DIY Batteries and Supercapacitors

A Practical Guide to Electrochemical Energy Storage

Build galvanic cells, primary batteries, and activated-carbon supercapacitors at home

Introduction

A battery converts chemical energy directly into electrical energy through redox reactions. A supercapacitor stores charge electrostatically at the surface of electrodes rather than through chemistry. Both are electrochemical devices, and both can be built at home from everyday materials — the main difference is how the energy is stored and how quickly it can be delivered.

This guide covers simple galvanic cells, a progression of battery chemistries, and a DIY electric double-layer capacitor (EDLC) using activated carbon and a sodium sulfate electrolyte.

The Basics of Galvanic Cells

A galvanic cell consists of two half-cells: each is a metal electrode immersed in an electrolyte. The more reactive metal (the anode) oxidises, releasing electrons into the external circuit. Those electrons travel through the wire to the cathode, where a reduction reaction consumes them.

The driving force is the difference in standard electrode potential between the two half-reactions, measured in volts:

\[E^\circ_\text{cell} = E^\circ_\text{cathode} - E^\circ_\text{anode}\]

Electrode Reaction E° (V)
Zinc (Zn²⁺/Zn) Zn → Zn²⁺ + 2e⁻ −0.76
Iron (Fe²⁺/Fe) Fe → Fe²⁺ + 2e⁻ −0.44
Hydrogen (reference) 2H⁺ + 2e⁻ → H₂ 0.00
Copper (Cu²⁺/Cu) Cu²⁺ + 2e⁻ → Cu +0.34
Silver (Ag⁺/Ag) Ag⁺ + e⁻ → Ag +0.80

Pairing zinc (anode) with copper (cathode) gives a theoretical cell voltage of 0.34 − (−0.76) = 1.10 V — the same as a Daniell cell.

In practice, the electrolyte must complete the ionic circuit. A salt bridge (saturated KCl or Na₂SO₄ in a gel or filter paper) allows ions to migrate between half-cells without mixing the solutions.

Battery Types

Lemon Cell (Zinc–Copper)

The classic classroom experiment. The lemon provides both the electrolyte (citric acid) and a physical separator.

Materials

  • Fresh lemon (or orange, potato, pickle)
  • Zinc strip or galvanised nail
  • Copper strip or piece of copper pipe
  • Multimeter

Procedure

  1. Push the zinc and copper electrodes into the lemon about 2 cm apart — close enough to share the electrolyte, far enough not to touch
  2. Connect the multimeter leads: red to copper, black to zinc
  3. Read the open-circuit voltage (typically 0.8–1.0 V)
  4. Connect a high-brightness LED across the terminals to see it glow (needs at least 1.8 V, so wire two or three lemons in series)

What limits it: Internal resistance is high (~500–2000 Ω), so current is tiny — a few milliamps at best. The electrolyte depletes and the zinc corrodes, so the cell is not rechargeable.

Series wiring: Connect the copper of one cell to the zinc of the next. With four lemons in series you can reach ~3.5 V and run a digital clock.

Salt-Water Zinc–Carbon Cell

A flat, open cell with no special chemicals — good for measuring electrode potential differences with different metal pairs.

Materials

  • Saturated sodium chloride solution (or sodium sulfate for fewer chlorine-evolution issues)
  • Zinc strip (or galvanised sheet metal)
  • Carbon rod (from a dead AA battery, or graphite pencil lead)
  • Small container
  • Multimeter, alligator clips

Procedure

  1. Dissolve 30 g NaCl or Na₂SO₄ in 100 mL water
  2. Immerse both electrodes without contact, a few centimetres apart
  3. Measure open-circuit voltage; with Zn and carbon you should see ~0.6–0.8 V
  4. Swap the carbon for copper, iron, aluminium, and record the voltage each time

This demonstrates the electrochemical series directly — the more separated the metals, the higher the voltage.

Warning

With salt (NaCl) electrolyte and any metal anode, chlorine can evolve at the cathode. Use Na₂SO₄ solution if you want to avoid this; it produces only oxygen at the anode.

Daniell Cell

The Daniell cell (1836) was the first practical battery, powering telegraphs for decades. It avoids the rapid voltage drop of single-electrolyte cells by keeping the two half-reactions spatially separated.

Materials

  • Copper sulfate solution — 100 mL at ~1 mol/L (25 g CuSO₄·5H₂O per 100 mL)
  • Zinc sulfate (or Na₂SO₄) solution — 100 mL at ~1 mol/L
  • Copper strip (cathode)
  • Zinc strip (anode)
  • Salt bridge: a U-tube or folded strip of filter paper saturated with saturated Na₂SO₄ solution
  • Two small glass jars
  • Multimeter

Procedure

  1. Fill one jar with CuSO₄ solution; place the copper strip in it (cathode, positive terminal)
  2. Fill the other jar with ZnSO₄ or Na₂SO₄ solution; place the zinc strip in it (anode, negative terminal)
  3. Bridge the two jars with the salt bridge — it must dip into both solutions
  4. Connect leads: red to copper, black to zinc
  5. Read voltage (~1.0–1.1 V open circuit)
  6. Connect a small load (100–1000 Ω resistor) and measure current

Reactions:

\[\ce{Zn(s) -> Zn^{2+}(aq) + 2e^-} \quad \text{(anode)}\]

\[\ce{Cu^{2+}(aq) + 2e^- -> Cu(s)} \quad \text{(cathode)}\]

The blue colour of the CuSO₄ solution fades as copper deposits on the cathode electrode; the zinc strip visibly dissolves. If you use a copper sulfate–only cell (no salt bridge), the cell still works briefly but the solutions mix quickly, reducing voltage.

Aluminium–Air Cell

Aluminium–air cells have very high theoretical energy density because oxygen from the air acts as the cathode material, so you don’t need to carry it.

Materials

  • Aluminium foil (heavy-duty) or aluminium strip
  • Carbon cloth or carbon paper (cathode), or a graphite rod
  • Sodium hydroxide solution: 40 g NaOH per litre (strongly alkaline — wear gloves)
  • Shallow container

Procedure

  1. Prepare the NaOH solution and pour into the container
  2. Suspend the carbon cathode so its lower face is submerged and its upper face is exposed to air
  3. Submerge the aluminium anode fully
  4. Connect: red lead to carbon, black to aluminium
  5. Open-circuit voltage: ~1.2–1.5 V; short-circuit current can be tens of milliamps per cm²

Reaction:

\[\ce{4Al + 3O2 + 4OH^- + 2H2O -> 4Al(OH)4^-} \quad E \approx 2.7\,\text{V theoretical}\]

In practice, aluminium’s native oxide layer causes a significant voltage gap. Sodium hydroxide dissolves the oxide continuously, keeping the surface active. The aluminium anode is consumed and cannot be recharged — replace it when depleted.

Warning

NaOH is corrosive. Wear gloves and eye protection. The reaction is exothermic and produces hydrogen gas — work in a ventilated space.

Magnesium–Copper Cell

Magnesium is more reactive than zinc (E° = −2.37 V), so a Mg–Cu cell in salt water gives a noticeably higher voltage than Zn–Cu.

Materials

  • Magnesium ribbon (available from lab suppliers; used in burning experiments)
  • Copper strip
  • Saturated Na₂SO₄ or NaCl solution

Procedure

Same as the salt-water zinc–carbon cell above. Expect ~1.5–1.7 V open circuit with the Mg–Cu pair. The magnesium surface oxidises quickly — scrub it with sandpaper before immersing to get a clean surface.

Note: Magnesium has a pronounced “parasitic corrosion” effect — it corrodes even when the cell is under load, consuming electrode without generating useful current. This limits practical efficiency but makes the chemistry interesting to observe.

Copper–Aluminium Salt-Water Cell

Aluminium is among the most reactive common metals (E° = −1.66 V), yet it is stable in air because it instantly grows a thin, tenacious oxide layer (Al₂O₃). Pairing it with copper in a salt-water electrolyte reveals that latent reactivity — but the oxide layer is also why measured voltage falls well short of the 2.0 V theoretical maximum.

Materials

  • Aluminium strip or heavy-duty aluminium foil folded into a stiff strip
  • Copper strip or piece of copper pipe
  • Saturated sodium sulfate (Na₂SO₄) or sodium chloride (NaCl) solution
  • Sandpaper (120–400 grit)
  • Shallow container
  • Multimeter

Procedure

  1. Scrub the aluminium surface firmly with sandpaper immediately before use — this breaks through the oxide layer; the cell voltage drops noticeably if you skip this step
  2. Prepare the salt solution: dissolve 30 g Na₂SO₄ (or NaCl) in 100 mL water
  3. Immerse both electrodes without contact, a few centimetres apart; connect the multimeter (red to copper, black to aluminium)
  4. Read open-circuit voltage — expect 0.7–1.1 V (the passive oxide layer on aluminium suppresses the voltage below the 2.0 V theoretical value)
  5. Connect a 100–500 Ω load and observe the current; aluminium visibly bubbles and begins to corrode at the surface
  6. Sand the aluminium again after a few minutes to remove the fresh oxide and watch the voltage recover

Reactions:

\[\ce{Al -> Al^{3+} + 3e^-} \quad \text{(anode)}\]

\[\ce{O2 + 2H2O + 4e^- -> 4OH^-} \quad \text{(cathode, oxygen reduction)}\]

In aerated salt water, dissolved oxygen is reduced at the copper surface rather than hydrogen ions — this is why copper is a better cathode here than carbon: its surface catalyses oxygen reduction more effectively. Hydrogen can also evolve at the cathode when oxygen is depleted.

What limits it: The Al₂O₃ passive layer regrows within seconds of sanding, acting as a resistive film. Chloride ions in NaCl electrolyte attack the oxide locally (pitting corrosion), which is why NaCl can actually give a slightly higher, less stable current than Na₂SO₄. The aluminium is consumed and the cell is not rechargeable.

Comparison with other simple cells: The Zn–Cu salt-water pair gives ~0.7–0.8 V; the Al–Cu pair gives slightly more but is less predictable due to the oxide layer. Aluminium–copper produces a strikingly larger current per unit area than the zinc–carbon pair when the surface is freshly abraded.

Nickel–Iron Cell (Edison Battery)

Thomas Edison patented the nickel–iron (NiFe) cell in 1901 as a rugged, rechargeable alternative to the fragile lead–acid battery, intending it for electric vehicles and industrial use. Edison cells were famously durable — some original units from the early 1900s still function. The chemistry is straightforward: an iron anode, a nickel oxyhydroxide cathode, and potassium hydroxide as the electrolyte.

Materials

  • Iron electrode: steel wool (0000 grade, tightly packed) or iron nails / steel strip
  • Nickel electrode: strips cut from a nickel anode (available from electroplating suppliers), or nickel-plated steel — needs to be electrochemically “formed” (see below)
  • Potassium hydroxide (KOH) solution: 20–25% by weight (~280 g KOH per litre of water) — wear gloves and eye protection; KOH is strongly corrosive
  • Separator: cotton cloth, felt, or several layers of filter paper
  • Two small glass jars with lids, or a single partitioned container
  • Charger: a 1.5–2 V regulated supply (two AA batteries or a USB supply through a resistor), plus alligator clips
  • Multimeter

Forming the nickel electrode

Fresh nickel metal does not contain the active material (NiOOH); it must be formed by repeated charge–discharge cycles:

  1. Immerse the nickel strip in KOH solution alongside a steel counter-electrode
  2. Apply 1.6–1.8 V (nickel strip = positive terminal) for 30 minutes — oxygen evolves at the nickel surface and the surface oxidises to Ni(OH)₂, then to NiOOH
  3. Reverse polarity for 15 minutes to reduce it back to Ni(OH)₂
  4. Repeat steps 2–3 at least five times; the nickel surface gradually darkens and the active layer builds up

After forming, the nickel electrode is ready to act as a cathode.

Procedure

  1. Fill the container with KOH solution
  2. Suspend the steel-wool iron electrode (anode, negative) and the formed nickel electrode (cathode, positive) in the same electrolyte, separated by the cloth separator — the electrodes must not touch
  3. Connect the multimeter: red to nickel, black to iron; open-circuit voltage of a partially charged cell is ~1.1–1.2 V
  4. Charge: apply 1.6–1.8 V across the cell for 1–2 hours; hydrogen evolves at the iron electrode, oxygen at the nickel electrode
  5. After charging, read the resting voltage (~1.3–1.4 V); it will settle to ~1.2 V as the surface charge equilibrates
  6. Discharge through a 10–50 Ω resistor and monitor voltage over time; the discharge plateau is flat at ~1.1–1.2 V until the cell is nearly depleted

Reactions:

\[\ce{Fe + 2OH^- -> Fe(OH)2 + 2e^-} \quad \text{(anode, discharge)}\]

\[\ce{2NiOOH + 2H2O + 2e^- -> 2Ni(OH)2 + 2OH^-} \quad \text{(cathode, discharge)}\]

\[\ce{Fe + 2NiOOH + 2H2O <=> Fe(OH)2 + 2Ni(OH)2} \quad E \approx 1.2\,\text{V}\]

The double arrow indicates the reaction is fully reversible — applying voltage drives it backwards, regenerating NiOOH at the cathode and iron at the anode.

Why it is so durable: Neither electrode material dissolves into the electrolyte in the way zinc does in acid cells. The KOH electrolyte is not consumed by the reactions; it acts only as an ion carrier. The cell can be fully discharged, overcharged, or left for months without permanent damage — unlike lead–acid or lithium cells.

What limits it: Self-discharge is significant (~1–2% per day); the cell also produces hydrogen on every charge cycle. Energy density (~30 Wh/kg) is low by modern standards. The forming process takes patience, and NiOOH active material builds up slowly over many cycles — performance improves with each charge–discharge cycle for the first 20–30 cycles.

Warning

KOH is strongly corrosive — more so than NaOH at equivalent concentrations. Wear nitrile gloves and eye protection whenever handling the electrolyte. Hydrogen gas is produced at the iron electrode during charging; charge in a well-ventilated area and keep sparks and flames away.

Supercapacitors (EDLC)

An electric double-layer capacitor (EDLC) stores charge not by chemical reaction but by adsorbing ions onto a high-surface-area electrode. When voltage is applied, a layer of positive ions forms on the negative electrode surface and a layer of negative ions on the positive electrode — the “double layer.” No chemical bonds form or break, so the cycle can be repeated hundreds of thousands of times with negligible degradation.

How They Differ from Batteries

Property Battery EDLC
Energy storage Chemical (bulk reaction) Electrostatic (surface adsorption)
Charge/discharge speed Slow (minutes–hours) Fast (seconds)
Cycle life 500–2000 cycles typical 100,000+ cycles
Energy density High Low (10–100× less than Li-ion)
Power density Low Very high
Self-discharge Low Moderate–high

EDLCs are complementary to batteries, not replacements: they handle rapid bursts of power (regenerative braking, camera flash) while a battery handles sustained energy delivery.

Na₂SO₄ + Activated Carbon EDLC

Sodium sulfate solution is an excellent DIY electrolyte: it is non-toxic, neutral pH, has good ionic conductivity (~70 mS/cm at 1 mol/L), and does not corrode carbon or common current collectors. Activated carbon provides the enormous surface area needed — typically 1000–3000 m²/g.

Rated voltage: ~0.8–1.0 V per cell (above ~1.0 V, water begins to electrolyse)

Materials

  • Activated carbon — 5–10 g (aquarium-grade activated carbon works; crush finely)
  • Sodium sulfate — saturated aqueous solution (~28 g/100 mL at 25°C)
  • Carbon black or graphite powder (optional, ~10% by weight, improves conductivity)
  • PTFE binder or CMC (sodium carboxymethylcellulose) — a few drops of PTFE dispersion, or 1–2% CMC solution as binder
  • Current collectors: graphite sheet, carbon cloth, or stainless-steel mesh (avoid aluminium with Na₂SO₄ — it corrodes)
  • Separator: tissue paper, filter paper, or polyethylene film with small holes
  • Small clamp or bulldog clip to hold the stack together
  • Multimeter (to measure voltage and charge/discharge)

Making Activated Carbon from Charcoal (Optional)

Commercial activated carbon is ideal, but you can improve ordinary charcoal:

  1. Crush hardwood charcoal (not briquettes) to a fine powder
  2. Soak overnight in zinc chloride solution (20 g ZnCl₂ per 100 mL water)
  3. Filter, then heat in a covered tin at ~400°C for 30–45 minutes in a gas flame or oven — this activates the pores
  4. Rinse thoroughly with water until the washings are neutral, then dry
  5. Surface area will be lower than commercial activated carbon but sufficient for a demonstration

Electrode Preparation

  1. Mix activated carbon with graphite powder (9:1 ratio by weight) in a small bowl
  2. Add CMC solution drop by drop, stirring until you have a thick, paste-like consistency that holds together but is not runny
  3. Spread a thin, even layer (~1–2 mm) onto the current collector (graphite sheet or carbon cloth)
  4. Press firmly and allow to dry for several hours at room temperature or 60°C in an oven
  5. Cut two electrodes of identical size — 3×3 cm is a convenient starting point

Assembly

  1. Soak the separator (filter paper) in Na₂SO₄ solution
  2. Stack: current collector / electrode / separator / electrode / current collector
  3. Clamp the stack firmly (the pressure ensures good contact)
  4. Connect leads to the two outer current collectors

Layer diagram:

[+ lead] ── [current collector] ── [carbon electrode]
                                         │
                                   [separator, wet with Na₂SO₄]
                                         │
              [current collector] ── [carbon electrode]
[− lead] ──────────────────────────────

Charging and Discharging

  1. Charge: Apply 0.8–0.9 V across the terminals using a USB power supply through a 100 Ω series resistor, or two AA batteries in series with a voltage divider. Charge for 30–60 seconds
  2. Disconnect the power source and immediately measure open-circuit voltage — it should hold near the charge voltage
  3. Discharge through a load resistor (100–1000 Ω) and monitor the voltage drop over time with a multimeter
  4. Plot voltage vs time: the roughly linear discharge is characteristic of a capacitor (contrast with a battery, which holds a flat plateau then drops sharply)

Estimating capacitance:

\[C = I \cdot \frac{\Delta t}{\Delta V}\]

Where \(I\) is the discharge current (A), \(\Delta t\) is the time (s) for the voltage to drop \(\Delta V\) volts. A 3×3 cm cell with good activated carbon might give 0.5–2 F.

Stacking Cells in Series

To reach a usable voltage (e.g., to light an LED):

  • Each cell is rated ~0.8 V
  • Wire 3 cells in series for ~2.4 V, enough to drive a red LED
  • Balance the cells by ensuring each charges to the same voltage (use matched electrode area and weight)

Improving Performance

Modification Effect
Thinner electrodes Lower internal resistance, faster charge/discharge
Higher surface-area carbon More capacitance
Add 5% carbon black Better conductivity within the electrode
Higher Na₂SO₄ concentration Lower electrolyte resistance
Apply stack pressure Better electrode–separator contact

Comparing What You Can Build

Cell type Voltage Current Rechargeable Difficulty
Lemon cell ~0.9 V ~1 mA No Very easy
Zinc–carbon (salt water) ~0.7 V ~5 mA No Easy
Daniell cell ~1.1 V ~10 mA No Easy
Aluminium–air ~1.2 V ~50 mA/cm² No (anode consumed) Medium
Magnesium–copper ~1.6 V ~10 mA No Easy
Copper–aluminium (salt water) ~0.9 V ~10 mA No Easy
Nickel–iron (Edison) ~1.2 V ~20 mA Yes (100 000+ cycles) Hard
Na₂SO₄/carbon EDLC ~0.8 V/cell High burst Yes (many times) Medium

Safety

Note

Most of these experiments use mild, low-toxicity chemicals. The main exceptions:

  • Sodium hydroxide (aluminium–air cell): Corrosive. Wear nitrile gloves and eye protection. Neutralise spills with dilute vinegar before cleaning up.
  • Potassium hydroxide (nickel–iron cell): Strongly corrosive — more aggressive than NaOH at equivalent concentrations. Wear nitrile gloves and eye protection; neutralise spills with dilute vinegar.
  • Copper sulfate: Toxic to aquatic organisms. Dispose of as hazardous waste, not down the drain.
  • Zinc chloride (if used in carbon activation): Corrosive. Handle with gloves; rinse electrodes thoroughly before use in a supercapacitor.
  • Hydrogen gas: Produced at the cathode of any aqueous electrochemical cell under sufficient voltage. Work in a ventilated area and avoid sparks near the apparatus.

Further Experiments

  • Internal resistance: Measure open-circuit voltage (\(V_\text{OC}\)) then voltage under a known load (\(V_L\), \(R_L\)). Internal resistance \(r = R_L \cdot (V_\text{OC}/V_L - 1)\).
  • Temperature effects: Chill a lemon cell in ice water vs warm water. Ionic conductivity decreases with temperature; observe the voltage drop.
  • Electrolyte concentration: Compare cell voltage and internal resistance for the Daniell cell at 0.1, 0.5, and 1.0 mol/L CuSO₄.
  • EDLC self-discharge: Charge the supercapacitor fully, disconnect, and measure voltage every 5 minutes for an hour. Plot the self-discharge curve.
  • Ragone plot: Measure energy density (charge × average voltage) and power density (peak current × voltage) for a battery and an EDLC and plot them on log–log axes — this is how engineers compare energy storage technologies.