Modern examples of battery technology: electrochemical cell examples from phones to EVs

The best way to understand electrochemistry is to start with batteries you actually touch. Several classic and modern cells show up again and again in textbooks because they’re also workhorses of industry. These are the real examples of battery technology: electrochemical cell examples that power consumer electronics, vehicles, and backup systems.

Think about a typical day:

• Your phone and laptop run on lithium‑ion cells.
• Your car (even if it’s not electric) uses a lead–acid starter battery.
• Remote controls and flashlights rely on alkaline cells.
• Grid‑scale battery farms increasingly use lithium‑ion or vanadium redox flow cells.

All of these are just different electrochemical designs that exploit specific redox couples and cell potentials.

Alkaline batteries: the everyday example of a primary electrochemical cell

Alkaline batteries are probably the most familiar example of battery technology: electrochemical cell examples don’t get more common than the AA or AAA cells in a game controller.

Typical overall cell reaction (simplified):

• Anode (oxidation):
  \[ \text{Zn(s)} + 2\text{OH}^- (aq) \rightarrow \text{ZnO(s)} + \text{H}_2\text{O(l)} + 2e^- \]

• Cathode (reduction):
  \[ 2\text{MnO}_2(s) + \text{H}_2\text{O(l)} + 2e^- \rightarrow \text{Mn}_2\text{O}_3(s) + 2\text{OH}^- (aq) \]

• Approximate cell potential: about 1.5 V under near‑standard conditions.

Alkaline cells are primary batteries: they’re designed for one‑time use. The electrochemical reactions are not easily reversed without damaging the structure of the electrodes. That’s why you toss them when they’re “dead.”

From a cell‑potential standpoint, the 1.5 V rating comes from the difference in standard reduction potentials between the zinc and manganese dioxide half‑reactions. In practice, the voltage drops under load because of internal resistance and concentration changes, which you can model with the Nernst equation.

Lead–acid batteries: classic automotive example of battery technology

Lead–acid batteries are among the oldest large‑scale examples of battery technology: electrochemical cell examples that still matter in 2024. They crank internal‑combustion engines, stabilize power in data centers, and sit in uninterruptible power supplies.

Half‑reactions in a typical 12 V car battery (each cell ~2.0 V):

• Anode (oxidation):
  \[ \text{Pb(s)} + \text{HSO}_4^- (aq) \rightarrow \text{PbSO}_4(s) + \text{H}^+ (aq) + 2e^- \]

• Cathode (reduction):
  \[ \text{PbO}_2(s) + 3\text{H}^+ (aq) + \text{HSO}_4^- (aq) + 2e^- \rightarrow \text{PbSO}_4(s) + 2\text{H}_2\text{O(l)} \]

• Overall cell reaction:
  \[ \text{Pb(s)} + \text{PbO}_2(s) + 2\text{H}_2\text{SO}_4(aq) \rightarrow 2\text{PbSO}_4(s) + 2\text{H}_2\text{O(l)} \]

• Standard cell potential: about 2.04 V per cell, so a 6‑cell battery is around 12.2–12.6 V at rest.

Lead–acid chemistry is a good example of how cell potential links to real‑world diagnostics. Mechanics routinely infer state of charge from open‑circuit voltage: around 12.6 V is fully charged, while below ~12.0 V signals significant discharge. That behavior tracks the Nernst equation because the activities of sulfuric acid and lead sulfate shift as the battery discharges.

Despite newer chemistries, lead–acid remains attractive for starter batteries because it delivers very high surge currents, tolerates abuse, and is highly recycled in the U.S. (recycling rates above 95% according to the U.S. EPA: https://www.epa.gov).

Lithium‑ion cells: flagship examples of modern battery technology

Lithium‑ion cells are the best examples of battery technology: electrochemical cell examples that dominate phones, laptops, power tools, and electric vehicles (EVs). They’re secondary (rechargeable) cells based on reversible intercalation of lithium ions between host materials.

A common smartphone chemistry uses a graphite anode and a lithium cobalt oxide (LiCoO₂) cathode.

Simplified half‑reactions during discharge:

• Anode (oxidation):
  \[ \text{LiC}_6 \rightarrow \text{C}_6 + \text{Li}^+ + e^- \]

• Cathode (reduction):
  \[ \text{Li}_{1-x}\text{CoO}_2 + x\text{Li}^+ + xe^- \rightarrow \text{LiCoO}_2 \]

• Nominal cell potential: about 3.6–3.7 V, depending on cathode chemistry.

Modern EVs often use:

• NMC (LiNiMnCoO₂) or NCA (LiNiCoAlO₂) cathodes for high energy density.
• LFP (LiFePO₄) cathodes for improved thermal stability and long cycle life.

The actual voltage curve versus state of charge is not flat; it reflects the thermodynamics of lithium intercalation and can be modeled with modified Nernst‑type relationships. That’s why a “3.7 V” lithium‑ion cell might charge up to 4.2 V and cut off around 3.0 V under normal use.

According to the U.S. Department of Energy (DOE: https://www.energy.gov), advanced lithium‑ion packs now reach 250–300 Wh/kg at the cell level in commercial EVs, with research pushing beyond 350 Wh/kg. Those numbers are a direct payoff from carefully engineered electrochemical cell potentials and minimized internal resistance.

Nickel–metal hydride (NiMH): a practical example of safer, moderate‑energy cells

Before lithium‑ion took over, NiMH cells were the go‑to rechargeable option for AA and AAA sizes and early hybrid vehicles like the Toyota Prius. They’re still widely used where safety and cost matter more than maximum energy density.

Typical NiMH cell reactions (simplified):

• Anode (oxidation): metal hydride alloy (MH) releases hydrogen:
  \[ \text{MH} + \text{OH}^- \rightarrow \text{M} + \text{H}_2\text{O} + e^- \]

• Cathode (reduction): nickel oxyhydroxide to nickel hydroxide:
  \[ \text{NiO(OH)} + \text{H}_2\text{O} + e^- \rightarrow \text{Ni(OH)}_2 + \text{OH}^- \]

• Nominal cell potential: about 1.2 V.

NiMH is a good example of how lower cell potential doesn’t necessarily mean poor performance. These cells tolerate deep discharge, handle many cycles, and avoid some of the thermal runaway risks associated with certain lithium‑ion combinations. For consumer devices where weight is less critical, NiMH remains a solid example of battery technology: electrochemical cell examples that balance safety and performance.

Zinc–carbon and zinc–air: low‑cost and high‑energy primary cell examples

Zinc‑based cells give two very different real‑world examples:

• Zinc–carbon cells: older, low‑cost primary batteries with ammonium chloride or zinc chloride electrolytes. They’re being displaced by alkaline cells but still appear in low‑drain devices.
• Zinc–air cells: used in hearing aids and some niche applications where high energy density and stable voltage matter.

For zinc–air, oxygen from the air acts as the cathode reactant.

Zinc–air half‑reactions (simplified):

• Anode (oxidation):
  \[ \text{Zn(s)} + 2\text{OH}^- \rightarrow \text{ZnO(s)} + \text{H}_2\text{O(l)} + 2e^- \]

• Cathode (reduction):
  \[ \text{O}_2(g) + 2\text{H}_2\text{O(l)} + 4e^- \rightarrow 4\text{OH}^- \]

• Cell potential: around 1.4 V.

Because one reactant is atmospheric oxygen, zinc–air cells pack a lot of energy in a small volume. That’s a textbook example of using thermodynamically favorable redox couples to boost cell potential and energy density.

Vanadium redox flow batteries: grid‑scale example of large electrochemical cells

When you scale up from a phone to a wind farm, you need different examples of battery technology: electrochemical cell examples that can store megawatt‑hours, not watt‑hours. That’s where redox flow batteries come in, particularly vanadium redox flow batteries (VRFBs).

In a VRFB, both half‑cell electrolytes are vanadium in different oxidation states, dissolved in sulfuric acid. The electrolytes are stored in tanks and pumped through electrochemical cells.

Typical half‑reactions:

• Positive electrode (catholyte during discharge):
  \[ \text{VO}_2^+ + 2\text{H}^+ + e^- \rightarrow \text{VO}^{2+} + \text{H}_2\text{O} \]

• Negative electrode (anolyte during discharge):
  \[ \text{V}^{2+} \rightarrow \text{V}^{3+} + e^- \]

• Cell potential: around 1.3–1.4 V.

The beauty of flow batteries is that power depends on stack size (number and area of electrochemical cells), while energy capacity depends on tank volume. That separation is a different style of electrochemical design, but the same thermodynamic rules apply: the cell potential comes from the difference in reduction potentials of the vanadium couples, adjusted by concentration via the Nernst equation.

The U.S. DOE’s Office of Electricity (https://www.energy.gov/oe) highlights redox flow systems as promising for long‑duration storage supporting solar and wind integration.

Solid‑state and lithium‑metal cells: cutting‑edge examples under development

Looking toward 2025 and beyond, some of the most hyped examples of battery technology: electrochemical cell examples involve solid‑state and lithium‑metal anodes.

The idea is to replace flammable liquid electrolytes with solid ion conductors and pair them with high‑capacity lithium‑metal anodes.

Conceptual half‑reactions for a lithium‑metal solid‑state cell:

• Anode (oxidation):
  \[ \text{Li(s)} \rightarrow \text{Li}^+ + e^- \]

• Cathode (reduction): similar to current lithium‑ion cathodes, for example:
  \[ \text{Li}_{1-x}\text{NiMnCoO}_2 + x\text{Li}^+ + xe^- \rightarrow \text{LiNiMnCoO}_2 \]

Because lithium metal has a very low standard reduction potential (~−3.04 V vs SHE), pairing it with high‑voltage cathodes can push cell potentials higher than today’s typical 3.6–3.7 V lithium‑ion cells. That means more energy per kilogram, at least in theory.

Research from universities and labs (for example, work summarized by the U.S. National Renewable Energy Laboratory, NREL: https://www.nrel.gov) shows progress in solid electrolytes like sulfides, oxides, and polymers, but challenges remain: dendrite formation, interface resistance, and manufacturing cost.

These are still pre‑mass‑market, but they are strong candidates for future best examples of battery technology, especially in EVs and high‑end portable electronics.

Connecting cell potential calculations to real battery performance

All of these examples of battery technology: electrochemical cell examples are governed by the same core formulas you see in an electrochemistry course.

Standard cell potential and the Nernst equation

For a generic cell:

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

This standard cell potential assumes 1 M concentrations, 1 atm gases, and 25 °C. Real batteries almost never operate under those conditions, which is where the Nernst equation comes in:

\[ E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.0592}{n} \log Q \]

Here:

• \( n \) is the number of electrons transferred in the balanced cell reaction.
• \( Q \) is the reaction quotient based on actual activities (often approximated by concentrations or partial pressures).

For an alkaline AA cell, as zinc is consumed and manganese dioxide is reduced, the concentrations of reactants and products shift, changing \( Q \) and nudging the cell potential downward. That’s why the voltage of a “fresh” alkaline cell under load is closer to 1.4–1.5 V and gradually sinks as it discharges.

In lithium‑ion packs, battery management systems (BMS) use empirical voltage‑versus‑state‑of‑charge curves that are underpinned by this same thermodynamic story, tuned by kinetic effects and internal resistance.

Gibbs free energy and energy content

The link between electrochemical cell potential and stored energy is given by:

\[ \Delta G = -n F E_{\text{cell}} \]

where \( F \) is Faraday’s constant (~96,485 C/mol·e⁻).

Higher cell potentials and more electrons transferred per mole translate to more negative \( \Delta G \), meaning more usable electrical work per unit of reactants. That’s why lithium‑ion, with cell potentials around 3.6–4.2 V, delivers much higher energy density than a 1.2 V NiMH cell of similar size.

Why different applications favor different electrochemical cell examples

When engineers choose among these examples of battery technology: electrochemical cell examples, they’re trading off several factors:

• Energy density: lithium‑ion and zinc–air stand out.
• Power density: lead–acid and some lithium‑ion chemistries can deliver huge bursts.
• Cycle life: LFP lithium‑ion and vanadium flow batteries excel.
• Safety: NiMH and well‑managed LFP systems are attractive.
• Cost and recyclability: lead–acid is still hard to beat on both.

That’s why:

• EVs lean toward lithium‑ion (NMC, NCA, LFP), with experimental solid‑state designs on the horizon.
• Grid storage increasingly uses lithium‑ion now, but vanadium redox flow and other flow chemistries are gaining attention for long‑duration storage.
• Medical implants (like pacemakers) use high‑reliability primary cells such as lithium–iodine or lithium–silver vanadium oxide, where predictable cell potential and low self‑discharge matter more than raw energy density. For background on implantable device power needs, see overviews from the U.S. National Institutes of Health (NIH: https://www.nih.gov).

Each of these is an example of how the same electrochemical principles show up in very different hardware.

FAQ: common questions about electrochemical cell examples in batteries

What are some real‑world examples of battery technology: electrochemical cell examples?

Real‑world examples include alkaline AA and AAA cells, lead–acid car batteries, lithium‑ion phone and EV packs, NiMH rechargeables, zinc–air hearing‑aid batteries, and vanadium redox flow batteries for grid storage.

Can you give an example of how to calculate cell potential for a battery?

Take a simplified galvanic cell based on Zn²⁺/Zn and Cu²⁺/Cu:

• \( E^\circ_{\text{cathode}} \) (Cu²⁺/Cu) = +0.34 V
• \( E^\circ_{\text{anode}} \) (Zn²⁺/Zn) = −0.76 V

Then:
\[ E^\circ_{\text{cell}} = 0.34\,\text{V} - (−0.76\,\text{V}) = 1.10\,\text{V} \]

This classic classroom example isn’t a commercial battery, but the method is exactly the same for practical electrochemical cell examples.

Which chemistry is the best example of a high‑energy battery today?

As of 2024–2025, high‑nickel lithium‑ion chemistries (like NMC and NCA) are among the best examples for high energy density in commercial products, reaching 250–300 Wh/kg at the cell level. Experimental lithium‑metal and solid‑state cells may eventually surpass them, but they’re still in development.

Why do some examples of batteries use liquid electrolytes while others use solids?

Liquid electrolytes are easy to manufacture and often have high ionic conductivity, which helps power output. Solid electrolytes can improve safety and enable lithium‑metal anodes, but they’re harder to engineer with low resistance and stable interfaces. Both types are legitimate examples of battery technology: electrochemical cell examples optimized for different design goals.

Are there examples of batteries that can be fully recycled?

Lead–acid batteries are a strong example of a highly recycled system, with U.S. recycling rates above 95%. Lithium‑ion recycling is improving, with growing efforts supported by national labs and industry. The electrochemistry doesn’t change during recycling, but the materials are recovered and fed back into new electrochemical cell examples.

The bottom line: once you understand how cell potential, reaction stoichiometry, and the Nernst equation fit together, all of these examples of battery technology—electrochemical cell examples from AA alkalines to grid‑scale flow batteries—start to look like variations on the same theme: controlled redox reactions engineered to deliver electrons where and when we want them.