the electrodes dissolving.
QUESTION 16 OF 20
What is the overall equation for the electrolysis of molten lead(II) bromide?
Electrolysis decomposes PbBr₂ into its elements: lead metal (Pb) and bromine gas (Br₂). Note that bromine forms as Br₂ molecules, not individual Br atoms.
QUESTION 17 OF 20
Why would using a.c. instead of d.c. in electrolysis not produce useful products?
A.C. reverses direction many times per second. This means the electrodes swap between positive and negative, and the ions keep changing direction. No consistent product accumulates at either electrode.
QUESTION 18 OF 20
When molten magnesium chloride (MgCl₂) is electrolysed, what forms at the cathode?
Mg²⁺ cations are attracted to the cathode and gain 2 electrons: Mg²⁺ + 2e⁻ → Mg. Magnesium metal is deposited. In a molten binary compound, the metal always forms at the cathode.
QUESTION 19 OF 20
In the electrolysis of molten PbBr₂, how many electrons does each Pb²⁺ ion gain at the cathode?
Pb²⁺ has a 2+ charge, so it needs to gain 2 electrons to become a neutral Pb atom: Pb²⁺ + 2e⁻ → Pb. The number of electrons gained always equals the charge on the ion.
QUESTION 20 OF 20
Which statement about electrolysis is correct?
Reduction (gain of electrons) occurs at the cathode; oxidation (loss of electrons) at the anode. Anions go to the anode (not cathode). Ions, not electrons, carry the current in the electrolyte. The electrolyte can be molten or aqueous.
4.2 Electrolysis of Aqueous Solutions
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Why Are Aqueous Solutions More Complex?
When an ionic compound is dissolved in water instead of being molten, something important changes: water itself produces ions. Water (H₂O) partially ionises into:
- H⁺ ions (hydrogen ions)
- OH⁻ ions (hydroxide ions)
This means that in an aqueous solution of, say, sodium chloride (NaCl), there are four types of ions present:
- Na⁺ and H⁺ competing at the cathode (both are cations)
- Cl⁻ and OH⁻ competing at the anode (both are anions)
So we need rules to predict which ion gets discharged at each electrode.
Rules for Predicting Products at the Cathode
At the cathode (negative electrode), we need to decide: is the metal ion or the hydrogen ion discharged?
The rule is based on the reactivity series:
- If the metal is more reactive than hydrogen (i.e. above hydrogen in the reactivity series - K, Na, Ca, Mg, Al, Zn), then hydrogen gas is produced. The metal ions stay in solution because they hold onto their electrons more tightly - they are harder to reduce.
- If the metal is less reactive than hydrogen (i.e. below hydrogen in the reactivity series - Cu, Ag, Au, Pt), then the metal is deposited. The metal ions are easier to reduce than H⁺ ions.
Memory Trick - Reactivity and the Cathode
Think of it this way: reactive metals are "clingy" - they do not want to become atoms (they prefer being ions). So if the metal is very reactive, it stays as an ion and hydrogen is discharged instead. Less reactive metals are "happy to be atoms" so they readily accept electrons and are deposited.
Rules for Predicting Products at the Anode
At the anode (positive electrode), we need to decide: is the anion from the compound or the hydroxide ion (OH⁻) discharged?
- If a halide ion (Cl⁻, Br⁻, or I⁻) is present in concentrated solution, then the halogen is produced (Cl₂, Br₂, or I₂).
- In all other cases (including dilute solutions or when sulfate/nitrate ions are present), oxygen gas is produced from the discharge of OH⁻ ions.
Four Key Aqueous Electrolysis Examples
1. Concentrated Hydrochloric Acid (HCl)
Ions present: H⁺, Cl⁻, OH⁻ (from water)
- Cathode: H⁺ ions gain electrons → hydrogen gas (H₂). Test: squeaky pop with a burning splint.
- Anode: Cl⁻ ions are in concentrated solution → chlorine gas (Cl₂). Test: bleaches damp litmus paper (turns it white).
2. Concentrated Sodium Chloride Solution (Brine)
Ions present: Na⁺, H⁺, Cl⁻, OH⁻
- Cathode: Na is more reactive than H, so hydrogen gas is produced (H⁺ ions are discharged). Na⁺ ions remain in solution.
- Anode: Concentrated Cl⁻ ions → chlorine gas (Cl₂).
- Remaining in solution: Na⁺ and OH⁻ → sodium hydroxide (NaOH) solution. This is a very important industrial process!
3. Copper(II) Sulfate Solution (CuSO₄)
Ions present: Cu²⁺, H⁺, SO₄²⁻, OH⁻
- Cathode: Cu is less reactive than H, so copper metal is deposited (a pink/brown coating on the electrode). Cu²⁺ ions are discharged preferentially.
- Anode: SO₄²⁻ is not a halide, so OH⁻ is discharged → oxygen gas. Test: relights a glowing splint.
- Observation: The blue colour of the solution fades over time as Cu²⁺ ions are removed and deposited as copper metal.
4. Dilute Sulfuric Acid (H₂SO₄)
Ions present: H⁺, SO₄²⁻, OH⁻
- Cathode: H⁺ ions gain electrons → hydrogen gas.
- Anode: SO₄²⁻ is not a halide, so OH⁻ is discharged → oxygen gas.
- Key observation: Both products are gases. The volume of hydrogen collected is twice the volume of oxygen (2:1 ratio), because of the equation: 2H₂O → 2H₂ + O₂
| Electrolyte | Cathode Product | Anode Product | Left in Solution |
|---|---|---|---|
| Conc. HCl | Hydrogen | Chlorine | Water |
| Conc. NaCl (brine) | Hydrogen | Chlorine | NaOH |
| CuSO₄(aq) | Copper | Oxygen | H₂SO₄ |
| Dilute H₂SO₄ | Hydrogen | Oxygen | H₂SO₄ |
Supplement
Half-Equations for Aqueous Electrolysis
Concentrated HCl:
Cathode: 2H⁺ + 2e⁻ → H₂
Anode: 2Cl⁻ → Cl₂ + 2e⁻
Concentrated NaCl (brine):
Cathode: 2H⁺ + 2e⁻ → H₂
Anode: 2Cl⁻ → Cl₂ + 2e⁻
Copper(II) sulfate solution:
Cathode: Cu²⁺ + 2e⁻ → Cu
Anode: 4OH⁻ → 2H₂O + O₂ + 4e⁻
Dilute sulfuric acid:
Cathode: 2H⁺ + 2e⁻ → H₂
Anode: 4OH⁻ → 2H₂O + O₂ + 4e⁻
Exam Tip - The OH⁻ Half-Equation
The anode half-equation for oxygen production (4OH⁻ → 2H₂O + O₂ + 4e⁻) is one of the trickiest to remember. Notice: 4 hydroxide ions produce 2 water molecules, 1 oxygen molecule, and release 4 electrons. Check: 4 O on left, 2+2=4 O on right. 4 H on left, 4 H on right. Charge: 4- on left, 4- (from 4e⁻) on right.
Selective Discharge of Ions
When there are competing ions at an electrode, which one gets discharged? Two factors determine this:
- Position in the electrochemical series (reactivity series): At the cathode, the less reactive metal ion (lower in the series) is discharged preferentially. At the anode, the ion that is most easily oxidised is discharged first.
- Concentration: If one ion is present in much higher concentration than its competitor, it may be discharged even if it would not normally be preferred. This is why concentrated NaCl produces chlorine at the anode, but dilute NaCl would produce oxygen.
For example, in dilute NaCl solution, OH⁻ is discharged at the anode (producing oxygen) because the Cl⁻ concentration is too low. In concentrated NaCl, the high concentration of Cl⁻ overcomes the preference, and chlorine is produced instead.
Industrial Uses of Electrolysis
1. Electroplating
Electroplating is coating a metal object with a thin layer of another metal using electrolysis. For example, coating a steel spoon with silver to make it look attractive and prevent rusting.
How it works:
- The object to be plated is made the cathode (negative electrode)
- The plating metal (e.g. silver) is made the anode (positive electrode)
- The electrolyte is a solution containing ions of the plating metal (e.g. silver nitrate solution)
During electroplating:
- Metal ions from the solution are deposited onto the object at the cathode
- The anode dissolves, replacing the metal ions in the solution
- The concentration of the solution stays constant
d.c. supply
(-) ________ (+)
| | | |
| |________| |
| |
____|____ _____|_____
| Object | | Plating |
| to be | | metal |
| plated | | (e.g. |
| CATHODE | | silver) |
|_________| | ANODE |
\ |___________|
\ /
------\----------/--------
| Solution of plating |
| metal ions |
| (e.g. AgNO3 solution) |
|_________________________|
Electroplating setup: the object is the cathode, the plating metal is the anode.
Supplement
2. Purification of Copper
Copper extracted from its ore is only about 99% pure. For electrical wiring, copper must be 99.99% pure because impurities increase electrical resistance. Electrolysis purifies it.
Setup:
- Anode: the impure copper
- Cathode: a thin sheet of pure copper
- Electrolyte: copper(II) sulfate solution
What happens:
- At the anode: impure copper dissolves. Cu atoms lose electrons: Cu → Cu²⁺ + 2e⁻ (oxidation). Impurities that are less reactive than copper (like silver, gold, platinum) do NOT dissolve - they fall to the bottom as anode sludge (which is valuable!).
- At the cathode: Cu²⁺ ions from the solution gain electrons: Cu²⁺ + 2e⁻ → Cu (reduction). Pure copper is deposited.
- The impure copper anode gets smaller, the pure copper cathode gets bigger.
- The concentration of Cu²⁺ in the solution stays roughly constant (copper dissolves from the anode at the same rate it deposits on the cathode).
Anode: Cu → Cu²⁺ + 2e⁻
Impure copper dissolves (oxidation)
Cathode: Cu²⁺ + 2e⁻ → Cu
Pure copper is deposited (reduction)
3. Extraction of Aluminium
Aluminium is too reactive to be extracted by reduction with carbon (unlike iron). Instead, it is extracted by electrolysis of molten aluminium oxide (Al₂O₃).
The problem: Al₂O₃ has a very high melting point (about 2072°C), which would make electrolysis extremely expensive. The solution: Al₂O₃ is dissolved in molten cryolite (Na₃AlF₆), which lowers the melting point to about 950°C, saving energy and cost.
Setup (the Hall-Heroult process):
- Electrolyte: Al₂O₃ dissolved in molten cryolite
- Cathode: carbon lining of the steel tank (where molten aluminium collects at the bottom)
- Anode: carbon (graphite) rods that dip into the electrolyte from above
Half-equations:
Cathode: Al³⁺ + 3e⁻ → Al
Aluminium ions gain electrons (reduction) - molten aluminium sinks to the bottom
Anode: 2O²⁻ → O₂ + 4e⁻
Oxide ions lose electrons (oxidation) - oxygen gas is produced
Key point: The carbon anodes burn away because the oxygen produced at high temperature reacts with the carbon electrodes: C + O₂ → CO₂. The anodes must be regularly replaced.
Why is aluminium expensive?
- Electrolysis requires huge amounts of electricity (continuous high current)
- The carbon anodes need regular replacement
- High temperatures must be maintained
- Aluminium smelters are usually built near cheap hydroelectric power sources
Worked Examples
Worked Example 1
Predict the products of electrolysis of concentrated sodium chloride solution (brine) using inert electrodes. Write the half-equations.
Show Solution
Step 1: Identify all ions
NaCl dissolves: Na⁺ + Cl⁻. Water provides: H⁺ + OH⁻. So four ions are present: Na⁺, H⁺, Cl⁻, OH⁻.
Step 2: Cathode prediction
Na⁺ and H⁺ compete. Sodium is more reactive than hydrogen (above H in reactivity series), so H⁺ is discharged. Half-equation: 2H⁺ + 2e⁻ → H₂. Product: hydrogen gas.
Step 3: Anode prediction
Cl⁻ and OH⁻ compete. Cl⁻ is a halide ion and the solution is concentrated, so Cl⁻ is discharged. Half-equation: 2Cl⁻ → Cl₂ + 2e⁻. Product: chlorine gas.
Step 4: What remains
Na⁺ and OH⁻ remain in solution, forming sodium hydroxide (NaOH).
Cathode: hydrogen gas. Anode: chlorine gas. Solution becomes NaOH.
Worked Example 2
Copper(II) sulfate solution is electrolysed using graphite electrodes. Describe what you observe at each electrode and explain why the blue colour fades.
Show Solution
Step 1: Identify ions
CuSO₄ provides: Cu²⁺ + SO₄²⁻. Water provides: H⁺ + OH⁻.
Step 2: Cathode
Cu²⁺ and H⁺ compete. Copper is less reactive than hydrogen, so Cu²⁺ is discharged. Observation: a pink/brown coating of copper metal appears on the cathode. Half-equation: Cu²⁺ + 2e⁻ → Cu.
Step 3: Anode
SO₄²⁻ and OH⁻ compete. Sulfate is not a halide, so OH⁻ is discharged. Observation: bubbles of colourless gas (oxygen) at the anode. Half-equation: 4OH⁻ → 2H₂O + O₂ + 4e⁻.
Step 4: Colour change
The blue colour of the solution is due to Cu²⁺ ions. As these are removed from solution and deposited as copper at the cathode, the blue colour fades. Eventually the solution becomes colourless (if electrolysis continues long enough).
Cathode: pink/brown copper metal deposited. Anode: colourless oxygen gas bubbles. Blue colour fades as Cu²⁺ ions are removed from solution.
Worked Example 3
In the purification of copper, explain why the impure copper is used as the anode and not the cathode.
Show Solution
Step 1: What happens at the anode
At the anode, atoms lose electrons and go into solution as ions. If impure copper is the anode, copper atoms dissolve: Cu → Cu²⁺ + 2e⁻. The impurities that are less reactive (Au, Ag, Pt) do not dissolve and fall as sludge. Impurities more reactive than copper dissolve but are not deposited at the cathode.
Step 2: What happens at the cathode
At the cathode, ions gain electrons and are deposited as pure metal. Only Cu²⁺ ions are deposited (not the impurity ions which are harder to reduce). So pure copper builds up on the cathode.
Step 3: Why not the other way round?
If impure copper were the cathode, it would just gain more copper on top (mixed with impurities). The separation would not occur. We need the impure copper to dissolve (anode) so the impurities can be left behind.
The impure copper must be the anode so it dissolves, releasing Cu²⁺ ions while impurities fall as sludge. Pure Cu²⁺ ions then deposit as pure copper on the cathode.
Exam Tips for 4.2
1. Always state the FOUR ions present when discussing aqueous electrolysis. E.g., for NaCl(aq): Na⁺, H⁺, Cl⁻, OH⁻. Examiners love to see this.
2. For the cathode rule, remember: metals ABOVE hydrogen in the reactivity series = hydrogen gas produced. Metals BELOW hydrogen = metal deposited.
3. For the anode rule: halide + concentrated = halogen produced. Everything else = oxygen.
4. When describing copper purification, always mention: impure copper anode gets smaller, pure copper cathode gets bigger, anode sludge forms.
5. For aluminium extraction, always mention cryolite and its purpose (lowers the melting point of Al₂O₃, saving energy). Also mention that carbon anodes burn away and need replacing.
6. When the blue colour of CuSO₄ fades, say "Cu²⁺ ions are removed from solution" - do not just say "copper is removed."
Apply It: Real-World Chemistry
Electrolysis of aqueous solutions is one of the most important industrial processes in the world. See how it connects to real factories, refineries, and everyday products.
1
The Hindalco aluminium smelter in Renukoot, Uttar Pradesh, is one of the largest in Asia. It is located near the Rihand Dam because electrolysis of aluminium oxide requires enormous amounts of electricity. The plant produces over 500,000 tonnes of aluminium per year.
Why is the smelter located near a hydroelectric dam, and why must the carbon anodes be replaced regularly?
Electricity cost
Extracting aluminium by electrolysis requires huge amounts of electrical energy because Al³⁺ ions need 3 electrons each to be reduced, and very high temperatures (~950°C) must be maintained. Hydroelectric power provides cheap, continuous electricity, making the process economically viable.
Carbon anodes
At the anode, oxygen gas is produced (2O²⁻ → O₂ + 4e⁻). At the high operating temperature, this oxygen reacts with the carbon anodes: C + O₂ → CO₂. The carbon literally burns away and must be replaced regularly.
Key Takeaway
Aluminium smelters are always built near cheap electricity sources. The anodes burn away because oxygen produced at ~950°C reacts with the carbon electrodes.
2
A jewellery workshop in Hatton Garden, London, electroplates brass rings with a thin layer of gold to create affordable gold-look jewellery. The rings are dipped into a solution of gold(III) chloride (AuCl₃) and connected to a power supply.
Which electrode should the brass ring be connected to, and why? What is the plating metal used as?
Setup
The brass ring must be the cathode (negative electrode). Au³⁺ ions from the gold chloride solution are attracted to the cathode. They gain 3 electrons: Au³⁺ + 3e⁻ → Au, depositing a thin layer of gold metal on the ring. A pure gold bar is used as the anode - it dissolves to replace Au³⁺ ions in the solution.
Key Takeaway
In electroplating, the object to be plated is always the cathode. Metal ions from the solution deposit on it. The plating metal (anode) dissolves to keep the ion concentration constant.
3
At the Sterlite Copper refinery in Tuticorin, Tamil Nadu, impure copper (about 99% pure) is refined to 99.99% purity for electrical wiring. They use large electrolytic cells with alternating anodes and cathodes immersed in copper sulfate solution.
What happens to the impurities like gold and silver during the purification process?
Anode sludge
When the impure copper anode dissolves (Cu → Cu²⁺ + 2e⁻), only the copper atoms are oxidised. Gold and silver are less reactive than copper, so they do not dissolve. Instead, they fall to the bottom of the tank as "anode sludge" or "anode mud." This sludge is collected and is actually very valuable - the recovered gold and silver often help pay for the refining process!
More reactive impurities
Impurities more reactive than copper (like iron and zinc) do dissolve into the solution as ions. However, because they are more reactive than copper, they are not deposited at the cathode - they remain in solution. Only Cu²⁺ ions are reduced at the cathode.
Key Takeaway
Anode sludge contains precious metals (Au, Ag) that did not dissolve. This is a bonus revenue stream for copper refineries!
4
The Chlor-alkali plant at INOVYN in Runcorn, Cheshire, UK, electrolyses concentrated brine (NaCl solution) on an enormous scale. The plant produces three valuable products simultaneously: chlorine, hydrogen, and sodium hydroxide.
Identify where each product forms and give one industrial use for each of the three products.
Products
Chlorine (Cl₂) forms at the anode from Cl⁻ ions. Hydrogen (H₂) forms at the cathode from H⁺ ions (Na is too reactive). Sodium hydroxide (NaOH) is left in solution from Na⁺ and OH⁻ ions that are not discharged.
Uses
Chlorine: water purification, making PVC plastic, bleach. Hydrogen: making margarine (hydrogenation of vegetable oils), fuel, making ammonia (Haber process). Sodium hydroxide: making soap, paper, textiles, and cleaning products.
Key Takeaway
The chlor-alkali industry is one of the most important applications of electrolysis. Brine (cheap raw material) gives three valuable products. This is why it is called the "chlor-alkali" process - chlorine + alkali (NaOH).
5
A student in Nairobi performs electrolysis of copper(II) sulfate solution using graphite electrodes. After 30 minutes, she notices the cathode has a reddish-brown coating, and the blue solution has become noticeably paler.
If she replaces the graphite electrodes with copper electrodes and repeats the experiment, will the blue colour still fade? Explain.
With graphite (inert) electrodes
Cu²⁺ ions are deposited at the cathode but nothing replaces them - OH⁻ is discharged at the anode producing O₂. So Cu²⁺ concentration decreases and the blue colour fades.
With copper electrodes
No, the blue colour will NOT fade. The copper anode dissolves (Cu → Cu²⁺ + 2e⁻) at the same rate as Cu²⁺ is deposited at the cathode (Cu²⁺ + 2e⁻ → Cu). So the concentration of Cu²⁺ in solution stays constant. The solution remains blue. This is the principle behind copper purification!
Key Takeaway
With inert electrodes, the electrolyte composition changes. With reactive (copper) electrodes, the anode replaces ions at the same rate the cathode removes them, keeping the solution constant.
Practice Questions: 4.2 Aqueous Electrolysis
Test your understanding - 20 questions
Your Score
0 / 20
QUESTION 1 OF 20
Why is electrolysis of aqueous solutions more complex than molten compounds?
In aqueous solutions, water partially ionises into H⁺ and OH⁻. These extra ions compete with the ions from the dissolved compound at each electrode, making prediction more complex.
QUESTION 2 OF 20
During electrolysis of CuSO₄(aq) with graphite electrodes, what forms at the cathode?
Copper is less reactive than hydrogen (below H in the reactivity series), so Cu²⁺ is discharged in preference to H⁺. Copper metal is deposited as a pink/brown layer.
QUESTION 3 OF 20
What gas is produced at the anode during electrolysis of dilute sulfuric acid?
SO₄²⁻ is not a halide ion, so OH⁻ is discharged at the anode instead: 4OH⁻ → 2H₂O + O₂ + 4e⁻. Oxygen gas is produced.
QUESTION 4 OF 20
What is produced at the anode during electrolysis of concentrated sodium chloride solution?
In concentrated NaCl solution, the high concentration of Cl⁻ means chloride ions are preferentially discharged at the anode: 2Cl⁻ → Cl₂ + 2e⁻. If the solution were dilute, oxygen would be produced instead.
QUESTION 5 OF 20
Why is hydrogen produced instead of sodium at the cathode during electrolysis of NaCl(aq)?
Sodium is above hydrogen in the reactivity series, meaning it is harder to reduce (more reluctant to accept electrons). So the easier-to-reduce H⁺ ions are discharged, producing hydrogen gas.
QUESTION 6 OF 20
During electrolysis of CuSO₄(aq) with inert electrodes, what happens to the blue colour?
Cu²⁺ ions give the solution its blue colour. As they are deposited as copper metal at the cathode, the concentration of Cu²⁺ decreases and the blue colour fades.
QUESTION 7 OF 20
In the purification of copper, the impure copper is used as the:
Impure copper = anode (it dissolves). Pure copper = cathode (pure copper deposits here). The anode gets smaller, the cathode gets bigger. Impurities fall as anode sludge.
QUESTION 8 OF 20
What is the purpose of cryolite in the extraction of aluminium?
Pure Al₂O₃ melts at ~2072°C. Dissolving it in molten cryolite (Na₃AlF₆) lowers the operating temperature to ~950°C, significantly reducing energy costs.
QUESTION 9 OF 20
The half-equation 4OH⁻ → 2H₂O + O₂ + 4e⁻ represents the reaction at which electrode?
Electrons appear on the product side, meaning the OH⁻ ions are losing electrons (oxidation). Oxidation occurs at the anode. This is the half-equation for oxygen production when no halide ions are present.
QUESTION 10 OF 20
In electroplating, the object to be plated is made the:
The object to be plated is the cathode so that metal ions from the solution are attracted to it and deposited as a thin metal layer. The plating metal is the anode.
QUESTION 11 OF 20
During electrolysis of dilute H₂SO₄, the volume ratio of hydrogen to oxygen collected is:
The overall equation is effectively 2H₂O → 2H₂ + O₂. Two volumes of hydrogen are produced for every one volume of oxygen, giving a 2:1 ratio.
QUESTION 12 OF 20
What three products are obtained from the electrolysis of brine (concentrated NaCl solution)?
Cathode: hydrogen (H⁺ discharged, not Na⁺). Anode: chlorine (concentrated Cl⁻ discharged). Remaining: Na⁺ + OH⁻ = sodium hydroxide (NaOH) solution.
QUESTION 13 OF 20
Why do the carbon anodes need to be replaced during aluminium extraction?
At ~950°C, the oxygen produced at the anodes reacts with the carbon: C + O₂ → CO₂. The anodes gradually burn away and must be replaced regularly.
QUESTION 14 OF 20
Which metal can be extracted by electrolysis but NOT by reduction with carbon?
Aluminium is more reactive than carbon, so carbon cannot reduce Al₂O₃. Electrolysis is the only viable method. Iron, copper, and lead are all less reactive than carbon and can be extracted by carbon reduction.
QUESTION 15 OF 20
In copper purification, what happens to the mass of the cathode during electrolysis?
The cathode gains mass because pure copper ions from the solution are deposited on it: Cu²⁺ + 2e⁻ → Cu. Meanwhile, the anode (impure copper) loses mass as it dissolves.
QUESTION 16 OF 20
What test would confirm that the gas produced at the cathode during electrolysis of dilute H₂SO₄ is hydrogen?
Hydrogen burns with a squeaky pop when a burning splint is held near it. Oxygen relights a glowing splint. CO₂ turns limewater milky. Chlorine bleaches litmus paper.
QUESTION 17 OF 20
When dilute NaCl solution (instead of concentrated) is electrolysed, what is produced at the anode?
In dilute NaCl, the Cl⁻ concentration is too low to be preferentially discharged. Instead, OH⁻ ions from water are discharged, producing oxygen. Concentration matters for anode products!
QUESTION 18 OF 20
The half-equation for the cathode reaction in the extraction of aluminium is:
At the cathode, Al³⁺ ions gain 3 electrons to form aluminium atoms (reduction): Al³⁺ + 3e⁻ → Al. The molten aluminium sinks to the bottom of the cell.
QUESTION 19 OF 20
In electroplating a fork with silver, what would the electrolyte be?
The electrolyte must contain ions of the plating metal. Silver nitrate solution contains Ag⁺ ions, which move to the cathode (fork) and deposit as silver metal: Ag⁺ + e⁻ → Ag.
QUESTION 20 OF 20
What is "anode sludge" in the purification of copper?
When the impure copper anode dissolves, impurities less reactive than copper (gold, silver, platinum) do not dissolve. They fall to the bottom as "anode sludge" and are collected - they are actually very valuable!
4.3 Hydrogen-Oxygen Fuel Cells
▼
What is a Fuel Cell?
A fuel cell is an electrochemical device that converts the chemical energy in a fuel directly into electrical energy. Unlike a battery, a fuel cell does not run out as long as fuel is continuously supplied.
The most common type is the hydrogen-oxygen fuel cell, where:
- Hydrogen gas is the fuel (supplied to one electrode)
- Oxygen gas (from air) is supplied to the other electrode
- They react to produce water and electricity
Think of it as the reverse of electrolysis: in electrolysis, we use electricity to split water. In a fuel cell, we combine hydrogen and oxygen to make water and generate electricity.
Hydrogen Oxygen
gas in gas in
| |
v v
___________ ___________
| Electrode | | Electrode |
| (anode) | | (cathode) |
| | Electrolyte| |
| H2 --> | membrane | <-- O2 |
| 2H+ + | ----------- | |
| 2e- | ions pass | O2 + 4H+ |
| | through | + 4e- --> |
|___________| | 2H2O |
| |___________|
| |
| e- flow |
|-----> LOAD (bulb) >------|
Only product: H2O (water!)
A hydrogen-oxygen fuel cell: hydrogen is oxidised at one electrode, oxygen is reduced at the other, producing water and electricity.
Advantages of Hydrogen Fuel Cells
- Clean energy: The only product is water - no carbon dioxide, no pollutants, no soot, no sulfur dioxide.
- Efficient: Fuel cells convert chemical energy directly to electrical energy without combustion. They are more efficient than burning fuels in engines (up to 60% vs ~25% for petrol engines).
- No greenhouse gas emissions during operation - no contribution to climate change at point of use.
- Quiet operation - no moving parts in the cell itself.
- Continuous power as long as fuel is supplied (unlike batteries which need recharging).
Disadvantages of Hydrogen Fuel Cells
- Hydrogen is difficult to store: It is a very light gas with a very low boiling point (-253°C), requiring either high-pressure tanks or cryogenic (extremely cold) storage. Both are expensive and bulky.
- Hydrogen is highly flammable: It has a very wide flammability range in air (4-75%) and can explode. This creates safety concerns for vehicles and filling stations.
- Hydrogen production may produce CO₂: Most hydrogen today is made by steam reforming of methane (CH₄ + H₂O → CO + 3H₂), which releases CO₂. So the overall process may not be carbon-free unless the hydrogen is produced by electrolysis using renewable energy.
- Expensive: Fuel cells often use platinum catalysts, which are very costly.
- Infrastructure: There are very few hydrogen filling stations compared to petrol stations.
Supplement
Overall Equation
2H₂ + O₂ → 2H₂O
Two molecules of hydrogen react with one molecule of oxygen to produce two molecules of water.
This is an exothermic reaction - the energy released is converted to electrical energy in the fuel cell.
This is essentially the same reaction as burning hydrogen, but in a fuel cell, the energy is released as electricity rather than as heat and light.
Worked Examples
Worked Example 1
A car manufacturer is choosing between a petrol engine, a battery-electric motor, and a hydrogen fuel cell. Compare the hydrogen fuel cell with the petrol engine in terms of emissions.
Show Solution
Petrol engine emissions
Burning petrol (a hydrocarbon) produces CO₂ (greenhouse gas contributing to climate change), H₂O, and may also produce CO (toxic), NOx (causes acid rain and smog), and particulates (soot). These pollutants harm health and the environment.
Fuel cell emissions
A hydrogen fuel cell produces only water (2H₂ + O₂ → 2H₂O). No CO₂, no CO, no NOx, no particulates. Zero harmful emissions at the point of use.
Important caveat
However, if the hydrogen was produced by steam reforming of methane, CO₂ was released during production. The fuel cell is only truly zero-carbon if the hydrogen was made using renewable energy (e.g. by electrolysis of water using solar/wind power).
Fuel cell: only water produced (zero emissions at point of use). Petrol: CO₂, CO, NOx, particulates. But hydrogen production method matters for overall carbon footprint.
Worked Example 2
Explain why hydrogen is described as both a clean fuel and a potentially dangerous one.
Show Solution
Clean fuel
When hydrogen reacts with oxygen in a fuel cell, the only product is water. There are no carbon emissions, no pollutants, and no greenhouse gases. This makes it one of the cleanest fuels available.
Dangerous fuel
Hydrogen is highly flammable and has a very wide explosive range in air (4-75% by volume). It is the lightest gas, so it escapes easily through small gaps. It burns with an almost invisible flame, making leaks difficult to detect visually. It must be stored under very high pressure or at extremely low temperatures (-253°C), both of which present safety challenges.
Clean: only produces water, no pollutants. Dangerous: highly flammable, wide explosive range, difficult to store safely, invisible flame.
Worked Example 3
A fuel cell produces 2.0 mol of water. How many moles of hydrogen and oxygen were consumed?
Show Solution
Step 1: Write the equation
2H₂ + O₂ → 2H₂O
Step 2: Use the mole ratio
From the equation: 2 mol H₂ : 1 mol O₂ : 2 mol H₂O. So 2 mol H₂O requires 2 mol H₂ and 1 mol O₂.
2.0 mol of hydrogen and 1.0 mol of oxygen were consumed.
Exam Tips for 4.3
1. When asked for advantages of fuel cells, always start with "the only product is water" - this is the most important point. Then mention efficiency and no CO₂ emissions.
2. For disadvantages, the three key points are: (1) hydrogen is difficult to store, (2) hydrogen is flammable, (3) hydrogen production may involve fossil fuels and produce CO₂.
3. Do not say fuel cells "produce no pollution" without qualification. Say "no pollution at the point of use" - because pollution may have occurred during hydrogen production.
4. The equation 2H₂ + O₂ → 2H₂O must be balanced with a 2 in front of H₂ and H₂O. Do not write H₂ + O₂ → H₂O (unbalanced!).
Apply It: Real-World Chemistry
Fuel cells are powering everything from buses to space missions. See how this clean technology is being used around the world right now.
1
The Tokyo Metropolitan Government operates a fleet of hydrogen fuel cell buses (Toyota Sora) for public transport. These buses produce only water vapour from their exhaust. Japan aims to become a "hydrogen society" and built hydrogen filling stations across Tokyo for the 2020 Olympics.
Why does Japan, which imports nearly all its fossil fuels, see hydrogen fuel cells as strategically important?
Energy security
Japan imports about 90% of its energy. Hydrogen can be produced domestically by electrolysis of water using renewable energy (solar, wind, geothermal). This reduces Japan's dependence on imported oil and gas, improving energy security.
Environmental benefits
Fuel cell buses in Tokyo produce only water (2H₂ + O₂ → 2H₂O) - no CO₂, no NOx, no particulates. This improves air quality in a densely populated city and helps Japan meet its climate targets.
Key Takeaway
Hydrogen fuel cells reduce dependence on fossil fuel imports AND eliminate tailpipe emissions. Countries with limited fossil fuel resources see hydrogen as a path to energy independence.
2
NTPC (India's largest power company) has launched a pilot project to run hydrogen fuel cell buses in Delhi and Leh. However, critics point out that most of India's hydrogen is currently produced from natural gas (methane) through steam reforming.
Why might these fuel cell buses not be as "green" as they appear? What would make them truly zero-carbon?
The problem
Steam reforming of methane produces hydrogen but also releases CO₂: CH₄ + H₂O → CO + 3H₂ (and then CO + H₂O → CO₂ + H₂). So while the bus itself only emits water, the hydrogen production process released greenhouse gases. This is called "grey hydrogen."
The solution
To be truly zero-carbon, hydrogen should be produced by electrolysis of water using renewable energy (solar or wind). This is called "green hydrogen." India has massive solar energy potential that could power electrolysers to produce green hydrogen.
Key Takeaway
A fuel cell vehicle is only as clean as the method used to produce its hydrogen. Grey hydrogen (from fossil fuels) still has a carbon footprint. Green hydrogen (from renewable electrolysis) is truly zero-carbon.
3
NASA has used hydrogen fuel cells in spacecraft since the 1960s (Gemini and Apollo missions). The fuel cells not only generated electricity for the spacecraft but also produced drinking water for the astronauts!
Explain why fuel cells were ideal for space missions, and how the water by-product was useful.
Why ideal for space
Fuel cells are lightweight for the energy they produce, have no moving parts (reliable), operate quietly, and produce electricity continuously as long as hydrogen and oxygen are supplied. Spacecraft already carry liquid hydrogen and oxygen as rocket propellant, so the fuel was readily available.
Water as a bonus
The reaction 2H₂ + O₂ → 2H₂O produces pure water. In space, water is extremely precious and expensive to transport. The fuel cells served a dual purpose: generating electricity AND providing drinking water for the crew.
Key Takeaway
In space, the "disadvantage" of producing water becomes an advantage! Fuel cells uniquely provided both electricity and drinking water from the hydrogen and oxygen already on board.
4
Aberdeen in Scotland operates the largest fleet of hydrogen fuel cell buses in Europe. The hydrogen is produced locally by electrolysis using electricity from an offshore wind farm in the North Sea.
Explain why this is considered a truly "zero-carbon" transport system, from production to use.
Production
The hydrogen is made by electrolysis of water (2H₂O → 2H₂ + O₂) powered by wind energy. Wind is renewable and produces no CO₂. So the hydrogen production step is carbon-free.
Use
The fuel cell buses combine hydrogen and oxygen: 2H₂ + O₂ → 2H₂O. The only emission is water. No CO₂ at any stage from production to use.
Key Takeaway
When hydrogen is produced by renewable-powered electrolysis and used in fuel cells, the entire cycle is zero-carbon: wind → electricity → electrolysis → hydrogen → fuel cell → water. A genuinely circular, clean system.
5
A student argues that battery-electric vehicles (like Tesla) are better than hydrogen fuel cell vehicles because you can charge a Tesla at home, but there are almost no hydrogen filling stations in India. Her lab partner argues that fuel cells are better because they refuel in 5 minutes vs several hours for battery charging.
Evaluate both arguments. What are the key trade-offs between battery-electric and hydrogen fuel cell vehicles?
Battery-electric advantages
Existing electrical infrastructure (plug in at home), higher energy efficiency (battery ~90% vs fuel cell ~60%), lower running costs, well-established charging networks in cities. No need to build entirely new infrastructure.
Fuel cell advantages
Much faster refuelling (5 minutes vs hours), longer range per fill-up, lighter than heavy battery packs (important for trucks and buses), no battery degradation over time, no rare earth mining for batteries.
Fuel cell disadvantages
Almost no hydrogen infrastructure, hydrogen storage is complex and expensive, platinum catalysts are costly, hydrogen production is often from fossil fuels.
Key Takeaway
Both technologies have their place. Batteries are better for short-range personal vehicles. Fuel cells may be better for long-range, heavy-duty transport (buses, trucks, trains) where quick refuelling and low weight matter most.
Practice Questions: 4.3 Fuel Cells
Test your understanding - 20 questions
Your Score
0 / 20
QUESTION 1 OF 20
In a hydrogen-oxygen fuel cell, the only product is:
A hydrogen-oxygen fuel cell combines hydrogen and oxygen to produce water and electricity: 2H₂ + O₂ → 2H₂O. Water is the only chemical product.
QUESTION 2 OF 20
What is the balanced overall equation for a hydrogen fuel cell?
The balanced equation is 2H₂ + O₂ → 2H₂O. Option B is not balanced (2 O atoms on left, 1 on right). Option D is the equation for electrolysis of water (reverse reaction).
QUESTION 3 OF 20
Which of the following is NOT an advantage of hydrogen fuel cells?
Hydrogen storage is actually a major disadvantage. Hydrogen must be stored at very high pressures or very low temperatures (-253°C), both of which are expensive and technically challenging.
QUESTION 4 OF 20
Why might the production of hydrogen for fuel cells still contribute to CO₂ emissions?
Most hydrogen is made by steam reforming of methane (natural gas), which releases CO₂. The fuel cell itself only produces water, but the upstream production process often involves fossil fuels.
QUESTION 5 OF 20
A fuel cell differs from a battery because a fuel cell:
A fuel cell receives fuel continuously from outside (hydrogen and oxygen are piped in). A battery has a fixed amount of reactants stored inside. When a battery's reactants are used up, it goes flat.
QUESTION 6 OF 20
Which safety concern is associated with hydrogen as a fuel?
Hydrogen is extremely flammable with a wide explosive range (4-75% in air). It is not toxic, radioactive, or fume-producing - it only makes water when it reacts with oxygen. The main concern is fire and explosion risk.
QUESTION 7 OF 20
A hydrogen fuel cell converts:
A fuel cell converts the chemical energy stored in the hydrogen-oxygen bond formation directly into electrical energy. This is the reverse of electrolysis (which converts electrical to chemical energy).
QUESTION 8 OF 20
If 4 mol of hydrogen are used in a fuel cell, how many moles of water are produced?
From the equation 2H₂ + O₂ → 2H₂O, the mole ratio of H₂ to H₂O is 2:2 = 1:1. So 4 mol H₂ produces 4 mol H₂O.
QUESTION 9 OF 20
Which method of hydrogen production would make fuel cells truly zero-carbon?
Electrolysis of water using renewable energy (solar, wind, hydroelectric) produces hydrogen with zero CO₂ emissions at any stage. This is called "green hydrogen." All other methods listed involve fossil fuels.
QUESTION 10 OF 20
Why is hydrogen difficult to store?
Hydrogen has a very low density (lightest element) and boils at -253°C. To store useful amounts, it must be compressed to very high pressures (700 bar) or cooled to cryogenic temperatures, both of which are expensive and technically demanding.
QUESTION 11 OF 20
Compared to a petrol engine, a hydrogen fuel cell is:
Fuel cells are about 60% efficient vs ~25% for petrol engines. They convert chemical energy directly to electrical energy without the wasteful intermediate step of combustion and mechanical conversion.
QUESTION 12 OF 20
The reaction in a hydrogen fuel cell is the reverse of:
Electrolysis of water: 2H₂O → 2H₂ + O₂ (uses electrical energy). Fuel cell: 2H₂ + O₂ → 2H₂O (produces electrical energy). They are exact reverse reactions.
QUESTION 13 OF 20
NASA used hydrogen fuel cells in spacecraft. An additional benefit beyond electricity was:
The pure water (H₂O) produced by the fuel cell was collected and used as drinking water by the astronauts. In space, every resource is precious, so this dual-purpose function was extremely valuable.
QUESTION 14 OF 20
What is "green hydrogen"?
"Green hydrogen" is hydrogen produced by electrolysis of water using renewable energy sources (wind, solar, hydroelectric). The entire production process generates zero CO₂ emissions.
QUESTION 15 OF 20
If 3 mol of oxygen react in a fuel cell, how many moles of hydrogen are needed?
From 2H₂ + O₂ → 2H₂O, the ratio H₂:O₂ = 2:1. So 3 mol O₂ requires 6 mol H₂.
QUESTION 16 OF 20
Which is a major barrier to widespread use of fuel cell cars?
Infrastructure is the biggest barrier. There are very few hydrogen filling stations worldwide compared to petrol stations or electric charging points. Building a hydrogen distribution network is expensive.
QUESTION 17 OF 20
In a fuel cell, what type of energy conversion takes place?
A fuel cell converts chemical energy (from the reaction of hydrogen and oxygen) directly into electrical energy. Electrolysis is the reverse: electrical to chemical.
QUESTION 18 OF 20
Why are fuel cells described as being more efficient than combustion engines?
Combustion engines waste a lot of energy as heat. Fuel cells bypass combustion entirely, converting chemical energy directly to electrical energy. This direct conversion is inherently more efficient.
QUESTION 19 OF 20
The flammability range of hydrogen in air is approximately:
Hydrogen has a very wide flammability range of 4-75% in air. This means it can ignite at concentrations as low as 4% and as high as 75%. By comparison, petrol vapour has a much narrower range (1.4-7.6%). This wide range makes hydrogen particularly hazardous.
QUESTION 20 OF 20
Which statement about hydrogen fuel cells is INCORRECT?
The fuel cell itself produces no CO₂. However, if hydrogen is produced by steam reforming of methane, CO₂ IS produced during hydrogen manufacture. So it is incorrect to say "never any CO₂ at any stage." It depends on how the hydrogen was made.
4.4 Electrochemical Cells [Supplement]
▼
Supplement - Extended Tier Only
What is an Electrochemical Cell?
An electrochemical cell (sometimes called a voltaic cell or galvanic cell) is a device that produces electricity from a chemical reaction. It is the opposite of electrolysis - instead of using electricity to drive a reaction, the reaction itself generates electricity.
The simplest electrochemical cell consists of:
- Two different metals (electrodes) placed in an electrolyte (a solution that conducts electricity, e.g. dilute sulfuric acid or salt solution)
- The metals are connected by a wire (external circuit) with a voltmeter or bulb
When two different metals are placed in an electrolyte and connected, a voltage (potential difference) is produced. The chemical reactions at the electrodes push electrons through the external wire, creating a current.
Voltmeter
V [ 1.1V ]
| |
wire | | wire
(e- flow this way -->)
| |
______|____ _____|_____
| Zinc || Copper |
| strip || strip |
| (more || (less |
| reactive)|| reactive)|
| NEGATIVE || POSITIVE |
|___________||___________|
\ /
--------\--------/----------
| Electrolyte |
| (e.g. dilute H2SO4 |
| or salt solution) |
|___________________________|
A simple electrochemical cell: two different metals in an electrolyte produce a voltage. Electrons flow from the more reactive metal to the less reactive one.
Which Metal is the Negative Electrode?
The more reactive metal is always the negative electrode. This is because the more reactive metal has a greater tendency to lose electrons and form ions. These electrons flow through the external wire to the less reactive metal.
For example, in a zinc-copper cell:
- Zinc is more reactive than copper, so zinc is the negative electrode
- Zinc atoms lose electrons: Zn → Zn²⁺ + 2e⁻ (oxidation)
- The electrons flow through the wire to the copper electrode
- At the copper electrode (positive), the electrons are accepted by ions in solution
The Greater the Reactivity Difference, the Greater the Voltage
The voltage produced by an electrochemical cell depends on how far apart the two metals are in the reactivity series:
| Metal Combination | Reactivity Difference | Approximate Voltage |
|---|---|---|
| Zinc + Copper | Medium | ~1.1 V |
| Magnesium + Copper | Large | ~2.7 V |
| Iron + Copper | Small | ~0.8 V |
| Zinc + Iron | Small | ~0.3 V |
| Magnesium + Zinc | Medium | ~1.6 V |
| Copper + Copper | None (same metal) | 0 V |
Key rule: If two strips of the same metal are used, there is no voltage produced (0 V) because there is no difference in reactivity.
Memory Trick - Reactivity and Voltage
Think of the reactivity series as a hill. The more reactive metal is higher up. The voltage is like the "height difference" between them. Bigger height difference = bigger voltage. Two metals at the same height = no voltage (flat = no flow).
Connecting to the Reactivity Series
The reactivity series (most to least reactive):
K > Na > Ca > Mg > Al > C > Zn > Fe > H > Cu > Ag > Au > Pt
Any two metals from this series can make an electrochemical cell. The further apart they are, the higher the voltage. The more reactive metal is always the negative electrode (it oxidises - loses electrons).
For example:
- Mg and Cu: Very far apart → high voltage (~2.7 V). Mg is negative, Cu is positive.
- Fe and Cu: Close together → low voltage (~0.8 V). Fe is negative, Cu is positive.
- Zn and Ag: Far apart → high voltage. Zn is negative, Ag is positive.
Worked Examples
Worked Example 1
A student sets up an electrochemical cell using magnesium and iron strips in dilute sulfuric acid. Which metal is the negative electrode, and predict the relative voltage compared to a zinc-copper cell.
Show Solution
Step 1: Identify the more reactive metal
In the reactivity series: Mg > Al > Zn > Fe > Cu. Magnesium is more reactive than iron.
Step 2: Identify the negative electrode
The more reactive metal is always the negative electrode. So magnesium is the negative electrode, and iron is the positive electrode.
Step 3: Predict the voltage
The gap between Mg and Fe in the reactivity series is larger than the gap between Zn and Cu. Therefore, the Mg-Fe cell will produce a higher voltage than the Zn-Cu cell (~1.9 V vs ~1.1 V).
Magnesium is the negative electrode. The voltage will be higher than a Zn-Cu cell because the reactivity difference between Mg and Fe is greater than between Zn and Cu.
Worked Example 2
Explain why two strips of the same metal placed in an electrolyte produce no voltage.
Show Solution
Explanation
Voltage is produced because the two metals have different tendencies to lose electrons (different reactivities). One metal loses electrons more readily and becomes the negative electrode, pushing electrons through the wire to the other metal. If both strips are the same metal, they have exactly the same tendency to lose electrons. There is no "driving force" to push electrons in either direction. The reactivity difference is zero, so the voltage is zero.
Same metals have the same reactivity, so there is no difference in tendency to lose electrons. No driving force for electron flow = no voltage (0 V).
Worked Example 3
Three cells are set up: (i) Zn and Cu, (ii) Mg and Cu, (iii) Fe and Cu. Arrange them in order of increasing voltage and explain your reasoning.
Show Solution
Step 1: Check the reactivity series
Order of reactivity: Mg > Zn > Fe > Cu. All three cells use copper as the positive electrode.
Step 2: Determine reactivity gaps
Fe-Cu: smallest gap (closest together in series). Zn-Cu: medium gap. Mg-Cu: largest gap (furthest apart in series).
Step 3: Order by voltage
Increasing voltage: Fe-Cu (~0.8 V) < Zn-Cu (~1.1 V) < Mg-Cu (~2.7 V).
Order of increasing voltage: Fe-Cu < Zn-Cu < Mg-Cu. The greater the difference in reactivity, the greater the voltage.
Exam Tips for 4.4
1. The more reactive metal is ALWAYS the negative electrode. This is because it has a greater tendency to lose electrons (oxidise).
2. If asked "what would happen if both electrodes were the same metal?" - the answer is always "no voltage would be produced because there is no difference in reactivity."
3. Electrons flow from the more reactive metal (negative) through the external wire to the less reactive metal (positive). Inside the cell, ions carry the current through the electrolyte.
4. Do not confuse electrochemical cells with electrolysis. In an electrochemical cell, the chemical reaction generates electricity. In electrolysis, electricity drives a chemical reaction.
5. Common exam question: "Predict which combination of metals would give the highest voltage." Choose the two metals that are furthest apart in the reactivity series.
Apply It: Real-World Chemistry
Electrochemical cells are the basis for every battery you have ever used. From the phone in your hand to emergency backup power, they are everywhere.
1
In a chemistry practical at a school in Mumbai, students test different pairs of metal strips (Mg-Cu, Zn-Cu, Fe-Cu, Zn-Fe) in lemon juice as the electrolyte. They record the voltmeter readings for each combination.
Which pair would give the highest reading? Why does lemon juice work as an electrolyte?
Highest voltage
Mg-Cu would give the highest reading because magnesium and copper are furthest apart in the reactivity series. The order would be: Mg-Cu > Zn-Cu > Zn-Fe > Fe-Cu.
Lemon juice as electrolyte
Lemon juice contains citric acid, which dissociates to produce H⁺ and citrate ions. These free-moving ions allow the lemon juice to conduct electricity and complete the circuit between the two metal electrodes.
Key Takeaway
Any acidic or ionic solution can serve as an electrolyte - even lemon juice, vinegar, or salt water. The electrolyte just needs free-moving ions to carry the current.
2
The original "Baghdad Battery" (dating to around 250 BCE, found in modern Iraq) consisted of a clay jar containing a copper cylinder and an iron rod, with evidence of an acidic electrolyte (possibly vinegar or grape juice). Some archaeologists believe it may have been used for electroplating gold onto silver objects.
Explain the chemistry behind how this ancient device could have produced a voltage.
Two different metals
Iron and copper are two different metals with different reactivities. Iron is more reactive than copper (higher in the reactivity series).
Electrolyte
The vinegar or grape juice provides an acidic electrolyte containing free-moving H⁺ ions that can carry the current.
Voltage produced
Iron (more reactive) would be the negative electrode, losing electrons: Fe → Fe²⁺ + 2e⁻. These electrons flow through the wire to the copper (positive electrode). This produces a small voltage of about 0.8 V - enough for simple electroplating if multiple cells were connected in series.
Key Takeaway
The principle of electrochemical cells - two different metals in an electrolyte - is so fundamental that people may have discovered it over 2000 years ago! The chemistry is the same whether it is a modern battery or an ancient artefact.
3
A marine engineer in Southampton notices that the steel hull of a ship corrodes faster near the bronze propeller. She recommends attaching blocks of zinc to the hull near the propeller as "sacrificial anodes."
Explain how zinc blocks protect the steel hull using your knowledge of electrochemical cells.
The problem
Steel (iron) and bronze (copper alloy) in seawater (electrolyte) form an electrochemical cell. Iron is more reactive, so it acts as the negative electrode and corrodes (oxidises): Fe → Fe²⁺ + 2e⁻. This accelerates the rusting of the hull.
The solution
Zinc is more reactive than iron. When zinc blocks are attached, the zinc-iron-seawater combination forms a new electrochemical cell where zinc is the negative electrode instead of iron. Zinc corrodes preferentially (Zn → Zn²⁺ + 2e⁻), "sacrificing" itself to protect the iron. The iron now acts as the positive electrode and is protected.
Key Takeaway
Sacrificial protection uses the electrochemical series: a more reactive metal corrodes instead of the protected metal. Zinc protects iron, magnesium protects iron too. The more reactive metal is "sacrificed" to save the less reactive one.
4
At the Amara Raja battery factory in Tirupati, Andhra Pradesh (India's second-largest battery manufacturer), lead-acid batteries are produced for cars and backup power systems. These batteries use lead and lead dioxide electrodes in sulfuric acid.
How does a car battery relate to the simple electrochemical cells you studied? Why does it eventually go flat?
Same principle
A lead-acid battery works on the same principle as a simple electrochemical cell: two different electrode materials (lead and lead dioxide) in an electrolyte (sulfuric acid). The chemical difference between the electrodes drives electron flow, producing a voltage of about 2 V per cell (6 cells = 12 V car battery).
Going flat
As the battery produces electricity, the reactants (lead, lead dioxide, and sulfuric acid) are gradually converted into products (lead sulfate and water). When the reactants are used up, the chemical reaction stops and no more electrons flow - the battery is "flat." Unlike a fuel cell, a battery has a fixed amount of reactants.
Key Takeaway
All batteries are electrochemical cells at heart. The voltage depends on the electrode materials, and the battery goes flat when the reactants are consumed. Rechargeable batteries reverse the reaction by passing current through them.
5
A student in Cambridge wants to build the highest-voltage cell possible using common lab metals: copper, iron, zinc, magnesium, and aluminium strips, with dilute sulfuric acid as the electrolyte.
Which two metals should she choose, and which will be the positive and negative electrodes?
Choosing the metals
For maximum voltage, choose the two metals that are furthest apart in the reactivity series. From the available metals: Mg > Al > Zn > Fe > Cu. Magnesium and copper are the furthest apart.
Electrode assignments
Magnesium (more reactive) = negative electrode. Copper (less reactive) = positive electrode. This will produce the highest voltage, approximately 2.7 V.
Key Takeaway
For maximum voltage: pick the MOST reactive metal as the negative electrode and the LEAST reactive as the positive. The bigger the gap in reactivity, the bigger the voltage.
Practice Questions: 4.4 Electrochemical Cells
Test your understanding - 20 questions
Your Score
0 / 20
QUESTION 1 OF 20
A simple electrochemical cell requires:
A simple electrochemical cell needs two DIFFERENT metals in an electrolyte. The difference in reactivity between the metals creates a voltage. Same metals = 0 V.
QUESTION 2 OF 20
In an electrochemical cell, the more reactive metal is:
The more reactive metal has a greater tendency to lose electrons. It oxidises (loses electrons), making it the negative electrode. Electrons flow from it through the wire to the less reactive metal.
QUESTION 3 OF 20
Which pair of metals would produce the highest voltage?
Mg and Cu are the furthest apart in the reactivity series (Mg > Al > Zn > Fe > Cu). The greater the gap, the greater the voltage. Mg-Cu gives about 2.7 V.
QUESTION 4 OF 20
What voltage would be produced by two copper strips in dilute sulfuric acid?
Two strips of the same metal produce 0 V. There is no difference in reactivity, so there is no driving force for electrons to flow in one direction.
QUESTION 5 OF 20
In a zinc-copper electrochemical cell, electrons flow through the external wire from:
Zinc is more reactive, so it loses electrons more readily: Zn → Zn²⁺ + 2e⁻. These electrons flow through the external wire from zinc (negative) to copper (positive).
QUESTION 6 OF 20
An electrochemical cell converts:
An electrochemical cell converts chemical energy (from the reactions at the electrodes) to electrical energy. This is the opposite of electrolysis, which converts electrical to chemical energy.
QUESTION 7 OF 20
In a magnesium-iron cell, which metal is the positive electrode?
Magnesium is more reactive than iron, so Mg is the negative electrode. Iron, being less reactive, is the positive electrode. No external battery is needed - the cell generates its own voltage.
QUESTION 8 OF 20
What is the role of the electrolyte in an electrochemical cell?
The electrolyte contains ions that carry the current inside the cell, completing the circuit. Electrons flow through the external wire, and ions flow through the electrolyte. Without the electrolyte, the circuit would be broken.
QUESTION 9 OF 20
Arrange these cells in order of INCREASING voltage: Zn-Cu, Mg-Cu, Fe-Cu
Reactivity order: Mg > Zn > Fe > Cu. Gap from Cu: Fe-Cu (smallest) < Zn-Cu (medium) < Mg-Cu (largest). Larger gap = higher voltage.
QUESTION 10 OF 20
Which of the following could NOT act as an electrolyte in an electrochemical cell?
Hexane is a covalent liquid with no ions, so it cannot conduct electricity or act as an electrolyte. The other options all contain free-moving ions that can carry the current.
QUESTION 11 OF 20
The difference between an electrochemical cell and electrolysis is:
Electrochemical cells generate electricity from spontaneous chemical reactions. Electrolysis uses electricity to drive non-spontaneous reactions. They are essentially opposite processes.
QUESTION 12 OF 20
In a Zn-Cu cell, what happens at the zinc electrode?
Zinc is the more reactive metal (negative electrode). Zinc atoms lose electrons: Zn → Zn²⁺ + 2e⁻. This is oxidation. The electrons flow through the wire to the copper electrode.
QUESTION 13 OF 20
A student replaces the copper electrode in a Zn-Cu cell with a silver electrode. What happens to the voltage?
Silver is less reactive than copper (lower in the reactivity series). So the gap between Zn and Ag is larger than between Zn and Cu. Greater reactivity difference = greater voltage.
QUESTION 14 OF 20
Why do zinc blocks protect steel hulls from corrosion?
Zinc is more reactive than iron. In the electrochemical cell formed (Zn-Fe in seawater), zinc acts as the negative electrode and oxidises preferentially. The iron is "protected" because the zinc corrodes instead. This is called sacrificial protection.
QUESTION 15 OF 20
What carries the electric current inside the electrochemical cell (through the electrolyte)?
Inside the electrolyte, the current is carried by moving ions. In the external wire, the current is carried by electrons. Do not confuse the two - this is a common exam mistake.
QUESTION 16 OF 20
A cell is made from aluminium and copper in dilute acid. Which is the negative electrode?
Aluminium is more reactive than copper (higher in the reactivity series). The more reactive metal is always the negative electrode because it has a greater tendency to lose electrons.
QUESTION 17 OF 20
To increase the voltage of an electrochemical cell, you should:
Voltage depends on the reactivity difference between the metals. Bigger pieces may make the cell last longer but do not increase the voltage. Choosing metals further apart in the reactivity series increases the voltage.
QUESTION 18 OF 20
A student uses lemon juice as the electrolyte in a Zn-Cu cell. Why does this work?
Lemon juice contains citric acid, which dissociates to produce H⁺ ions and citrate ions. These free-moving ions allow the current to flow through the electrolyte, completing the circuit.
QUESTION 19 OF 20
What would happen if the wire connecting the two metals in an electrochemical cell were cut?
Cutting the wire breaks the external circuit. Electrons cannot flow from the negative to the positive electrode. A voltmeter would still detect a potential difference (voltage), but no current would flow. The chemical reactions at the electrodes would stop.
QUESTION 20 OF 20
Which of the following statements about electrochemical cells is correct?
Greater reactivity difference = greater voltage. The MORE reactive metal is negative (not less). Same metals give 0 V (not highest). The electrolyte can be any solution with ions (acid, alkali, or salt solution).