Watch the current graph in the first minute: it falls as hydrogen bubbles cover the copper (polarization). Use 1 s = 1 day to see the zinc and the acid get used up.
Changing the metals, the acid or the depth sets up a fresh cell.
All readings are model values for a typical school set-up: strips 2 cm wide, 200 cm³ of acid, a 5 Ω milliammeter. Real cells give similar sizes but vary. Cell voltage (e.m.f.) is what a voltmeter connected straight across the two strips would read (fresh cell: 0.8 × the difference in E° of the metals, a model rule); it falls as hydrogen covers the positive strip. While current flows, the p.d. across the ammeter is only I × R.
Any two different metals in an acid make a cell. The more reactive metal (higher in the reactivity series, more negative standard electrode potential E°) gives up electrons more easily, so it becomes the negative electrode. The further apart the two metals are, the bigger the voltage. Two strips of the same metal give 0 V. Pick two metals:
E° values are data-book standard electrode potentials (M²⁺ + 2e⁻ → M, 25 °C, 1 mol/dm³). The difference in E° is the voltage of a cell in which each metal stands in a solution of its own ions (for zinc and copper that is the Daniell cell, 1.10 V). A simple cell in acid is not that cell: the positive strip makes hydrogen, not metal, so 1.10 V does not apply and the reading is lower. The difference in E° is used here only as a measure of how far apart the metals are in the reactivity series. The lab’s model rule is: fresh-cell reading = 0.8 × the difference (zinc and copper: 0.88 V; real fresh zinc–copper cells read roughly 0.8–1.0 V and then fall as hydrogen collects). It shows the trend correctly but is not a calculation from theory, and pairs far apart are only rough: a real magnesium–copper cell usually reads less than the model’s 2.17 V. Lead is left out on purpose: in sulfuric acid it gets coated with insoluble lead sulfate.
The diagram you draw in a test: name the electrodes, show which way the electrons flow and where the hydrogen forms. Switch between the labeled diagram, a blank one to test your memory, and a quiz where you place each label yourself. You can also download both versions or print a label worksheet.
Tap a word, then tap the numbered box it belongs to (on the diagram or in the list below). Tap a filled box to take the word back. On a phone, swipe the diagram sideways to see every number.
Ten questions (from a bank of twenty) on the electrodes, the half-equations, electron flow, polarization and the reactivity series.
Tap a card, then tap the box it belongs in. Check when all twelve are placed.
Put a zinc strip and a copper strip into dilute sulfuric acid without letting them touch, and join them with wires to an ammeter. The needle moves: a current flows. Bubbles of gas appear on the copper, and over a long time the zinc gets thinner. Open the switch and the current stops. This set-up is called a simple cell (or simple voltaic cell, after Alessandro Volta, who made the first battery in 1800).
Zinc is more reactive than copper. At the zinc strip, zinc atoms lose two electrons each and go into the acid as zinc ions: Zn(s) → Zn²⁺(aq) + 2e⁻. The electrons left behind make the zinc the negative electrode. They flow through the wire and the ammeter to the copper. In the acid there are H⁺ ions and SO₄²⁻ ions. The H⁺ ions move to the copper, take the electrons and pair up: 2H⁺(aq) + 2e⁻ → H₂(g). The copper itself does not change. The sulfate ions do not react at all (spectator ions), while zinc ions build up and H⁺ ions are used up. Overall: Zn + H₂SO₄ → ZnSO₄ + H₂.
Loss of electrons is oxidation; the electrode where it happens is the anode (zinc, negative). Gain of electrons is reduction; it happens at the cathode (copper, positive). Electrons flow zinc → wire → copper; conventional current is drawn the other way, copper → zinc. In the acid the charge is carried by moving ions, not electrons. The voltage depends on the two metals: the further apart in the reactivity series (the bigger the difference in E°), the bigger the voltage. The current also depends on the acid concentration, the area of the strips and the resistance of the circuit. It falls when hydrogen bubbles cover the copper (polarization), when impure zinc reacts by itself (local action), and it stops when the zinc or the acid is used up.
Press Textbook cell and watch the current graph for a minute: it falls to about a third. Press Tap the copper strip and it jumps back up. Turn on Conventional current to see the arrows flip. In the particle view, count the ions: SO₄²⁻ never changes. Choose copper for both strips: 0 V. Then try Magnesium | copper and the 2 weeks in 14 s time-lapse.
| Where | What happens | Equation |
|---|---|---|
| Zinc strip: negative electrode, anode | oxidation (loses electrons); zinc dissolves | Zn(s) → Zn²⁺(aq) + 2e⁻ |
| Copper strip: positive electrode, cathode | reduction (gains electrons); hydrogen gas forms; copper unchanged | 2H⁺(aq) + 2e⁻ → H₂(g) |
| Overall (ionic) | add the half-equations; the electrons cancel | Zn(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂(g) |
| Overall (full) | sulfate is a spectator ion | Zn + H₂SO₄ → ZnSO₄ + H₂ |
| Wire (external circuit) | electrons flow zinc → copper; conventional current copper → zinc | charge Q = I × t |
| Solution (internal circuit) | ions carry the charge: H⁺ and Zn²⁺ drift towards the copper, SO₄²⁻ towards the zinc | no electrons in the solution |
| Difference in E° | more positive E° minus more negative E°: tells you which metal is negative and how far apart the metals are (a simple cell in acid reads less than this) | ΔE° = E°(+) − E°(−) |
| Faraday | moles of electrons = charge ÷ Faraday constant (F = 96 485 C/mol, often rounded to 96 500) | n(e⁻) = Q ÷ F; n(Zn) = n(H₂) = n(e⁻) ÷ 2 |
The reaction that runs is Zn + 2H⁺ → Zn²⁺ + H₂, so a purist would say the driving couples are Zn²⁺/Zn (−0.76 V) and H⁺/H₂ (0.00 V): standard cell potential 0.76 V, and the 1.10 V of the Daniell cell does not apply. Yet a fresh zinc–copper cell usually reads more than 0.76 V (roughly 0.8–1.0 V), and the reading does depend on the positive metal. The reason is that a strip in acid is not at an equilibrium potential. Before much hydrogen has formed, the copper’s potential is a mixed potential set by whatever can take electrons at its surface: dissolved oxygen (O₂ + 4H⁺ + 4e⁻ → 2H₂O, E° = +1.23 V), traces of copper oxide, and H⁺ ions (slowly, because making H₂ on copper needs an extra push, the overpotential). So it sits well above 0.00 V, and a more reactive positive metal (iron, tin) sits lower. Once current flows, the oxygen near the copper is used up and hydrogen covers it, so the copper’s potential drops towards (and, while current flows, below) the hydrogen value and the reading falls: that is polarization. That is why school courses, and this lab, link the voltage to how far apart the two metals are in the reactivity series. The lab’s 0.8 × ΔE° is an empirical model rule that reproduces this trend and gives a realistic fresh Zn/Cu reading; it is not a thermodynamic calculation.
Making hydrogen gas on a metal needs an extra push (the hydrogen overpotential), and bubbles stuck to the copper block part of its surface. Both lower the voltage the cell can deliver while current flows. A depolarizer (an oxidizing agent such as hydrogen peroxide or potassium dichromate; manganese(IV) oxide in the old Leclanché dry cell) removes the hydrogen as water. The Daniell cell avoids the problem completely: its copper electrode stands in copper(II) sulfate solution, so copper metal is deposited (Cu²⁺ + 2e⁻ → Cu) instead of hydrogen gas, and it gives a steady 1.10 V.
School equations write H₂SO₄ → 2H⁺ + SO₄²⁻. In fact the first H⁺ is lost completely but the second only partly: HSO₄⁻ ⇌ H⁺ + SO₄²⁻, Ka2 ≈ 1.0 × 10⁻². In 0.5–1 mol/dm³ acid many sulfate groups are present as HSO₄⁻ ions, and H⁺ is really H₃O⁺. None of this changes the electrode reactions; the lab uses the simple school picture.
Mass dissolved = Q × M ÷ (z × F), with M the molar mass, z = 2 for Zn²⁺ and F = 96 485 C/mol. The lab uses exactly this: every 2 electrons through the ammeter = 1 zinc atom dissolved = 1 H₂ molecule on the copper. Zinc lost by local action is extra: it dissolves without sending any current through the ammeter.
A simple cell is two different metals (usually zinc and copper) dipped into an electrolyte such as dilute sulfuric acid and joined by a wire. The more reactive metal gives away electrons, which flow through the wire, so the cell turns chemical energy into electrical energy. It is also called a simple voltaic cell, after Alessandro Volta.
Key takeaway: two different metals + an electrolyte + a wire = a current.Zinc is the anode: oxidation happens there (Zn → Zn²⁺ + 2e⁻) and it is the negative electrode, because zinc is more reactive than copper and its atoms give up electrons. Copper is the cathode: reduction happens there (2H⁺ + 2e⁻ → H₂) and it is the positive electrode.
Key takeaway: zinc = negative = anode; copper = positive = cathode.The bubbles are hydrogen made from the acid, not from the copper. Electrons from the zinc arrive at the copper through the wire; H⁺ ions from the acid collect at the copper surface, take the electrons and pair up into H₂ molecules. The copper only provides the surface, so its mass does not change.
Key takeaway: the hydrogen comes from H⁺ ions in the acid.No. Electrons flow only through the metals: from the zinc, through the wire and the ammeter, to the copper. Inside the acid the charge is carried by ions: H⁺ (and Zn²⁺) ions drift towards the copper and SO₄²⁻ ions towards the zinc side. Together the wire and the solution make a complete circuit.
Key takeaway: electrons in the wire, ions in the solution.A zinc atom has two outer electrons and loses both when it reacts, so the ion has two more protons than electrons: a charge of 2+. That is why the half-equation has 2e⁻, and why each zinc atom that dissolves sends two electrons round the circuit and makes one H₂ molecule (which needs two electrons).
Key takeaway: Zn loses 2 electrons → Zn²⁺.Nothing. Sulfate ions (SO₄²⁻) do not react at either strip: they are spectator ions and stay in the solution. What changes is the other ions: H⁺ ions are used up and Zn²⁺ ions build up, so the solution slowly turns from sulfuric acid into zinc sulfate solution. The particle view counts them for you.
Key takeaway: SO₄²⁻ ions stay in the solution unchanged.Mainly because of polarization: hydrogen bubbles stick to the copper, cover the surface where H⁺ ions take electrons, and push back against the cell. In the lab the current falls from 27.6 mA to about 8.4 mA in a couple of minutes. Tapping the strip or adding a depolarizer brings it back up. Over days the current also falls because the zinc and the acid are used up.
Key takeaway: bubbles on the copper (polarization) lower the current.Ordinary zinc contains tiny bits of other metals. Each bit forms a tiny cell on the zinc surface, so hydrogen bubbles form on the zinc itself and the zinc dissolves even when the switch is open. That zinc is wasted, because its electrons never go through the ammeter. Very pure zinc (or zinc coated with mercury, in old batteries) does not do this.
Key takeaway: local action = impure zinc reacting by itself, wasting zinc.A fresh zinc–copper cell in dilute sulfuric acid reads roughly 0.8–1.0 V on a voltmeter (the lab’s model gives 0.88 V), and the voltage falls as hydrogen collects on the copper. The data-book difference in E° between zinc and copper, 1.10 V, is what a Daniell cell gives, where the copper stands in copper sulfate solution; it does not apply to the simple cell, because there the positive strip makes hydrogen instead of copper.
Key takeaway: about 0.8–1.0 V at first, then less.Use two metals that are further apart in the reactivity series. Magnesium and copper (difference in E° 2.71 V) give more than zinc and copper (1.10 V), which give more than iron and copper (0.78 V); the simple cells read less than these data-book differences, but in the same order. Two strips of the same metal give 0 V. Bigger strips, stronger acid or a smaller resistance give a bigger current, but not a bigger voltage. For more voltage you can also join cells in series.
Key takeaway: further apart in the reactivity series = bigger voltage.