Electrochemical Cells ⚡

Electrochemical cells are systems that use a redox reaction to either produce electricity or force a non-spontaneous reaction to happen. In this lesson, students, you will learn how these cells work, how to tell a galvanic cell from an electrolytic cell, and how to connect the ideas to redox chemistry in IB Chemistry HL. The key objectives are to understand the main terms, explain the flow of electrons and ions, apply cell notation and standard electrode potentials, and see why electrochemical cells are such an important example of how chemical change can be controlled and measured.

A real-world hook: every battery in a phone, laptop, flashlight, or electric car depends on electrochemical ideas 🔋. Some cells make electricity from chemistry, while others use electricity to drive chemistry. That connection between chemical change and electrical energy is central to Reactivity 3.

What an Electrochemical Cell Is

An electrochemical cell is a setup in which oxidation and reduction happen in separate places, so electrons move through an external wire instead of directly from one substance to another. This movement of electrons is what creates or uses electrical energy.

There are two main kinds of electrochemical cells:

Galvanic cell: a spontaneous redox reaction produces electrical energy.
Electrolytic cell: electrical energy is supplied to force a non-spontaneous redox reaction.

The important redox ideas are always the same:

Oxidation means loss of electrons.
Reduction means gain of electrons.
The reducing agent is oxidized.
The oxidizing agent is reduced.

A useful memory tool is OIL RIG: oxidation is loss, reduction is gain.

In an electrochemical cell, the reaction is split into half-equations. One half-cell carries out oxidation, and the other carries out reduction. Because electrons cannot travel through the solution very efficiently, they move through a wire or external circuit instead. That electron flow can power a device, like a small motor or a light bulb 💡.

Galvanic Cells: Turning Chemical Energy into Electrical Energy

A galvanic cell is also called a voltaic cell. It uses a spontaneous redox reaction. Since spontaneous reactions release energy, that energy can be captured as electricity.

A classic example is the zinc-copper cell. Zinc metal is placed in a solution containing zinc ions, and copper metal is placed in a solution containing copper ions. The two half-cells are connected by a wire and a salt bridge.

The oxidation half-equation is:

$$\mathrm{Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-}$$

The reduction half-equation is:

$$\mathrm{Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)}$$

The overall equation is:

$$\mathrm{Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)}$$

In this example, zinc is oxidized and acts as the reducing agent. Copper ions are reduced and act as the oxidizing agent. The electrons flow from zinc to copper through the wire.

Important electrode terms:

Anode: where oxidation occurs
Cathode: where reduction occurs

In a galvanic cell, the anode is negative because it releases electrons, and the cathode is positive because it attracts electrons. This often confuses students, so students, remember: the names anode and cathode tell you oxidation and reduction, not the sign directly.

The salt bridge is also essential. It allows ions to move so that charge does not build up in either half-cell. Without it, the reaction would quickly stop because the solutions would become too positively or negatively charged. Usually, the salt bridge contains an inert electrolyte such as potassium nitrate.

Cell Notation and Standard Electrode Potentials

Chemists use cell notation to write electrochemical cells in a compact form. For the zinc-copper cell, the notation is:

$$\mathrm{Zn(s)\ |\ Zn^{2+}(aq)\ ||\ Cu^{2+}(aq)\ |\ Cu(s)}$$

The single line $|$ shows a phase boundary, and the double line $||$ shows the salt bridge.

Cell notation is written with the anode on the left and the cathode on the right. This is a standard IB convention.

To predict whether a cell will work spontaneously, you can use standard electrode potentials. These are measured under standard conditions:

temperature of $298\ \mathrm{K}$
concentration of $1.0\ \mathrm{mol\ dm^{-3}}$
pressure of $100\ \mathrm{kPa}$ for gases

Each half-equation has a standard reduction potential, written as $E^\circ$. More positive values mean a stronger tendency to be reduced.

For a galvanic cell, the standard cell potential is:

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

If $E^\circ_{\text{cell}} > 0$, the reaction is spontaneous under standard conditions. If $E^\circ_{\text{cell}} < 0$, it is not spontaneous.

Example: if the reduction potential for $\mathrm{Cu^{2+}/Cu}$ is more positive than that for $\mathrm{Zn^{2+}/Zn}$, then copper is reduced and zinc is oxidized. That is why the zinc-copper cell can produce electricity.

This connects directly to Reactivity 3 because it links chemical reactivity to measurable energy changes. The more favorable the electron transfer, the more likely the redox process is to occur spontaneously.

Electrolytic Cells: Using Electricity to Force Change

An electrolytic cell does the opposite job of a galvanic cell. It uses an external power supply to make a non-spontaneous redox reaction happen. This is extremely important in industry and in the lab ⚙️.

A common example is the electrolysis of molten sodium chloride. When sodium chloride is molten, the ions are free to move.

At the cathode, reduction occurs:

$$\mathrm{Na^+(l) + e^- \rightarrow Na(l)}$$

At the anode, oxidation occurs:

$$\mathrm{2Cl^-(l) \rightarrow Cl_2(g) + 2e^-}$$

The overall reaction is:

$$\mathrm{2NaCl(l) \rightarrow 2Na(l) + Cl_2(g)}$$

In an electrolytic cell, the anode is positive and the cathode is negative because the power supply pulls electrons away from the anode and pushes electrons toward the cathode. This is another common place where students get mixed up. The rule still holds: oxidation at the anode, reduction at the cathode.

Electrolysis is used in real life to extract reactive metals, plate objects with metals, and produce chemicals such as chlorine and sodium hydroxide in the chlor-alkali process. That shows how electrochemistry is not just a classroom idea, but a major part of modern manufacturing.

How to Analyze a Cell in IB Chemistry HL

When solving cell questions, students, use a clear method:

Identify the oxidized species and the reduced species.
Write the two half-equations.
Label anode and cathode correctly.
Determine the direction of electron flow.
Write the overall equation.
If given standard electrode potentials, calculate $E^\circ_{\text{cell}}$.

A useful logic chain is: the species with the more negative reduction potential is more likely to be oxidized, while the species with the more positive reduction potential is more likely to be reduced.

Example reasoning: if half-cell A has a very strong tendency to lose electrons and half-cell B has a stronger tendency to gain electrons, then electrons will flow from A to B. That means A is the anode and B is the cathode.

Sometimes IB questions ask about mass changes. In a galvanic cell, the anode metal may dissolve because atoms become ions. At the cathode, metal ions may be deposited as solid metal. This is evidence of redox taking place. If a metal electrode gets smaller, that side is being oxidized; if it gets larger, reduction is happening there.

Why Electrochemical Cells Matter in Reactivity 3

Electrochemical cells bring together several big ideas from Reactivity 3:

Redox chemistry: oxidation and reduction happen together.
Energy changes: spontaneous redox reactions can do electrical work.
Reaction pathways: electrons move by a controlled route through a wire rather than directly between substances.
Chemical reactivity: the electrode potential helps compare how easily substances are oxidized or reduced.

This topic also connects to broader chemistry. For example, in acid-base chemistry, some redox reactions depend on $\mathrm{H^+}$ concentration. In electrochemistry, changing concentration can change the position of equilibrium and the cell voltage. That means conditions matter, not just the identities of the substances.

Electrochemical cells are also a bridge to analytical chemistry. Measurements of voltage can help identify whether a reaction is feasible and compare the relative strengths of oxidizing and reducing agents.

Conclusion

Electrochemical cells show how chemical change can be turned into useful energy or driven by an external source. In a galvanic cell, a spontaneous redox reaction produces electricity. In an electrolytic cell, electricity forces a non-spontaneous redox reaction. The key ideas are oxidation at the anode, reduction at the cathode, electron flow through the external circuit, ion movement through the electrolyte or salt bridge, and the use of standard electrode potentials to predict behavior.

For IB Chemistry HL, students, the most important skill is to explain the process clearly using correct terminology and evidence. If you can identify the half-reactions, assign the electrodes, and interpret $E^\circ_{\text{cell}}$, you can handle most electrochemical cell questions confidently. ⚡

Study Notes

Electrochemical cells involve redox reactions and the transfer of electrons.
In a galvanic cell, a spontaneous reaction produces electrical energy.
In an electrolytic cell, electrical energy forces a non-spontaneous reaction.
Oxidation occurs at the anode.
Reduction occurs at the cathode.
In a galvanic cell, the anode is negative and the cathode is positive.
In an electrolytic cell, the anode is positive and the cathode is negative.
Electrons travel through the external wire; ions travel through the electrolyte or salt bridge.
Cell notation places the anode on the left and the cathode on the right.
Standard cell potential is calculated using $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$.
If $E^\circ_{\text{cell}} > 0$, the reaction is spontaneous under standard conditions.
Electrochemical cells connect redox chemistry to energy, reactivity, and real-world technology.