19.3 Galvanic Cells
PRE-HEALTH PREP
Galvanic Cells
Anode, Cathode, and Salt Bridge
General Chemistry 19.3
Lesson Overview
Galvanic (voltaic) cells convert the energy of spontaneous redox reactions into electrical current by separating oxidation and reduction half-reactions into distinct compartments connected by an external wire. The anode, where oxidation occurs, loses mass as metal dissolves into solution, while the cathode gains mass as metal ions plate out — with a salt bridge maintaining electrical neutrality by allowing ion migration between half-cells. Electrode polarity, electron flow direction, the role of inert electrodes such as platinum, and the distinguishing features of electrolytic cells are also addressed.
Table of Contents
- The 2 Types of Electrochemical Cells
- Spontaneity: Delta G and Cell Potential
- Galvanic vs. Electrolytic: Energy In and Out
- 2 Electrodes: Anode and Cathode
- Electron Flow Through the Wire
- Electrode Polarity: 2 Rules to Know
- The Zinc-Copper Galvanic Cell
- The 2 Half-Reactions
- Setting Up the Half-Cell Solutions
- The Salt Bridge
- Ion Flow vs. Electron Flow: A Critical Distinction
- Electrolytic Cells: A Brief Contrast
- Inert Electrodes
- 10 Key Takeaways
The 2 Types of Electrochemical Cells
A galvanic cell — also called a voltaic cell — is one of two primary types of electrochemical cells used in electrochemistry. The terms are completely synonymous; you will encounter "galvanic cell" more often in a general chemistry course, but treat them interchangeably. The other major type is the electrolytic cell, in which electrolysis occurs. This lesson focuses on galvanic cells, but understanding where they fit in the broader picture is essential.
The defining difference between these two cell types comes down to spontaneity.
| Property | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| Reaction type | Spontaneous | Non-spontaneous |
| Energy flow | Produces electrical energy | Consumes electrical energy |
| Delta G | Negative | Positive |
| Cell potential (Ecell) | Positive | Negative |
| Anode sign | Negative (−) | Positive (+) |
| Cathode sign | Positive (+) | Negative (−) |
| Common application | Batteries, fuel cells | Industrial production of elements (Cl₂, Na) |
Spontaneity: Delta G and Cell Potential
Through thermodynamics, you learned to assess reaction spontaneity using Gibbs free energy (ΔG): a negative ΔG signals a spontaneous reaction, while a positive ΔG signals a non-spontaneous one. For redox reactions, there is a second — and often more practical — measure of spontaneity: cell potential.
Cell potential goes by several names. You may see it written as EMF (electromotive force), Ecell, or simply "voltage," measured in volts (V). The abbreviation Ecell uses either the Greek letter epsilon (ε) or a capital E, depending on the textbook.
Sign Rule: When Ecell is positive, the reaction is spontaneous. When Ecell is negative, the reaction is non-spontaneous. Delta G and Ecell always carry opposite signs — this is not a coincidence. They are mathematically related by the equation ΔG = −nFEcell, which you will explore in a later lesson in this chapter.
Galvanic vs. Electrolytic: Energy In and Out
Because a galvanic cell runs a spontaneous reaction, it produces electrical energy. This is precisely how a battery works: when a battery discharges in its forward, spontaneous direction, it is operating as a galvanic cell. The chemical energy released by the spontaneous redox reaction is converted into electrical energy that powers your device.
When you recharge that same battery, you are forcing the spontaneous reaction to run in reverse — a non-spontaneous process. During recharging, the cell is functioning as an electrolytic cell, consuming electrical energy from an external source to reverse the chemistry.
Electrolytic cells have important industrial applications beyond rechargeable batteries. The large-scale production of elemental chlorine and elemental sodium, for instance, relies on electrolysis of molten sodium chloride. These examples will be covered in detail in the final lessons of this chapter.
2 Electrodes: Anode and Cathode
Every electrochemical cell contains two electrodes — typically metal conductors suspended in a solution. These electrodes carry specific names tied to the chemistry occurring at each one.
- Anode — the electrode at which oxidation takes place (electrons are lost)
- Cathode — the electrode at which reduction takes place (electrons are gained)
Memory Aid: Notice that anode and oxidation both begin with vowels, while cathode and reduction both begin with consonants. This simple pattern is a reliable way to keep the definitions straight.
Characteristics of Electrochemical Cells
1. The anode is ALWAYS the site of oxidation. (an ox)
2. The cathode is ALWAYS the site of reduction. (red cat)
3. Electrons ALWAYS flow from anode to cathode.
4. Anions flow to the anode and cations to the cathode through the salt bridge (in galvanic cells).
5. In galvanic/voltaic cells (spontaneous), the cathode is + and the anode is −. In electrolytic cells (nonspontaneous), the cathode is − and the anode is +.
6. Often, the cathode gains mass, while the anode loses mass (for metal/metal salt galvanic cells).
By formal convention, when drawing a diagram of a galvanic cell the anode is always placed on the left and the cathode on the right. This matters more than it might seem: certain calculations reference "left" and "right" as shorthand for anode and cathode. Not every instructor adheres to this convention, but knowing it will serve you well on standardized exams and in more advanced coursework.
Electron Flow Through the Wire
Since electrons are lost at the anode and gained at the cathode, they must travel from the anode to the cathode. The wire connecting the two electrodes provides the path for this electron transfer.
Key Fact: Electrons travel through the wire from anode to cathode — not through the solution. This distinction is frequently tested. A useful memory trick: anode to cathode follows alphabetical order (A before C).
In a galvanic cell, this direction of electron flow is the spontaneous direction — it is the direction the electrons "want" to travel. In an electrolytic cell, an external power source forces electrons to move in the opposite direction, against their natural tendency.
Electrode Polarity: 2 Rules to Know
In a galvanic cell, the anode carries a negative sign and the cathode carries a positive sign. A practical way to remember this: negatively charged electrons want to move away from the negative electrode (anode) and toward the positive electrode (cathode), just as opposite charges attract.
A word of clarification on what "negative" and "positive" actually mean here. The signs on the electrodes refer to their relative cell potentials, not to static charge buildup. The anode has a lower cell potential and the cathode has a higher cell potential. Because electrons have higher potential energy at lower potential (the opposite of the convention for positive charges), electrons spontaneously flow from lower potential to higher potential — or equivalently, from higher potential energy to lower potential energy. The physics works out consistently; the terminology is just counterintuitive at first.
In an electrolytic cell the signs are reversed: the anode is positive and the cathode is negative. Because the spontaneous electron-flow direction is now opposite to the forced direction, an external power source is required to drive the reaction.
| Electrode | Sign in Galvanic Cell | Sign in Electrolytic Cell |
|---|---|---|
| Anode | − (negative) | + (positive) |
| Cathode | + (positive) | − (negative) |
Electron flow through the wire is always from anode to cathode in both cell types — what changes between galvanic and electrolytic cells is the polarity assigned to each electrode.
The Zinc-Copper Galvanic Cell
The classic textbook galvanic cell pairs zinc and copper. Under standard conditions, the overall reaction is spontaneous — it carries a positive cell potential and a negative ΔG. You will learn to calculate Ecell quantitatively in the next lesson; for now, focus on the structure and logic of the cell.
The fundamental design principle of a galvanic cell is separation. The oxidation and reduction half-reactions are physically divided into two compartments, called half-cells. By forcing the electrons to travel through an external wire to get from one half-cell to the other, the cell can do useful electrical work. If both half-reactions occurred in the same solution, electrons would transfer directly between species in solution and no current would flow through an external circuit — you would have chemistry, but no battery.
The 2 Half-Reactions
In the zinc-copper cell, the two half-reactions are:
Example: Zinc-Copper Half-Reactions
Anode (oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻
Zinc metal loses two electrons. Its oxidation state rises from 0 to +2. The solid zinc electrode is consumed in the process.
Cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)
Copper(II) ions in solution gain two electrons. The oxidation state falls from +2 to 0. Copper metal plates out onto the electrode surface.
The electrode material is determined by which species in each half-reaction exists as a solid metal, since a solid metal can conduct electricity. At the anode, solid zinc serves as the electrode — it is the reactant being consumed. At the cathode, solid copper is the product being deposited; it serves as (or builds up on) the electrode. Whether the solid metal appears on the reactant side or the product side does not affect this rule — as long as a solid, electrically conductive metal is present in the half-reaction, that is the electrode material.
Mass Changes at the Electrodes: The anode dissolves over time (zinc metal is consumed), so the anode electrode loses mass. The cathode gains mass as copper metal plates out of solution. This is a hallmark of many galvanic cells and is testable on exams.
Setting Up the Half-Cell Solutions
Each half-cell requires a solution containing the ionic species for that half-reaction. For the zinc-copper cell, that means Zn²⁺ ions in the anode compartment and Cu²⁺ ions in the cathode compartment. Under standard conditions, both concentrations are 1.00 M.
These ions cannot exist in solution without accompanying counterions — solutions must be electrically neutral. The practical approach is to dissolve the corresponding ionic salts. A typical setup uses zinc sulfate (ZnSO₄) in the anode compartment and copper(II) sulfate (CuSO₄) in the cathode compartment. The sulfate ions (SO₄²⁻) are spectator ions — they are present to balance charge but do not participate in the half-reactions.
Keep this detail in mind: the presence of spectator anions (such as sulfate) in each compartment becomes important when analyzing what the salt bridge needs to do once the cell begins operating.
The Salt Bridge
A galvanic cell with two separate half-cells but no connection between the solutions is a dead battery. The cell requires a salt bridge to function. Without it, the half-cell solutions quickly develop opposing charge imbalances that bring the reaction to a halt before it even starts.
A salt bridge is a tube filled with an aqueous ionic solution — commonly NaNO₃ (sodium nitrate) — often suspended in a gel matrix to prevent bulk mixing of solutions. The bridge physically connects the two half-cells, allowing ions to flow between them while keeping the solutions from mixing directly.
Purpose of the Salt Bridge: Maintaining Electrical Neutrality
As the cell operates, charge imbalances develop in each compartment:
- At the anode, zinc metal dissolves to produce Zn²⁺ ions. This continuously adds positive charge to the solution. Without intervention, the anode compartment would become increasingly positive, which would be energetically unfavorable and would stop the reaction.
- At the cathode, Cu²⁺ ions are consumed as copper metal plates out. This removes positive charge from the solution, leaving an excess of negative sulfate ions. Without intervention, the cathode compartment would become increasingly negative.
The salt bridge neutralizes both imbalances simultaneously by allowing ions to migrate between compartments:
Ion Migration Through the Salt Bridge
Anions flow into the anode compartment: Nitrate ions (NO₃⁻) migrate from the salt bridge into the anode compartment to counteract the buildup of positive zinc ions.
Cations flow into the cathode compartment: Sodium ions (Na⁺) migrate from the salt bridge into the cathode compartment to counteract the depletion of positive copper ions.
A useful memory pattern: anions flow into the anode compartment; cations flow into the cathode compartment. Both words even share their starting letters.
Ion Flow vs. Electron Flow: A Critical Distinction
Common Exam Trap: Electrons flow through the wire from anode to cathode. Ions flow through the salt bridge between the two half-cells. These are two completely separate pathways carrying charge in the circuit. Do not confuse them.
Together, these two pathways complete the electrical circuit. Electrons carry charge through the external wire, while ions carry charge through the internal solution pathway (the salt bridge). Both are necessary for sustained current flow.
Electrolytic Cells: A Brief Contrast
An electrolytic cell runs a non-spontaneous reaction by consuming electrical energy from an external source. Because it does not generate electricity, there is no need to separate the reaction into two half-cells in separate compartments — the goal is not to harness electron flow through a wire, but simply to force a reaction to occur.
Example: Electrolysis of Molten Sodium Chloride
Molten (liquid) NaCl is electrolyzed to produce elemental sodium and elemental chlorine:
Anode (oxidation): 2Cl⁻ → Cl₂ + 2e⁻ (chloride ions lose electrons)
Cathode (reduction): 2Na⁺ + 2e⁻ → 2Na (sodium ions gain electrons)
The electrons are pumped through the wire by the external power source. Chlorine is produced as a gas; sodium may be solid or liquid depending on the operating temperature. No salt bridge is needed because there is only one compartment.
Notice that in the electrolytic cell, the negatively charged chloride ions are attracted to the positive anode, and the positively charged sodium ions are attracted to the negative cathode. This reinforces a consistent pattern across both cell types: anions are oxidized at the anode; cations are reduced at the cathode.
Electrolytic cells and the broader topic of electrolysis will be covered in depth in the final lessons of this chapter.
Inert Electrodes
The electrode material used in the examples above — solid zinc and solid copper — participates directly in the half-reactions. This is the common case, but not the only one. Sometimes no species in a half-reaction exists as a solid metal that can serve as the electrode. In these situations, an inert electrode is used.
An inert electrode conducts electrons to or from the solution without participating chemically in the half-reaction. The most common material is platinum, which is chemically unreactive under most electrochemical conditions.
Example: The Standard Hydrogen Electrode (SHE)
The standard hydrogen electrode involves the following half-reaction:
2H⁺(aq) + 2e⁻ → H₂(g)
Neither hydrogen ions (aqueous) nor hydrogen gas can serve as a solid electrode. A platinum electrode is used instead. The platinum conducts electrons to the hydrogen ions in solution, which are reduced to hydrogen gas. The platinum itself is neither consumed nor produced — it is purely a conduit for electron transfer.
The SHE is the reference electrode against which all standard reduction potentials are measured. It will reappear in the next lesson on calculating cell potentials.
Inert Electrodes and Mass Change: Unlike the zinc-copper cell, an inert electrode does not gain or lose mass during the reaction. The earlier generalization — that the anode loses mass and the cathode gains mass — applies only when the electrode material participates in the half-reaction. When an inert electrode is used, electrode mass remains constant.
10 Key Takeaways
- Two types of electrochemical cells: Galvanic (voltaic) cells run spontaneous reactions and produce electrical energy; electrolytic cells run non-spontaneous reactions and consume electrical energy from an external source.
- Spontaneity indicators: A negative ΔG and a positive Ecell both signal a spontaneous reaction; the two quantities always carry opposite signs and are mathematically related.
- Anode and cathode definitions: Oxidation always occurs at the anode; reduction always occurs at the cathode — in all electrochemical cells, both galvanic and electrolytic. By convention, the anode is drawn on the left and the cathode on the right.
- Electrode polarity: In a galvanic cell, the anode is negative (−) and the cathode is positive (+). In an electrolytic cell, the signs are reversed.
- Electron flow: Electrons travel through the external wire from anode to cathode in all electrochemical cells.
- Half-cells in galvanic cells: The oxidation and reduction half-reactions occur in separate compartments (half-cells), forcing electrons through the wire where they can do electrical work.
- Salt bridge function: The salt bridge maintains electrical neutrality in both half-cell solutions by allowing anions to migrate into the anode compartment and cations to migrate into the cathode compartment.
- Electron vs. ion pathways: Electrons carry charge through the wire; ions carry charge through the salt bridge. These are two distinct conduction pathways that together complete the circuit.
- Electrode mass changes: In cells with reactive electrodes, the anode loses mass (metal dissolves) and the cathode gains mass (metal deposits). Inert electrodes (e.g., platinum in the SHE) do not change in mass.
- Inert electrodes: When no species in a half-reaction is a solid conductor, an inert electrode such as platinum is used to transfer electrons without participating in the chemistry.