5. Cell Potential

What Is Cell Potential?

A galvanic cell converts:

chemical energy → electrical energy

The redox reaction causes electrons to move through an external circuit from the:

anode → cathode

The driving force that pushes electrons through the circuit is related to the difference in electrical potential between the two electrodes.

This difference is called the:

cell potential.

Cell potential is also commonly called:

• cell voltage
• electromotive force (emf)
• potential difference

The symbol normally used is:

Ecell

Cell potential is measured in:

volts (V).

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What Does Voltage Tell Us?

Voltage tells us how much electrical potential difference exists between the two electrodes.

A larger positive cell potential generally means a stronger thermodynamic driving force for the cell reaction as written.

For example:

Cell A:

Ecell = +0.40 V

Cell B:

Ecell = +1.50 V

Cell B has the greater potential difference.

However, voltage does not directly tell us:

• how much current the cell can provide
• how long the cell will operate
• how much reactant is present
• how quickly the reaction occurs

These are different properties.

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Where Does Cell Potential Come From?

Different substances have different tendencies to:

gain or lose electrons.

Some species are relatively easy to oxidize.

Others are relatively easy to reduce.

When two suitable half-cells are connected, this difference in their tendencies to undergo reduction produces a:

potential difference.

The greater the difference between the electrode potentials, the greater the cell voltage.

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Electrode Potentials

Each half-cell has an electrode potential.

Electrode potentials describe the tendency of a half-reaction to occur as a:

reduction.

For example:

Cu²⁺(aq) + 2e⁻ → Cu(s)

has a standard reduction potential of approximately:

E° = +0.34 V

While:

Zn²⁺(aq) + 2e⁻ → Zn(s)

has:

E° = −0.76 V

The ° symbol indicates:

standard conditions.

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Standard Electrode Potentials

Electrode potentials cannot be measured independently.

Instead, they are measured relative to a reference electrode.

The standard reference is the:

standard hydrogen electrode (SHE).

Its standard electrode potential is defined as:

E° = 0.00 V

Other half-cells are compared with this reference.

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Standard Conditions

Standard electrode potentials are tabulated under specified standard conditions.

Commonly this means approximately:

• dissolved species at standard-state concentration/activity
• gases at standard-state pressure
• pure solids and liquids
• specified temperature, often 25°C or 298 K

At introductory level, solution concentrations are commonly treated as:

1 mol/L

and gases approximately:

1 bar.

If conditions change, the measured cell potential may also change.

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Standard Cell Potential

The standard cell potential is represented by:

E°cell

It can be calculated from standard reduction potentials using:

E°cell = E°cathode − E°anode

Remember:

cathode = reduction

anode = oxidation

The values in a standard electrode-potential table are normally written as:

reduction potentials.

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A Useful Memory Rule

For a galvanic cell:

Cathode = reduction

Anode = oxidation

Then:

E°cell = E°cathode − E°anode

A useful sequence is:

identify cathode → identify anode → subtract

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Example: Zinc-Copper Cell

Consider the familiar cell:

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Standard reduction potentials:

Cu²⁺ + 2e⁻ → Cu E° = +0.34 V

Zn²⁺ + 2e⁻ → Zn E° = −0.76 V

Copper is the cathode.

Zinc is the anode.

Therefore:

E°cell = E°cathode − E°anode

E°cell = (+0.34) − (−0.76)

E°cell = +1.10 V

The standard cell potential is:

+1.10 V.

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Why Do We Subtract?

The electrode potentials are listed as reduction potentials.

At the cathode, reduction actually occurs.

For copper:

Cu²⁺ + 2e⁻ → Cu

At the anode, however, the listed reduction reaction is reversed.

Listed:

Zn²⁺ + 2e⁻ → Zn E° = −0.76 V

Actual reaction:

Zn → Zn²⁺ + 2e⁻

Reversing the reaction reverses the sign:

E°oxidation = +0.76 V

We could therefore calculate:

E°cell = E°reduction + E°oxidation

E°cell = +0.34 + 0.76

E°cell = +1.10 V

This gives the same answer.

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Two Valid Calculation Methods

You can use either method.

Method 1

Using reduction potentials:

E°cell = E°cathode − E°anode

Method 2

Convert the anode value into an oxidation potential:

E°cell = E°reduction + E°oxidation

Both methods give the same result.

For consistency, the first method is often easier.

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Standard Reduction Potential Table

A simplified table might contain:

Ag⁺ + e⁻ → Ag +0.80 V

Cu²⁺ + 2e⁻ → Cu +0.34 V

2H⁺ + 2e⁻ → H₂ 0.00 V

Fe²⁺ + 2e⁻ → Fe −0.44 V

Zn²⁺ + 2e⁻ → Zn −0.76 V

Mg²⁺ + 2e⁻ → Mg −2.37 V

More positive values indicate a greater tendency for the species on the left side of the reduction half-equation to:

gain electrons under standard conditions.

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Which Species Is Reduced?

When comparing two standard reduction potentials, the half-reaction with the more positive reduction potential normally occurs as:

reduction.

It therefore forms the:

cathode.

For example:

Cu²⁺/Cu:

+0.34 V

Zn²⁺/Zn:

−0.76 V

Because +0.34 V is more positive:

Cu²⁺ is reduced.

Therefore copper is the:

cathode.

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Which Species Is Oxidized?

The half-reaction with the less positive reduction potential is reversed and occurs as:

oxidation.

For zinc:

Zn²⁺ + 2e⁻ → Zn −0.76 V

is reversed:

Zn → Zn²⁺ + 2e⁻

Therefore zinc is the:

anode.

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Example: Magnesium-Copper Cell

Standard reduction potentials:

Cu²⁺ + 2e⁻ → Cu +0.34 V

Mg²⁺ + 2e⁻ → Mg −2.37 V

Copper has the more positive reduction potential.

Therefore:

Cathode = Cu

Anode = Mg

Calculate:

E°cell = E°cathode − E°anode

E°cell = +0.34 − (−2.37)

E°cell = +2.71 V

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Comparing Zinc-Copper and Magnesium-Copper Cells

Zinc-copper:

E°cell = +1.10 V

Magnesium-copper:

E°cell = +2.71 V

Therefore, under standard conditions:

Mg-Cu has the larger cell potential.

This occurs because the difference between the Mg²⁺/Mg and Cu²⁺/Cu electrode potentials is greater.

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Example: Iron-Copper Cell

Standard reduction potentials:

Cu²⁺ + 2e⁻ → Cu +0.34 V

Fe²⁺ + 2e⁻ → Fe −0.44 V

Copper is the cathode.

Iron is the anode.

Therefore:

E°cell = +0.34 − (−0.44)

E°cell = +0.78 V

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Example: Zinc-Silver Cell

Standard reduction potentials:

Ag⁺ + e⁻ → Ag +0.80 V

Zn²⁺ + 2e⁻ → Zn −0.76 V

Silver is reduced.

Zinc is oxidized.

Therefore:

E°cell = +0.80 − (−0.76)

E°cell = +1.56 V

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Do We Multiply Electrode Potentials?

Consider the zinc-silver cell.

Oxidation:

Zn → Zn²⁺ + 2e⁻

Reduction:

Ag⁺ + e⁻ → Ag

To balance electrons, the silver half-reaction must be multiplied by 2:

2Ag⁺ + 2e⁻ → 2Ag

However:

do NOT multiply the electrode potential by 2.

The silver potential remains:

+0.80 V

Therefore:

E°cell = +0.80 − (−0.76)

E°cell = +1.56 V

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Why Isn't Voltage Multiplied?

Electrode potential is an:

intensive property.

It does not depend on how much material is present.

Doubling a half-equation doubles quantities such as:

• number of moles
• number of electrons
• amount of chemical change

but it does not double the:

electrode potential.

This is a very important electrochemistry rule.

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Cell Potential and Spontaneity

The sign of the standard cell potential tells us whether a redox reaction is thermodynamically favorable under standard conditions.

If:

E°cell > 0

the reaction as written is:

spontaneous under standard conditions.

If:

E°cell < 0

the reaction as written is:

nonspontaneous under standard conditions.

If:

E°cell = 0

the system is at:

equilibrium under standard conditions.

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Example: Predicting Spontaneity

Suppose:

E°cathode = +0.50 V

and:

E°anode = −0.30 V

Then:

E°cell = +0.50 − (−0.30)

E°cell = +0.80 V

Because E°cell is positive, the reaction is:

spontaneous under standard conditions.

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What If the Cell Potential Is Negative?

Suppose:

E°cell = −0.80 V

This means the proposed reaction is not spontaneous in the direction written under standard conditions.

However, reversing the reaction reverses the sign:

E°cell = +0.80 V

The reverse reaction is therefore thermodynamically favored under standard conditions.

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Cell Potential and Gibbs Free Energy

Cell potential is connected to another thermodynamic quantity:

Gibbs free energy.

The relationship is:

ΔG° = −nFE°cell

where:

ΔG° = standard Gibbs free-energy change

n = number of moles of electrons transferred

F = Faraday constant

E°cell = standard cell potential

The Faraday constant is approximately:

96,485 C/mol

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What Does This Relationship Mean?

If:

E°cell > 0

then:

ΔG° < 0

and the reaction is spontaneous under standard conditions.

If:

E°cell < 0

then:

ΔG° > 0

and the reaction is nonspontaneous under standard conditions.

Therefore:

positive E°cell ↔ negative ΔG°

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Cell Potential Does Not Tell Us Reaction Speed

A positive cell potential tells us about:

thermodynamic favorability.

It does not tell us how quickly the reaction occurs.

A reaction may be thermodynamically spontaneous but proceed very slowly because of:

kinetic barriers.

Therefore:

spontaneous ≠ fast.

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Factors Affecting Cell Potential

The actual cell potential can depend on several factors.

Important factors include:

• identity of the electrodes and redox couples
• ion concentrations
• gas pressures
• temperature
• reaction conditions

The standard potential assumes:

standard-state conditions.

Changing these conditions can change the measured voltage.

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Factor 1: Electrode Materials

Different electrode systems have different reduction potentials.

For example:

Ag⁺/Ag = +0.80 V

Cu²⁺/Cu = +0.34 V

Zn²⁺/Zn = −0.76 V

Mg²⁺/Mg = −2.37 V

Choosing half-cells farther apart in reduction potential generally produces a:

larger standard cell voltage.

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Factor 2: Ion Concentration

Changing ion concentrations can change the:

cell potential.

As a galvanic cell operates:

• reactants are consumed
• products form
• ion concentrations change

Therefore the voltage may gradually change.

Eventually, when equilibrium is reached:

Ecell = 0.

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The Nernst Equation

The relationship between concentration and cell potential can be described using the:

Nernst equation.

At 25°C, a commonly used form is:

E = E° − (0.0592/n) log Q

where:

E = cell potential under current conditions

E° = standard cell potential

n = number of electrons transferred

Q = reaction quotient

You may encounter the Nernst equation in more advanced electrochemistry.

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Factor 3: Gas Pressure

If gases participate in the electrode reactions, changing their pressures can change the:

cell potential.

For example, hydrogen gas is involved in the standard hydrogen electrode.

Changing hydrogen pressure away from standard conditions affects the electrode potential.

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Factor 4: Temperature

Temperature can also affect electrode potentials and therefore:

cell voltage.

The effect depends on the particular chemical reaction.

For this reason, temperature should be specified when accurate electrochemical measurements are required.

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Does Electrode Size Affect Voltage?

Increasing the surface area of an electrode can affect:

• possible current
• reaction rate
• internal resistance effects

However, under ideal equilibrium conditions, simply using a larger piece of the same electrode material does not automatically increase the equilibrium:

cell potential.

A bigger electrode is not equivalent to a higher voltage.

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Voltage vs Current

Voltage and current are different.

Voltage

is the electrical potential difference.

Measured in:

volts (V).

Current

is the rate of flow of electric charge.

Measured in:

amperes (A).

A cell can have a relatively high voltage but provide limited current.

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Voltage vs Electrical Energy

Voltage also should not be confused with total energy.

A small cell and a large cell might have similar voltages but contain very different amounts of reactants.

The larger cell may be able to provide electrical energy for:

much longer.

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Measuring Cell Potential

Cell potential can be measured using a:

voltmeter.

The voltmeter is connected across the two electrodes.

For a functioning galvanic cell:

• electrons leave the anode
• travel through the external circuit
• reach the cathode

A high-resistance voltmeter can measure the potential difference while drawing very little current.

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Connecting the Voltmeter

In a conventional galvanic cell:

anode = negative electrode

cathode = positive electrode

Connecting the positive voltmeter terminal to the cathode and the negative terminal to the anode should produce a:

positive voltage reading.

If the leads are reversed, the reading will have the opposite sign.

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Comparing Galvanic Cells

Suppose three cells have the following standard voltages:

Cell A:

+0.78 V

Cell B:

+1.10 V

Cell C:

+2.71 V

The order of increasing standard cell potential is:

A < B < C

Cell C has the greatest potential difference between its electrodes.

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Example Comparison

Consider:

Fe-Cu cell = +0.78 V

Zn-Cu cell = +1.10 V

Mg-Cu cell = +2.71 V

The copper cathode is the same in each case.

The anodes differ.

Because magnesium has a much more negative standard reduction potential than zinc or iron, the Mg-Cu combination produces the:

largest standard potential difference.

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Predicting the Best Pair

Suppose the available reduction potentials are:

Ag⁺/Ag = +0.80 V

Cu²⁺/Cu = +0.34 V

Zn²⁺/Zn = −0.76 V

Mg²⁺/Mg = −2.37 V

Which pair would produce the largest standard cell potential?

Choose the most positive value as the cathode:

Ag⁺/Ag = +0.80 V

Choose the most negative value as the anode:

Mg²⁺/Mg = −2.37 V

Then:

E°cell = +0.80 − (−2.37)

E°cell = +3.17 V

Therefore the Mg-Ag combination has the largest standard cell potential among these choices.

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Writing the Cell

The Mg-Ag cell can be written:

Mg(s) | Mg²⁺(aq) || Ag⁺(aq) | Ag(s)

Anode:

Mg

Cathode:

Ag

Electron direction:

Mg → Ag

Standard cell potential:

+3.17 V

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Worked Example 1

Calculate E°cell for:

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Given:

Zn²⁺/Zn = −0.76 V

Cu²⁺/Cu = +0.34 V

Use:

E°cell = E°cathode − E°anode

E°cell = +0.34 − (−0.76)

E°cell = +1.10 V

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Worked Example 2

Calculate E°cell for:

Fe(s) | Fe²⁺(aq) || Ag⁺(aq) | Ag(s)

Given:

Fe²⁺/Fe = −0.44 V

Ag⁺/Ag = +0.80 V

Therefore:

E°cell = +0.80 − (−0.44)

E°cell = +1.24 V

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Worked Example 3

Calculate E°cell for a Mg-Zn galvanic cell.

Given:

Mg²⁺/Mg = −2.37 V

Zn²⁺/Zn = −0.76 V

The more positive potential is Zn.

Therefore:

Cathode = Zn

Anode = Mg

Calculate:

E°cell = −0.76 − (−2.37)

E°cell = +1.61 V

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Worked Example 4

A student calculates:

E°cell = −0.50 V

for a proposed reaction.

What does this mean?

The reaction as written is:

nonspontaneous under standard conditions.

The reverse reaction would have:

E°cell = +0.50 V

and would be thermodynamically favored under standard conditions.

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Worked Example 5

A reaction has:

E°cell = +0.65 V

Is it spontaneous?

Under standard conditions:

yes.

Because:

E°cell > 0.

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Worked Example 6

A galvanic cell uses:

Cu²⁺/Cu = +0.34 V

and:

Ag⁺/Ag = +0.80 V

Which electrode is the cathode?

The more positive reduction potential is:

Ag⁺/Ag.

Therefore:

Ag is the cathode.

Calculate:

E°cell = +0.80 − (+0.34)

E°cell = +0.46 V

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Worked Example 7

Compare:

Cell A:

Zn-Cu = +1.10 V

Cell B:

Cu-Ag = +0.46 V

Which has the greater standard potential difference?

Zn-Cu.

Difference:

1.10 − 0.46 = 0.64 V

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Worked Example 8

Why does a battery's voltage sometimes decrease as it operates?

As the reaction proceeds:

• reactants are consumed
• products accumulate
• concentrations change

These changes can alter the:

cell potential.

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Worked Example 9

Two students build identical Zn-Cu cells.

Student A uses a zinc electrode twice as large as Student B.

Will Student A automatically obtain twice the voltage?

No.

Electrode size does not simply multiply the equilibrium cell potential.

It can affect factors such as current capability and reaction surface area.

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Worked Example 10

Why does an Mg-Cu cell have a greater standard voltage than a Zn-Cu cell?

The difference between:

Mg²⁺/Mg and Cu²⁺/Cu

is greater than the difference between:

Zn²⁺/Zn and Cu²⁺/Cu.

A larger difference in electrode potentials produces a larger:

E°cell.

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Cell Potential and Equilibrium

As a spontaneous galvanic-cell reaction proceeds, the concentrations of reactants and products change.

Eventually the system may approach:

equilibrium.

At equilibrium:

Ecell = 0

There is then no net thermodynamic driving force for the reaction in either direction.

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Cell Potential and Batteries

Commercial batteries use carefully selected redox reactions to produce useful:

voltages.

Engineers consider much more than voltage when designing batteries.

Other important factors include:

• energy density
• power output
• lifetime
• rechargeability
• safety
• cost
• operating temperature
• environmental impact

Cell potential is therefore one important property of a battery, but not the only one.

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Cells Connected in Series

Individual electrochemical cells can be connected in:

series.

When cells are connected correctly in series, their voltages approximately add.

For example:

Three 1.5 V cells:

1.5 + 1.5 + 1.5 = 4.5 V

This is why batteries may contain multiple electrochemical:

cells.

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Cell vs Battery

These words are sometimes used interchangeably in everyday speech, but scientifically:

Cell

A single electrochemical unit.

Battery

Technically, a collection of cells connected together, although everyday and commercial terminology often calls single cells "batteries."

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Common Mistake: Adding Reduction Potentials Directly

Suppose:

Cathode:

+0.34 V

Anode reduction potential:

−0.76 V

Do not calculate:

+0.34 + (−0.76) = −0.42 V

Instead use:

E°cell = E°cathode − E°anode

= +0.34 − (−0.76)

= +1.10 V

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Common Mistake: Multiplying Electrode Potentials

If you multiply a half-reaction by 2 to balance electrons:

do not multiply E° by 2.

Electrode potential is an:

intensive property.

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Common Mistake: More Positive Means Anode

For standard reduction potentials, the more positive half-reaction has a stronger tendency to occur as:

reduction.

Therefore it normally forms the:

cathode.

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Common Mistake: Positive Voltage Means Fast Reaction

A positive E°cell indicates that the reaction is thermodynamically favorable under standard conditions.

It does not tell us:

reaction rate.

Thermodynamics and kinetics are different concepts.

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Common Mistake: Voltage and Current Are the Same

They are different quantities.

Voltage = potential difference

Current = rate of charge flow

A cell's voltage does not by itself tell you how much current it can supply.

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Common Mistake: Larger Electrode Means Larger Voltage

Increasing electrode surface area does not automatically increase the equilibrium voltage.

Voltage is primarily determined by the:

electrochemical potential difference between the half-cells and the conditions.

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Common Mistake: Standard Potential Is Always the Actual Voltage

E°cell applies to:

standard-state conditions.

Actual cell voltage can differ because of:

• concentration
• gas pressure
• temperature
• internal resistance and operating conditions

Therefore:

Ecell and E°cell are not always identical.

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A Reliable Problem-Solving Strategy

When solving standard cell-potential problems:

Step 1: Write both reduction half-reactions.

Step 2: Compare their standard reduction potentials.

Step 3: The more positive reduction potential becomes the cathode.

Step 4: The other half-reaction is reversed to represent oxidation at the anode.

Step 5: Balance electrons if you need the overall equation.

Step 6: Calculate:

E°cell = E°cathode − E°anode

Step 7: Check the sign.

If positive:

spontaneous as written under standard conditions.

If negative:

nonspontaneous as written under standard conditions.

Step 8: Never multiply electrode potentials when balancing half-reactions.

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Check Your Understanding

1. Define cell potential.

2. What unit is used to measure cell potential?

3. What symbol is commonly used for cell potential?

4. What does E°cell represent?

5. What is an electrode potential?

6. Why are electrode potentials measured relative to a reference electrode?

7. What is the standard electrode potential of the standard hydrogen electrode?

8. At which electrode does reduction occur?

9. At which electrode does oxidation occur?

10. State the equation used to calculate standard cell potential from standard reduction potentials.

11. Calculate E°cell if E°cathode = +0.60 V and E°anode = −0.40 V.

12. Calculate E°cell if E°cathode = +0.25 V and E°anode = −0.55 V.

13. Which half-reaction normally becomes the cathode: the one with the more positive or more negative reduction potential?

14. Why is the anode reduction potential subtracted?

15. Calculate E°cell for Zn-Cu using +0.34 V and −0.76 V.

16. Calculate E°cell for Mg-Cu using +0.34 V and −2.37 V.

17. Calculate E°cell for Fe-Cu using +0.34 V and −0.44 V.

18. Calculate E°cell for Zn-Ag using +0.80 V and −0.76 V.

19. Which has the greater standard voltage: Zn-Cu or Mg-Cu?

20. Explain why.

21. What does a positive E°cell indicate?

22. What does a negative E°cell indicate?

23. What does Ecell = 0 indicate?

24. If a reaction has E°cell = −0.75 V, what is E°cell for the reverse reaction?

25. Why must electrode potentials not be multiplied when balancing electrons?

26. Explain the difference between voltage and current.

27. Does increasing electrode size automatically increase cell voltage?

28. Name four factors that can affect actual cell potential.

29. How can concentration affect cell voltage?

30. What equation describes the effect of nonstandard conditions on cell potential?

31. Why can battery voltage change as the battery operates?

32. Explain the relationship between E°cell and ΔG°.

33. Does a positive cell potential guarantee a fast reaction? Explain.

34. Which pair produces the larger standard potential difference: Cu-Ag or Mg-Ag?

35. Calculate the standard cell potential for Mg-Ag using −2.37 V and +0.80 V.

36. Write the cell notation for the Mg-Ag galvanic cell.

37. In which direction do electrons move in that cell?

38. Three 1.5 V cells are connected correctly in series. What is the total voltage?

39. Explain why the measured voltage of a real cell may differ from its tabulated E°cell.

40. Explain how electrode potentials can be used to predict both the direction and spontaneity of an electrochemical reaction.

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Key Terms

• Cell potential: Electrical potential difference between the electrodes of an electrochemical cell.
• Voltage: Electrical potential difference, measured in volts.
• Electrode potential: Measure of the tendency of a half-cell to undergo reduction relative to a reference.
• Standard electrode potential: Electrode potential measured relative to the standard hydrogen electrode under standard-state conditions.
• E°cell: Standard cell potential.
• Standard hydrogen electrode (SHE): Reference electrode assigned a standard potential of 0.00 V.
• Anode: Electrode where oxidation occurs.
• Cathode: Electrode where reduction occurs.
• Reduction potential: Electrode potential written for a reduction half-reaction.
• Spontaneous reaction: Thermodynamically favorable reaction under the stated conditions.
• Gibbs free energy: Thermodynamic quantity related to the ability of a process to occur spontaneously.
• Faraday constant: Charge carried by approximately one mole of electrons, about 96,485 C/mol.
• Nernst equation: Equation relating cell potential to standard potential and reaction conditions.
• Reaction quotient (Q): Quantity describing the relative amounts or activities of products and reactants at a particular point.
• Current: Rate of flow of electric charge.
• Voltmeter: Instrument used to measure electrical potential difference.

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Key Takeaways

• Cell potential is the electrical potential difference between the two electrodes of an electrochemical cell.
• Cell potential is measured in volts (V).
• Standard cell potential is represented by E°cell.
• Electrode potentials are measured relative to the standard hydrogen electrode, defined as 0.00 V.
• Standard electrode-potential tables normally list reduction potentials.
• The more positive reduction potential normally becomes the cathode in a spontaneous galvanic cell.
• The less positive reduction potential is reversed and becomes the anode reaction.
• Standard cell potential is calculated using E°cell = E°cathode − E°anode.
• A positive E°cell indicates a spontaneous reaction as written under standard conditions.
• A negative E°cell indicates a nonspontaneous reaction as written under standard conditions.
• At equilibrium, Ecell = 0.
• Electrode potentials are not multiplied when half-equations are multiplied.
• Larger differences between electrode potentials produce larger standard cell voltages.
• Actual cell potential can be affected by concentration, gas pressure, temperature, and reaction conditions.
• The Nernst equation describes cell potential under nonstandard conditions.
• Cell potential is related to Gibbs free energy by ΔG° = −nFE°cell.
• A positive cell potential indicates thermodynamic favorability, not necessarily a fast reaction.
• Voltage and current are different physical quantities.
• Increasing electrode size does not simply increase the equilibrium voltage.
• Comparing electrode potentials allows us to predict which species is oxidized, which is reduced, the direction of electron flow, the cell voltage, and whether the reaction is thermodynamically spontaneous.