Electrochemical Cells

An electrochemical cell is a contained system where an oxidation-reduction (redox) reaction takes place. Every electrochemical cell falls into one of three categories — galvanic (voltaic) cells, electrolytic cells, or concentration cells — and the same underlying rules for electron flow, oxidation, and reduction apply across all three.

Key Takeaways

• An electrochemical cell is a contained system where a redox reaction occurs; the three types are galvanic (voltaic), electrolytic, and concentration cells.

• Galvanic and concentration cells run spontaneous reactions (positive emf, ΔG < 0); electrolytic cells run nonspontaneous reactions and need an external power source (negative emf, ΔG > 0).

• In every cell type, electrons flow anode → cathode (oxidation at the anode, reduction at the cathode), while conventional current flows cathode → anode.

• The Daniell cell (Zn anode, Cu cathode, salt bridge, ~1.1 V) is the standard galvanic cell example; its cell diagram, Zn(s) | Zn²⁺(1 M) || Cu²⁺(1 M) | Cu(s), uses a single line for a phase boundary and a double line for a salt bridge.

• Electrolytic cells flip electrode polarity relative to galvanic cells (anode positive, cathode negative) and require an external power source, as in the electrolysis of molten NaCl to produce Na and Cl₂.

• Faraday's laws relate charge passed to substance liberated; the Faraday constant (F ≈ 96,485 C/mol e⁻) lets you calculate moles (It/nF) and mass (It·M/nF) deposited during electrolysis.

• Concentration cells use identical electrodes and are driven by a concentration gradient between half-cells; the neuron cell membrane and its resting membrane potential are a biological example.

• Rechargeable batteries like lead-acid (Pb/PbO₂/H₂SO₄) and nickel-cadmium (Cd/NiO(OH)/KOH) function as both galvanic and electrolytic cells depending on charge state.

What Is an Electrochemical Cell?

Galvanic cells and concentration cells both run on spontaneous redox reactions: the reaction proceeds on its own, without any external energy input, and produces a positive electromotive force (emf) — the voltage, or electrical potential difference, generated by the cell. A positive emf means the cell releases energy (ΔG < 0). Batteries in flashlights and other everyday devices are common galvanic cells.

Concentration cells are a special type of galvanic cell: the two electrodes are chemically identical, but the electrolyte concentrations surrounding them differ. That concentration difference — not a difference in electrode material — is what drives electron flow and generates current. Like galvanic cells, concentration cells produce a positive emf and involve spontaneous reactions.

Electrolytic cells, by contrast, run on nonspontaneous reactions. They require an external power source to force the reaction forward, and they have a negative emf, meaning they absorb energy (ΔG > 0). The electrolysis of water — splitting it into hydrogen and oxygen gas using electrical energy — is a classic electrolytic cell.

Temperature also affects how well a cell performs: lead-acid car batteries, a type of galvanic cell, are notorious for failing in cold weather, because low temperatures slow the chemical reactions inside the battery and reduce its ability to generate current.

Anode, Cathode, and the Direction of Electron Flow

In every electrochemical cell — galvanic, electrolytic, or concentration — electrons flow from the anode to the cathode. The anode is where oxidation (loss of electrons) happens; the cathode is where reduction (gain of electrons) happens.

This is where a lot of students get tripped up: while electrons physically flow from anode to cathode, conventional current is defined as flowing the opposite way, from cathode to anode. That's because current, by physics convention, tracks the flow of positive charge — which moves opposite to electron flow. So:

• Electron flow: anode → cathode

• Conventional current: cathode → anode

In a galvanic cell, the anode is the negatively charged electrode and the cathode is the positively charged electrode, and electrons flow naturally from anode to cathode because the redox reaction is spontaneous.

Galvanic Cells: The Daniell Cell

The Daniell cell is the textbook example of a galvanic cell — it generates electrical energy from a spontaneous redox reaction between zinc and copper.

• The anode is solid zinc, sitting in a zinc sulfate (ZnSO₄) solution. Zinc metal is oxidized, losing electrons to form zinc ions: Zn → Zn²⁺ + 2e⁻. Those electrons travel out into the external circuit.

• The cathode is solid copper, sitting in a copper sulfate (CuSO₄) solution. Copper ions in solution gain those electrons and are reduced to copper metal: Cu²⁺ + 2e⁻ → Cu.

• A salt bridge — typically filled with a salt like potassium nitrate (KNO₃) that won't react with the cell components — connects the two half-cells. It maintains electrical neutrality by letting ions flow between the compartments: negative ions move toward the anode side, positive ions move toward the cathode side, without letting the two solutions mix directly. Without it, charge would build up in each half-cell and the reaction would stall.

As zinc oxidizes at the anode, the released electrons travel through the external wire to the copper cathode (where they can power a device along the way), and copper ions are reduced and deposit as solid copper. Electrons flow anode → cathode; conventional current flows cathode → anode. The whole process generates an emf of about 1.1 V, consistent with a spontaneous reaction (ΔG < 0). Adding the two half-reactions together gives the net overall reaction for the cell — the actual voltage calculation is covered separately, once reduction potentials are introduced.

Reading a Cell Diagram

A cell diagram (or cell notation) is shorthand for describing an electrochemical cell's components and layout without drawing the whole setup. For the Daniell cell, it's written as:

Zn(s) | Zn²⁺(1 M) || Cu²⁺(1 M) | Cu(s)

The rules for building one:

• List species from left to right in the order: anode | anode solution (concentration) || cathode solution (concentration) | cathode.

• A single vertical line (|) marks a phase boundary — a change between different physical states or phases.

• A double vertical line (||) marks a salt bridge (or another separator) between the two half-cells.

Reading the Daniell cell's diagram: Zn(s) | Zn²⁺(1 M) represents solid zinc being oxidized into dissolved zinc ions at the anode, and Cu²⁺(1 M) | Cu(s) represents dissolved copper ions being reduced into solid copper at the cathode.

Electrolytic Cells: Electrolysis of Molten NaCl

An electrolytic cell uses electrical energy to drive a nonspontaneous chemical reaction — the opposite relationship of a galvanic cell, which generates electricity from a reaction that already wants to happen. Electrolytic cells power industrial processes like electroplating, water electrolysis, and the production of chemicals like chlorine and sodium hydroxide.

An electrolytic cell still has an anode and a cathode, but their charge polarity flips relative to a galvanic cell:

• Anode: the positive electrode — oxidation still happens here, even though it's positively charged.

• Cathode: the negative electrode — reduction still happens here.

• Electrolyte: a molten salt or aqueous solution that lets ions move between the electrodes, completing the circuit.

• External power source: required to force electrons to flow from anode to cathode, since the reaction won't proceed on its own.

The electrolysis of molten sodium chloride (NaCl) is a standard industrial example, used to produce sodium metal and chlorine gas. When NaCl is molten, it dissociates into Na⁺ and Cl⁻ ions, and an external power source drives the reaction:

• Cathode (reduction): Na⁺ + e⁻ → Na(l)

• Anode (oxidation): 2Cl⁻ → Cl₂(g) + 2e⁻

• Overall reaction: 2NaCl(l) → 2Na(l) + Cl₂(g)

This is the main industrial method for producing sodium and chlorine.

Faraday's Laws of Electrolysis

Michael Faraday laid the groundwork for electrochemistry in the early 19th century by formulating the laws of electrolysis — quantitative rules connecting how much electric charge passes through an electrolyte to how much of a substance is chemically altered at an electrode.

• First law: The amount of chemical change produced by an electric current is directly proportional to the quantity of electric charge (Q) passed through the electrolyte.

• Second law: When the same quantity of electricity passes through different electrolytes, the amounts of different substances liberated are proportional to their equivalent weights.

Faraday also defined the Faraday constant (F) — the charge carried by one mole of electrons, F ≈ 96,485 C/mol e⁻.

From the first law, the moles of substance produced or consumed at an electrode during electrolysis are given by:

moles = It / nF

where I is the current, t is the time the current runs, n is the number of moles of electrons transferred per unit of the substance (its valence in the reaction), and F is the Faraday constant. To convert that to the mass deposited, multiply by the substance's molar mass (M):

mass = (It × M) / nF

This equation is what lets you calculate how much metal will plate out during electroplating, or how much gas will be produced during electrolysis, for a given current and time.

Concentration Cells

A concentration cell is a galvanic cell where both electrodes are made of the same material — for example, two copper electrodes — so they share the same reduction potential. Instead of a chemical difference between the electrodes driving the reaction, the driving force is a concentration gradient between the two solutions surrounding them.

That concentration gradient creates a potential difference, which pushes electrons to flow in the direction that equalizes the ion concentrations between the two half-cells. Once the concentrations become equal, the gradient — and the current — disappears, and the cell's voltage drops to zero. (The exact voltage at any given concentration difference is calculated with the Nernst equation, covered separately.)

A concentration cell doesn't have to be a lab setup — the cell membrane of a neuron is a biological example. The membrane maintains different concentrations of sodium (Na⁺), potassium (K⁺), and chloride (Cl⁻) ions on either side, and that difference in ion concentration creates a membrane potential. The sodium-potassium pump actively transports Na⁺ out of the cell and K⁺ into the cell (against their concentration gradients, powered by ATP), maintaining the gradient responsible for the neuron's resting membrane potential. A large enough disturbance to that resting potential can trigger an action potential — the electrical signal behind nerve impulses.

Rechargeable Cells: Lead-Acid and Nickel-Cadmium Batteries

A rechargeable cell (rechargeable battery) can function as both a galvanic cell (while discharging/providing power) and an electrolytic cell (while being recharged with an external power source).

Lead-acid batteries, widely used in vehicles, operate as galvanic cells when fully charged:

• Anode: lead (Pb)

• Cathode: porous lead dioxide (PbO₂)

• Electrolyte: concentrated sulfuric acid (H₂SO₄), around 4 M

During discharge, both electrodes react with the sulfuric acid to form lead sulfate (PbSO₄), and the acid becomes diluted as the reaction proceeds. Lead-acid batteries have a relatively low energy density (energy stored per unit weight), which makes them bulkier for the power they deliver — acceptable in cars, where weight is less of a constraint than in portable electronics.

Nickel-cadmium (Ni-Cd) batteries are another rechargeable type, familiar as AA/AAA cells:

• Anode: cadmium (Cd)

• Cathode: nickel(III) oxide-hydroxide

• Electrolyte: potassium hydroxide (KOH)

Ni-Cd batteries have higher energy density than lead-acid batteries and can deliver high surge currents — bursts of large current early in the discharge cycle — which suits devices like power tools and remote controls. Modern devices have largely moved to nickel-metal hydride (NiMH) batteries instead, which offer even higher energy density, lower cost, and significantly less toxicity than Ni-Cd.

Common MCAT Mistakes

• Mixing up electron flow with conventional current. Electrons physically flow anode → cathode in every cell type; conventional current is defined as flowing the opposite direction, cathode → anode, because it tracks positive charge.

• Assuming the anode is always negative. True in a galvanic cell, but an electrolytic cell's external power source flips the polarity — the anode becomes positive and the cathode becomes negative, even though oxidation still happens at the anode and reduction still happens at the cathode.

• Misreading cell diagram notation. A single vertical line (|) marks a phase boundary; a double vertical line (||) marks a salt bridge. Swapping the two misreads which part of the diagram is the physical separator between half-cells.

• Thinking a concentration cell needs two different electrode materials. It doesn't — both electrodes are chemically identical. The driving force is a concentration gradient between the two solutions, not a difference in reduction potential between different metals.

MCAT-Style Concept Check

Question: A current of 4.825 A is passed through molten NaCl for 4,000 seconds to electrolytically deposit sodium metal at the cathode (Na⁺ + e⁻ → Na(l), n = 1). Using F ≈ 96,485 C/mol e⁻, how many moles of sodium metal are deposited?

• A) 0.10 mol

• B) 0.20 mol

• C) 0.40 mol

• D) 2.00 mol

Answer: B

Explanation: moles = It/nF = (4.825 A × 4,000 s) / (1 × 96,485 C/mol) = 19,300 / 96,485 ≈ 0.20 mol. This is a direct application of Faraday's first law, which ties the amount of charge passed through an electrolyte to the amount of substance liberated at an electrode.

FAQ

What's the difference between a galvanic cell and an electrolytic cell?

A galvanic cell runs a spontaneous redox reaction and generates a positive emf on its own (ΔG < 0), like a battery. An electrolytic cell runs a nonspontaneous reaction and needs an external power source to force it forward, producing a negative emf (ΔG > 0), like the electrolysis of molten NaCl.

Why does conventional current flow opposite to electron flow?

Electrons are the actual physical charge carriers, and they move from anode to cathode in every electrochemical cell. Conventional current is defined by physics convention as the flow of positive charge, which is opposite in direction to the flow of negative electrons — so current flows cathode to anode.

What is a salt bridge for?

A salt bridge connects the two half-cells of a galvanic cell like the Daniell cell and maintains electrical neutrality by letting ions flow between the compartments — without it, charge would build up in each half-cell and the reaction would stall almost immediately.

How do concentration cells generate voltage without different electrode materials?

The two electrodes are chemically identical and share the same reduction potential, so the driving force instead comes from a concentration gradient between the two solutions surrounding them. Electrons flow in the direction that equalizes the ion concentrations, and the voltage drops to zero once the concentrations equalize.