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A battery is a device that stores chemical energy, and converts it to electricity. This is known as electrochemistry and the system that underpins a battery is called an electrochemical cell. A battery can be made up of one or several (like in Volta's original pile) electrochemical cells. Each electrochemical cell consists of two electrodes (anode and cathode) separated by an electrolyte.

Batteries come in all shapes and sizes to suit various needs. They range from miniature batteries used in wristwatches and hearing aids to several meters wide, serving as an emergency power resource or storing renewable energy from solar power plants and wind farms.

Benjamin Franklin coined the word “battery” to refer to linked capacitors in his experiments, but it was Alexandro Volta who developed the first “real” battery, in 1800. Since then the usefulness and ubiquity of these power cells has increased steadily.

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The galvanic cell has a long history. Its name comes from Italian physicist Luigi Galvani (1737 – 1798), who observed that dissected frog leg muscles twitched when an electric current was applied to them. A galvanic cell uses the energy released from a spontaneous redox reaction to produce an electric current.

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Another common name for galvanic cells is voltaic cells, which is named after another Italian physicist, Alessandro Volta (1745 – 1827), who invented the galvanic (voltaic) cell. Volta invented the “Voltaic pile”. He made the electrodes in the form of discs, and put a put a cardboard or felt disc soaked with electrolyte in between each copper-zinc pair – and then, stacked them up into a pile.

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A battery operates by converting chemical energy into electrical energy through electrochemical reactions. It comprises two electrodes—an anode (negative) and a cathode (positive)—separated by an electrolyte. When the battery is connected to an external circuit, oxidation occurs at the anode, releasing electrons, while reduction takes place at the cathode, accepting electrons. This flow of electrons through the external circuit generates electric current, powering connected devices. Simultaneously, ions move through the electrolyte to balance the charge. In rechargeable batteries, applying an external electrical current reverses these chemical reactions, restoring the original composition of the electrodes and allowing the battery to be used multiple times.

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For easy understanding of the redox reaction happening in the batteries.

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Primary batteries are non-rechargeable and designed for single-use applications. Once their chemical energy is depleted, they cannot be recharged and must be replaced. These batteries are commonly used in devices where recharging is impractical or unnecessary, such as remote controls, flashlights, and clocks. They are known for their high energy density, long shelf life, and cost-effectiveness in low-consumption applications.

Characteristics of Primary Batteries:

Non-rechargeable: Irreversible chemical reactions.

High energy density: Provides long-lasting power for low-drain devices.

Long shelf life: Retains charge for extended periods.

Cost-effective: Ideal for low-consumption applications.

Examples of Primary Batteries:

Alkaline batteries (used in remote controls, flashlights).

Zinc-carbon batteries (used in clocks, basic toys).

Lithium batteries (used in smoke detectors, cameras).

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Secondary batteries are rechargeable, designed for multiple cycles of charge and discharge. They can be restored to their original charge state by applying an external electric current. These batteries are widely used in devices that require frequent and long-term use.

Characteristics of Secondary Batteries:

Rechargeable: Reversible chemical reactions.

Higher upfront cost: But economical over time due to reusability.

Energy-efficient: Suitable for high-drain and long-term applications.

Shorter shelf life: Compared to primary batteries.

Examples of Secondary Batteries:

Lead-acid batteries (used in vehicles, backup power systems).

Lithium-ion (Li-ion) batteries (used in smartphones, laptops, electric vehicles).

Nickel-cadmium (NiCd) batteries (used in power tools, medical equipment).

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Difference between Primary and Secondary Battery

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Difference between Primary and Secondary Battery

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Primary batteries are non-rechargeable and designed for single use, with irreversible chemical reactions, commonly used in low-drain devices like remote controls and flashlights. They have a lower upfront cost and longer shelf life but generate more waste. Secondary batteries, in contrast, are rechargeable, allowing for multiple use cycles. While they have a higher initial cost, they are more economical in the long run and suitable for high-drain devices like smartphones and electric vehicles. Secondary batteries are more environmentally friendly due to their reusability. Examples of primary batteries include alkaline and lithium, while secondary batteries include lithium-ion and lead-acid.

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Primary vs. Secondary Cells (Batteries) - IB Physics

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References

11.4.1.1: The Voltaic Cell. (2022, June 20). Engineering LibreTexts. https://eng.libretexts.org/Sandboxes/jhalpern/Energy_Alternatives/11%3A_Energy_Storage/11.04%3A_Energy_Storage_Technologies/11.4.01%3A_Battery_Storage/11.4.1.01%3A_The_Voltaic_Cell

Australian Academy of Science. (2016, February 24). How a battery works. Australian Academy of Science. https://www.science.org.au/curious/technology-future/batteries

Bramer, B. (2021, January 31). Galvanic Cells & Voltaic Cells | Electrochemical Cells. ChemTalk. https://chemistrytalk.org/electrochemical-galvanic-cells/

Chapman, B. (2019, September 23). How does a lithium-Ion battery work? Let’s Talk Science. https://letstalkscience.ca/educational-resources/stem-explained/how-does-a-lithium-ion-battery-work

Rouse, M. (2019). What is a Battery? - Definition from Techopedia. Techopedia.com. https://www.techopedia.com/definition/16316/battery

Stubbings, J. (2019). Lead-Acid Batteries Chemistry Tutorial. Ausetute.com.au. https://www.ausetute.com.au/pbbattery.html

University of Cambridge. (2025a). Primary batteries. Doitpoms.ac.uk. https://www.doitpoms.ac.uk/tlplib/batteries/primary.php?

University of Cambridge. (2025b). Secondary batteries. Doitpoms.ac.uk. https://www.doitpoms.ac.uk/tlplib/batteries/secondary.php?

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Alkaline Batteries for remote controls, flashlights, etc.

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Lithium Batteries in smoke detectors, cameras, etc.

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Lead-Acid Batteries for starting the engine, powering electrical systems, voltage stabilization, energy storage in cars.

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Lithium-ion Batteries for powering the device, enabling portability, quick charging, maintaining charge efficiency in cellphones.

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Alkaline Batteries

Zinc (Oxidizes)

Zn(s)+2OH−(aq)→Zn(OH)2​(s)+2e−

Manganese Dioxide (Reduced)

2MnO2​(s)+2e−+2H2​O(l)→Mn2​O3​(s)+2OH−(aq)

Overall Reaction

Zn(s)+2MnO2​(s)+2H2​O(l)→Zn(OH)2​(s)+Mn2​O3​(s)

Zinc Carbon Batteries

Zinc (Oxidizes)

Zn(s) → Zn²⁺(aq) + 2 e⁻

Manganese Dioxide (Reduced)

2MnO₂(s) + 2 H⁺(aq) + 2 e⁻ → Mn₂O₃(s) + H₂O(l)

The H⁺ comes from the NH₄⁺(aq):

NH₄⁺(aq) + H₂O(l) → H⁺(aq) + NH₃(aq)

Overall Reaction

Zn(s) + 2 MnO₂(s) + 2 NH₄⁺(aq) → Mn₂O₃(s) + Zn(NH₃)₂ ²⁺(aq)

Lithium Batteries

Lithium (Oxidized)

Li(s)→Li+(aq)+e−

Cobalt Oxide (Reduced)

Lix​CoO2​+xe−→LiCoO2​

Overall Reaction

Li(s)+Lix​CoO2​(s)→LiCoO2​(s)+Lix​O2

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Lead-Acid Battery

Lead (Oxidized)

Pb(s)+ SO₄ ²⁻(aq) → PbSO4(s)+ 2e⁻

Lead Dioxide (Reduced)

PbO2(s) + 4H⁺(aq) + 2e⁻ + SO₄ ²⁻(aq) → PbSO4(s) + 2H₂O(l)

Overall Reaction

Pb(s) + PbO2(s) + 4H⁺(aq) + 2SO₄ ²⁻(aq)→2PbSO4(s) + 2H₂O(l)

Lithium-ion Battery

Graphite (Oxidized)

LiC6​→Li+(aq)+e−+C6​

Lithium Cobalt (Reduced)

LiCoO2​+Li+(aq)+e−→Li2​CoO2​

Overall Reaction

LiC6​+LiCoO2​→C6​+Li2​CoO2​

Nickel-cadmium Battery

Cadmium (Oxidized)

Cd(s)+2OH−(aq)→Cd(OH)2​(s)+2e−

Nickel Oxyhydroxide (Reduced)

2NiOOH(s)+2e−+2H2​O(l)→2Ni(OH)2​(s)+2OH−(aq)

Overall Reaction

Cd(s)+2NiOOH(s)+2H2​O(l)→Cd(OH)2​(s)+2Ni(OH)2​(s)