Lesson Sources and induction · Conduction and cells
Cells and batteries
A cell converts chemical energy into electrical energy using two unlike electrodes and an electrolyte. A battery is two or more cells connected together (though people often say “battery” for a single AA). Primary cells are use-once; secondary cells recharge. Cell voltage comes from the metal pair; current capacity comes mainly from plate area and chemistry — not from “bigger voltage alone.”
Chemical action produces voltage
Early experiments showed that dissimilar metals in an electrolyte produce electricity — not from “animal muscle,” but from chemical reaction at the electrodes. Volta’s voltaic pile stacked zinc and silver discs separated by brine-soaked cardboard; each layer is one cell; voltage adds when cells stack in series.
Any practical cell needs:
- Two unlike electrode materials
- An electrolyte (liquid, paste, or gel)
- A completed external circuit to deliver power
Simple demos: potato with copper and aluminum wires; wet paper between a penny and a nickel — small voltage measurable with a high-impedance meter.
Cell vs battery — words that matter on the job
| Term | Meaning |
|---|---|
| Cell | One unit: one set of electrodes + electrolyte |
| Battery | Two or more cells combined (e.g., 6 × 1.5 V = 9 V rectangular pack) |
Schematic: long/short lines = one cell; multiple pairs = battery symbol. Label and order parts by cell count when sizing chargers or replacement packs.
What sets cell voltage: electromotive series
Voltage depends on how eagerly one metal gives electrons compared with the other. The electromotive series ranks metals (lithium and potassium at the top — strongest electron donors; gold and platinum at the bottom).
Rule of thumb: farther apart on the list → higher cell voltage. In a zinc–copper acid cell, zinc (higher on the list) is the negative electrode; copper is positive. Zinc corrodes as it supplies electrons; hydrogen bubbles at the copper side in the classic wet cell.
Not every metal pair is practical — some corrode instantly or build insulating films.
Primary vs secondary cells
| Type | Can recharge? | What happens when discharged |
|---|---|---|
| Primary | No — replace | Electrode material consumed (zinc can dissolves) |
| Secondary | Yes | Chemical reaction reversible with charging current |
Common primary voltages per cell (approximate):
| Chemistry | Volts per cell |
|---|---|
| Carbon-zinc / alkaline | 1.5 |
| Mercury button | 1.35 |
| Silver-zinc button | 1.6 |
| Zinc-air | ~1.4 |
| Lithium (varies) | 1.9 – 3.6 |
When a primary cell is “dead,” chemistry is spent — charging will not restore it.
Primary cell families you will see
Carbon-zinc (Leclanché): zinc can (negative), carbon rod (positive), paste electrolyte — “dry cell.” Zinc container erodes until it leaks; discard promptly when spent.
Alkaline: same 1.5 V nominal, potassium hydroxide paste, longer life and better high-rate performance, higher cost. Generally treated as non-rechargeable despite specialty rechargeables with limited cycles.
Button cells: watches, hearing aids, calculators — mercury, silver, lithium chemistries trade cost for energy density and stable voltage.
Lithium primary: very high energy density; some types last decades at microamp loads; wrong charge current on rechargeable lithium types can be explosive — match charger to chemistry.
Current capacity and internal resistance
Current capacity = how much energy the cell can deliver over time. Ratings include:
- mAh (milliampere-hours): 10,000 mAh ≈ 10 mA for 1000 h or 100 mA for 100 h (same chemistry, idealized)
- Wh (watt-hours): mAh × terminal voltage
Larger D vs AA vs AAA = more plate area → more current, not higher voltage per cell.
Every cell has internal resistance. Under load, terminal voltage sags. Short-circuit a fresh alkaline D: ammeter might show several amperes but voltage collapses — most energy heats the cell, not the load. As cells age, internal resistance rises and terminal voltage under load drops even when “full” by open-circuit check.
Secondary preview — same box, different life cycle
Lead-acid, nickel-cadmium, lithium-ion, and others are secondary — discharge converts plates/compounds; charging reverses the reaction. Lead-acid (~2 V per cell), NiCd/NiMH (~1.2 V), Li-ion (~3.6 V) each have distinct charging rules. This lesson names the split; following lessons cover lead-acid, connections, and alternate small sources in depth.
Field case
Situation. Night shift replaces “dead” 9 V smoke-detector batteries with cheap carbon-zinc cells from the lunch-room remote. Within two months every detector chirps low-battery again.
Why. Same 1.5 V × 6 cells = 9 V geometry, but alkaline packs hold several times the mAh of carbon-zinc at smoke-detector drain rates. Open-circuit voltage looked fine on install; under pulsed load, sag came early.
Applied lesson. Match chemistry and capacity to duty cycle, not just terminal voltage. For safety gear, use manufacturer-specified type (usually alkaline or lithium primary).
In the field
Symptom
Device dies quickly or works only briefly after battery change
Where to look
Cell count, chemistry label, date code, corrosion at contacts
Likely causes
- Wrong primary type, mixed old/new cells, high internal resistance, leakage from overdue primary
What to measure
- Loaded vs open-circuit voltage
- compare mAh/Wh rating to OEM
What not to do
- Recharge primary cells
- mix chemistries in one battery holder
Checklist
- I distinguish cell from battery
- I explain primary vs secondary in plain language
- I use the electromotive series idea (farther apart → more voltage)
- I cite common per-cell voltages (1.5, 1.2, 2, 3.6)
- I define mAh and why plate area affects current, not voltage
- I explain terminal sag from internal resistance