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{{T}}
[[Category:English]]
[[Category:Electrochemical Cells]]

{{BR}}
= Introduction =

'''Electrochemical cells''' convert energy between chemical and electrical forms through [[English:Redox reaction|redox reactions]]. They connect chemistry with electricity, thermodynamics, materials science, batteries, corrosion, electrolysis, and modern energy technology. This aiMOOC is designed for '''Grades 11–13'''. You will move from particle-level explanations to quantitative calculations and evaluation of real electrochemical systems.

[[File:Galvanic Cell Diagram.svg|500px|frameless|center]]

In a working electrochemical cell, oxidation and reduction are physically or functionally separated so that electrons can be directed through an external circuit. The ions in the electrolyte complete the internal charge-transfer pathway. A cell may operate spontaneously and deliver electrical energy, or it may consume electrical energy to force a non-spontaneous reaction.

{{#ev:youtube|https://www.youtube.com/watch?v=inXgVjz1BDQ|500|center}}

{{BR}}
= Learning Goals =

By the end of this course, you should be able to:
# [[English:Redox reaction|Redox reaction]]: Identify oxidation, reduction, oxidizing agents, and reducing agents from electron transfer and oxidation-state changes.
# [[English:Galvanic cell|Galvanic cell]]: Explain the roles of the anode, cathode, electrolyte, salt bridge, and external circuit.
# [[English:Cell potential|Cell potential]]: Calculate standard cell potentials from standard reduction potentials and interpret the sign of the result.
# [[English:Nernst equation|Nernst equation]]: Predict how non-standard concentrations and pressures affect cell potential.
# [[English:Electrolysis|Electrolysis]]: Relate current, time, charge, moles of electrons, and amount of product.
# [[English:Battery|Battery]]: Apply electrochemical ideas to rechargeable cells, fuel cells, corrosion, and energy storage.

{{BR}}
= Redox Foundations =

Every electrochemical cell involves two linked half-reactions. '''Oxidation''' is loss of electrons, while '''reduction''' is gain of electrons. The two processes must occur together because electrons released by one species must be accepted by another.

A useful memory rule is '''OIL RIG''': Oxidation Is Loss, Reduction Is Gain. The rule describes electron transfer, but oxidation states provide a second method. An increase in oxidation state indicates oxidation; a decrease indicates reduction.

Consider the reaction:

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Zinc is oxidized:
Zn(s) → Zn²⁺(aq) + 2e⁻

Copper ions are reduced:
Cu²⁺(aq) + 2e⁻ → Cu(s)

The zinc is the '''reducing agent''' because it donates electrons. Cu²⁺ is the '''oxidizing agent''' because it accepts electrons.

{{BR}}
== Why Separate the Half-Reactions? ==

If zinc metal is placed directly into a copper-ion solution, electron transfer can occur at the metal-solution interface and much of the released chemical energy becomes thermal energy. If the oxidation and reduction processes are separated into half-cells and linked by a wire, electrons can travel through an external circuit and perform electrical work. This separation is the central design idea of many electrochemical cells.

{{BR}}
= Anatomy of an Electrochemical Cell =

A typical cell contains two '''electrodes''' and one or more '''electrolytes'''. An electrode is an electronic conductor at which a half-reaction occurs. The electrolyte contains mobile ions that carry charge through the liquid, gel, molten salt, or solid ionic conductor.

The '''anode''' is always the electrode where oxidation occurs. The '''cathode''' is always the electrode where reduction occurs. These reaction-based definitions never change.

In a galvanic cell, the anode is negative and the cathode is positive because the spontaneous reaction pushes electrons from anode to cathode through the external circuit. In an electrolytic cell, an external power source forces electron transfer, so the anode is positive and the cathode is negative.

A '''salt bridge''' or porous separator can connect two liquid half-cells. Its key role is not to transport electrons. Instead, mobile ions move so that large charge imbalances do not build up in either half-cell. In a simple Daniell cell, anions migrate toward the anode compartment, where positive ions are being produced, while cations migrate toward the cathode compartment, where positive ions are being consumed.

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= Galvanic and Voltaic Cells =

A [[English:Galvanic cell|galvanic cell]], also called a voltaic cell, uses a thermodynamically favorable redox reaction to generate electrical energy. A classic example is the [[English:Daniell cell|Daniell cell]], based on zinc and copper.

At the zinc anode:
Zn(s) → Zn²⁺(aq) + 2e⁻

At the copper cathode:
Cu²⁺(aq) + 2e⁻ → Cu(s)

Overall:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Electrons move through the external wire from the zinc anode to the copper cathode. Conventional current in the wire is defined in the opposite direction.

[[File:EC 3.1 Daniell Cell.webm|500px|center]]

{{BR}}
== Cell Notation ==

Electrochemists use shorthand notation to describe cells. A vertical line represents a phase boundary, and a double vertical line represents the junction between half-cells, often a salt bridge.

For the Daniell cell:
Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

By common convention, the anode half-cell is written on the left and the cathode half-cell on the right for a galvanic cell. Species in the same phase can be separated by commas. An inert conductor such as platinum may be included when no reacting solid conductor is present.

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== What You Can Observe in a Daniell Cell ==

As the cell discharges, zinc atoms enter solution as Zn²⁺, so the zinc electrode tends to lose mass. Cu²⁺ ions gain electrons and deposit as copper metal, so the copper electrode tends to gain mass. The concentrations of dissolved species change with time, and the cell potential eventually changes as the system moves toward equilibrium.

{{BR}}
= Measuring and Predicting Cell Potential =

Electrical potential difference is measured in volts. One volt is one joule per coulomb. In electrochemistry, the cell potential indicates the driving force per unit charge for the redox process.

A single half-cell potential cannot be measured in isolation. Electrode potentials are therefore defined relative to a reference electrode. The most important reference is the [[English:Standard hydrogen electrode|standard hydrogen electrode]], whose standard electrode potential is assigned a value of 0 V.

[[File:Standard hydrogen electrode 2009-02-06.svg|500px|frameless|center]]

Standard reduction potentials are tabulated for half-reactions written as reductions. Under standard-state conventions, solute activities are 1, gases have a standard pressure of 1 bar, and pure solids and liquids have activity 1. In school calculations, an aqueous activity of 1 is often approximated by a concentration of 1 mol dm⁻³, and data tables are commonly given for 298.15 K.

{{BR}}
== Standard Cell Potential ==

When both tabulated values are written as reduction potentials:

'''E°cell = E°cathode − E°anode'''

For the Daniell cell:
Cu²⁺ + 2e⁻ → Cu, E° = +0.34 V
Zn²⁺ + 2e⁻ → Zn, E° = −0.76 V

Therefore:
E°cell = +0.34 V − (−0.76 V) = +1.10 V

Do not multiply an electrode potential by a stoichiometric coefficient. Potential is an intensive quantity. You balance electrons in the half-equations, but the E° values themselves are not scaled.

{{BR}}
== Interpreting the Sign ==

For a cell reaction written in the forward direction, a positive E°cell corresponds to a negative standard Gibbs free-energy change and therefore a thermodynamically favorable reaction under standard-state conditions.

The relationship is:
'''ΔG° = −nFE°cell'''

Here, n is the number of moles of electrons transferred per mole of the balanced cell reaction and F is the Faraday constant, approximately 96485 C mol⁻¹.

At equilibrium, ΔG = 0 and there is no net thermodynamic driving force. Standard cell potential is also connected with the equilibrium constant:
'''ln K = nFE°cell / RT'''

A large positive E°cell therefore generally corresponds to a large equilibrium constant for the reaction as written.

{{BR}}
= Non-Standard Conditions and the Nernst Equation =

Real cells rarely operate with every species in its standard state. The [[English:Nernst equation|Nernst equation]] connects cell potential with composition:

'''Ecell = E°cell − (RT / nF) ln Q'''

Q is the reaction quotient for the balanced cell reaction. At 298.15 K, a commonly used base-10 form is:

'''Ecell = E°cell − (0.05916 V / n) log Q'''

Pure solids and pure liquids are omitted from Q because their activities are taken as 1. Dissolved species and gases can influence Q and therefore the cell potential.

For the Daniell cell:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

an approximate concentration-based reaction quotient is:
'''Q = [Zn²⁺] / [Cu²⁺]'''

If Q increases, the logarithmic correction becomes more positive and Ecell decreases. This matches the chemical idea that the cell loses driving force as products accumulate relative to reactants.

{{#ev:youtube|https://www.youtube.com/watch?v=WqsJNtBEKtg|500|center}}

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== Concentration Cells ==

A concentration cell can use the same redox couple in both half-cells but at different activities. The potential comes from the tendency of the concentration difference to decrease. As the two sides approach equilibrium, the cell potential approaches zero. Concentration cells show clearly that voltage can arise from a chemical-potential difference even when the electrode materials are identical.

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= Electrolytic Cells =

An [[English:Electrolytic cell|electrolytic cell]] uses electrical energy from an external source to drive a non-spontaneous redox process. The definitions of the electrodes remain unchanged: oxidation occurs at the anode and reduction occurs at the cathode.

In an electrolytic cell:
'''anode = oxidation = positive electrode'''
'''cathode = reduction = negative electrode'''

This polarity is opposite to a galvanic cell because the external power supply pulls electrons away from the anode and pushes electrons toward the cathode.

[[File:Electrolytic Cell Diagram.jpg|500px|frameless|center]]

Electrolytic cells are used in metal extraction, metal refining, electroplating, chemical manufacture, and charging some rechargeable batteries. The actual products of electrolysis depend on the electrolyte, electrode materials, concentrations, and applied potential. In aqueous systems, water can compete with dissolved ions for oxidation or reduction, so simply choosing the ion with the most favorable tabulated potential may not always predict the observed product.

{{BR}}
== Quantitative Electrolysis and Faraday's Law ==

Electrical charge is related to current and time:
'''Q = It'''

where Q is charge in coulombs, I is current in amperes, and t is time in seconds.

The amount of electrons transferred is:
'''n(e⁻) = Q / F'''

If formation of one mole of product requires z moles of electrons, then:
'''n(product) = Q / zF'''

For a product with molar mass M, the ideal deposited mass is:
'''m = ItM / zF'''

These relationships let you predict the mass of metal deposited during electroplating or the amount of gas produced during electrolysis. Real systems may have less than 100 percent current efficiency if side reactions also consume charge.

[[File:Electrolysis diagram.png|500px|frameless|center]]

{{BR}}
= Practical Investigation: Building and Testing Cells =

A useful school investigation compares cell voltages produced by different metal and ion combinations. You can construct microscale half-cells, connect them with a salt bridge or electrolyte-soaked paper, and measure open-circuit voltage with a high-resistance voltmeter.

{{#ev:youtube|https://www.youtube.com/watch?v=igpwN3m_4Sw|500|center}}

A good investigation controls electrode surface condition, solution concentration, temperature, immersion depth, salt-bridge composition, and measurement time. Repeat measurements and record uncertainty. Compare the observed voltage with the value predicted from electrode potentials.

'''Safety:''' Use eye protection and follow your teacher's or laboratory's risk assessment. Use small quantities of solutions, avoid skin contact, do not taste chemicals, wash hands after practical work, and dispose of metal-ion solutions as instructed. Do not connect school cells to mains electricity. For electrolysis, use only approved low-voltage power supplies and suitable electrodes.

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== Explaining Differences Between Theory and Experiment ==

Measured voltages can differ from ideal predictions. Possible causes include non-standard concentrations, temperature differences, contaminated or oxidized electrode surfaces, liquid-junction potentials, internal resistance, polarization, and measurement uncertainty.

Thermodynamics also does not tell you how fast a reaction will occur. A reaction can be thermodynamically favorable but kinetically slow. '''Overpotential''' is the extra applied potential often needed to drive an electrode reaction at a practical rate. '''Passivation''' occurs when a surface layer slows further reaction. These ideas are especially important in corrosion, electrolysis, and battery engineering.

{{BR}}
= Batteries and Other Applications =

A [[English:Battery|battery]] contains one or more electrochemical cells. During discharge, a battery operates galvanically: spontaneous redox reactions provide electrical energy. A rechargeable battery can be driven in the reverse direction during charging, so the device then operates electrolytically.

[[File:Lithium ion battery.png|500px|frameless|center]]

In a lithium-ion cell, lithium ions move through an electrolyte between host materials while electrons move through the external circuit. The exact electrode reactions depend on the cell chemistry. Good battery design must balance cell voltage, capacity, power, cycle life, safety, cost, and environmental impact.

[[File:Lead aicd battery.png|500px|frameless|center]]

Lead-acid cells use lead, lead dioxide, and sulfuric acid and remain important for vehicle starting and backup power. Their chemistry is rechargeable, but lead is toxic and requires controlled recycling.

{{BR}}
== Fuel Cells ==

A [[English:Fuel cell|fuel cell]] is an electrochemical cell in which reactants are supplied continuously from outside the device. In a hydrogen fuel cell, hydrogen is oxidized and oxygen is reduced; the overall chemical product is water under typical operation. Fuel cells can convert chemical energy to electrical energy without a combustion step, but their environmental impact depends strongly on how the fuel is produced and transported.

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== Corrosion as an Electrochemical Process ==

[[English:Corrosion|Corrosion]] often involves coupled anodic and cathodic reactions on a metal surface. In iron corrosion, some regions of the metal are oxidized while reduction reactions occur elsewhere. Water and dissolved ions can provide ionic conduction. Protective coatings, sacrificial anodes, impressed-current systems, alloying, and careful materials selection are used to reduce corrosion.

Understanding corrosion requires both thermodynamics and kinetics. Electrode potentials can indicate whether a process is thermodynamically possible, but rates depend on surface condition, transport, oxygen availability, pH, and other factors.

{{BR}}
= Comparing Galvanic and Electrolytic Cells =

{| class="wikitable"
! Feature
! Galvanic cell
! Electrolytic cell
|-
| Energy conversion
| Chemical energy to electrical energy
| Electrical energy to chemical change
|-
| Overall reaction
| Thermodynamically favorable in the operating direction
| Driven in a non-spontaneous direction by an external source
|-
| Anode process
| Oxidation
| Oxidation
|-
| Cathode process
| Reduction
| Reduction
|-
| Anode polarity
| Negative
| Positive
|-
| Cathode polarity
| Positive
| Negative
|-
| Typical example
| Discharging battery
| Electroplating cell
|}

The safest way to identify an electrode is by the reaction, not by the sign. Remember: '''anode means oxidation''' and '''cathode means reduction''' in every electrochemical cell.

{{BR}}
= Common Misconceptions =

'''Misconception: electrons travel through the salt bridge.''' They do not. Electrons travel through the external electronic conductor; ions carry charge through the electrolyte and salt bridge.

'''Misconception: the anode is always negative.''' The anode is negative in a galvanic cell but positive in an electrolytic cell.

'''Misconception: the larger standard potential must be multiplied when a half-equation is multiplied.''' Electrode potentials are not multiplied by stoichiometric coefficients.

'''Misconception: a positive standard cell potential guarantees a fast reaction.''' It indicates favorable thermodynamics under the stated conditions, not a fast rate.

'''Misconception: a salt bridge is just a wire for ions.''' Its ion motion maintains approximate electroneutrality and closes the ionic circuit while limiting bulk mixing between half-cell solutions.

{{BR}}
= Worked Example =

A galvanic cell uses the half-reactions:
Ag⁺ + e⁻ → Ag, E° = +0.80 V
Cu²⁺ + 2e⁻ → Cu, E° = +0.34 V

The more positive reduction potential is the silver half-reaction, so silver ions are reduced at the cathode. Copper is oxidized at the anode.

Balanced half-reactions:
2Ag⁺ + 2e⁻ → 2Ag
Cu → Cu²⁺ + 2e⁻

Overall:
Cu + 2Ag⁺ → Cu²⁺ + 2Ag

The standard cell potential is:
'''E°cell = +0.80 V − +0.34 V = +0.46 V'''

The positive value shows that the balanced reaction as written is thermodynamically favorable under standard-state conditions. Notice that the silver potential remains +0.80 V even though its half-reaction was multiplied by two.

{{BR}}
= Interactive Tasks =

{{BR}}
== Quiz: Test Your Knowledge ==

{{MC}}

'''At which electrode does oxidation always occur?'''
(Anode)
(!Cathode)
(!Salt bridge)
(!Voltmeter)

{{E}}
<br>

{{MC}}

'''What is the main role of a salt bridge in a two-half-cell galvanic cell?'''
(Maintain charge balance by ion movement)
(!Carry electrons between electrodes)
(!Increase the electrode surface area)
(!Supply electrical power)

{{E}}
<br>

{{MC}}

'''In a Daniell cell, in which direction do electrons move through the external wire?'''
(From zinc anode to copper cathode)
(!From copper cathode to zinc anode)
(!From salt bridge to zinc electrode)
(!From electrolyte to salt bridge)

{{E}}
<br>

{{MC}}

'''Which expression gives the standard cell potential from standard reduction potentials?'''
(E cathode minus E anode)
(!E anode minus E cathode)
(!E cathode plus E anode)
(!E anode divided by E cathode)

{{E}}
<br>

{{MC}}

'''What happens to the zinc electrode in a discharging Daniell cell?'''
(It tends to lose mass)
(!It gains copper metal)
(!It becomes the cathode)
(!It receives electrons from copper)

{{E}}
<br>

{{MC}}

'''What does a positive standard cell potential indicate for the reaction as written?'''
(It is thermodynamically favorable under standard conditions)
(!It must react rapidly)
(!It requires an external power supply)
(!It has zero Gibbs free energy change)

{{E}}
<br>

{{MC}}

'''Which quantity appears in the Nernst equation to represent mixture composition?'''
(Reaction quotient)
(!Activation energy)
(!Molar mass)
(!Specific heat capacity)

{{E}}
<br>

{{MC}}

'''What is the polarity of the anode in an electrolytic cell?'''
(Positive)
(!Negative)
(!Always neutral)
(!Alternating)

{{E}}
<br>

{{MC}}

'''Which relationship connects charge, current, and time?'''
(Q equals I times t)
(!Q equals I divided by t)
(!Q equals F times voltage)
(!Q equals resistance times time)

{{E}}
<br>

{{MC}}

'''Why are standard reduction potentials not multiplied by stoichiometric coefficients?'''
(Potential is an intensive quantity)
(!Electrons cannot be balanced)
(!Half-reactions must contain one electron)
(!Voltage has no units)

{{E}}
<br>

{{BR}}
== Memory Game ==

<div class="memo-quiz">
{|
|-
| Anode || Electrode where oxidation occurs
|-
| Cathode || Electrode where reduction occurs
|-
| Electrolyte || Ionic medium that conducts charge
|-
| Salt bridge || Ionic connection that helps maintain electroneutrality
|-
| Cell potential || Electrical driving force per unit charge between electrodes
|-
| Nernst equation || Relationship between cell potential and reaction quotient
|-
| Faraday constant || Charge carried by one mole of electrons
|}
{{E}}
<br>

{{BR}}
== Drag and Drop ==

<div class="lueckentext-quiz">

{| class="wikitable"
! Match the correct terms.
! Topic
|-
| '''Oxidation'''
| Loss of electrons
|-
| '''Reduction'''
| Gain of electrons
|-
| '''Galvanic cell'''
| Spontaneous redox reaction produces electrical energy
|-
| '''Electrolytic cell'''
| External electrical energy drives chemical change
|-
| '''Salt bridge'''
| Ion movement maintains approximate charge balance
|}
{{E}}

<br>
...
<br />

{{BR}}
== Crossword Puzzle ==

<div class="kreuzwort-quiz">
{|
|-
| Anode || Which electrode is the site of oxidation?
|-
| Cathode || Which electrode is the site of reduction?
|-
| Electrolyte || What ionic medium conducts charge inside a cell?
|-
| Oxidation || What process involves loss of electrons?
|-
| Reduction || What process involves gain of electrons?
|-
| Voltmeter || What instrument measures potential difference?
|}
{{E}}
<br>

{{BR}}
== LearningApps ==

<iframe> https://learningapps.org/index.php?s=Electrochemical+Cells </iframe>

{{BR}}
== Cloze Text ==

<quiz display=simple>
{'''Complete the text.'''<br>
|type="{}"}
An electrochemical cell couples oxidation and { reduction }. Oxidation always occurs at the { anode }. Electrons move through the external circuit toward the { cathode }. A salt bridge completes the ionic circuit by allowing { ions } to move. In a galvanic cell, a spontaneous reaction generates { electrical } energy. Standard cell potential is found by subtracting the anode reduction potential from the { cathode } reduction potential. The reference assigned a standard potential of zero volts is the standard { hydrogen } electrode. The Nernst equation uses the reaction quotient { Q } to describe non-standard composition. During electrolysis, charge is calculated from current multiplied by { time }. One mole of electrons carries approximately one { Faraday } of charge.
</quiz>

<br>

{{BR}}
= Open-Ended Tasks =

{{BR}}
=== Easy ===
# [[English:Electrode identification|Electrode identification]]: Draw a galvanic cell and label the anode, cathode, electron direction, ion path, and the oxidation and reduction half-reactions.
# [[English:Redox storyboard|Redox storyboard]]: Create a six-frame visual story that follows one electron from a zinc atom at the anode to a copper ion at the cathode.
# [[English:Cell vocabulary interview|Cell vocabulary interview]]: Interview a classmate about anode, cathode, electrolyte, and salt bridge, then write a short correction of any misconception you discover.
# [[English:Battery audit|Battery audit]]: Photograph or list five battery-powered devices in your environment and classify the likely battery role as primary, rechargeable, or built-in rechargeable.

{{BR}}
=== Standard ===
# [[English:Daniell cell investigation|Daniell cell investigation]]: Build or observe a teacher-approved microscale zinc-copper cell, measure its voltage, repeat the measurement, and explain differences from the predicted standard value.
# [[English:Electrolysis data study|Electrolysis data study]]: Use a safe teacher-provided electrolysis setup to measure current and time, predict the amount of product from Faraday's law, and compare prediction with observation.
# [[English:Concentration and voltage|Concentration and voltage]]: Design an investigation that changes one ion concentration while controlling other variables, then predict and graph how cell potential should respond.
# [[English:Battery explainer video|Battery explainer video]]: Produce a two-minute video explaining how a rechargeable cell switches between galvanic behavior during discharge and electrolytic behavior during charging.

{{BR}}
=== Advanced ===
# [[English:Nernst modeling project|Nernst modeling project]]: Build a spreadsheet or program that calculates cell potential from E°cell, temperature, electron number, and reaction quotient, then test it with at least three scenarios.
# [[English:Corrosion field study|Corrosion field study]]: Visit or inspect a safe local metal structure, document evidence of corrosion or protection, and explain the observations using coupled anodic and cathodic processes.
# [[English:Battery technology comparison|Battery technology comparison]]: Compare lithium-ion and lead-acid cells using chemistry, voltage, energy storage, cycle life, safety, recycling, and material-resource considerations, then defend a choice for one application.
# [[English:Electrochemical design challenge|Electrochemical design challenge]]: Propose a cell for a realistic task, justify the half-reactions and materials, estimate voltage, identify kinetic or safety limitations, and present your design as a poster or short research pitch.

{{:Open Task - Create a MOOC}}

{{BR}}
= Learning Assessment =

# [[English:Cell mechanism analysis|Cell mechanism analysis]]: Given a new cell diagram and half-reactions, determine the anode and cathode, trace electron and ion movement, and justify every label from the reaction chemistry.
# [[English:Potential calculation and interpretation|Potential calculation and interpretation]]: Calculate E°cell from a table of reduction potentials, balance the overall equation, determine n, and explain what the sign of E°cell means thermodynamically.
# [[English:Non-standard transfer|Non-standard transfer]]: Predict qualitatively and quantitatively how changing one reactant concentration affects Q and Ecell, then connect the result with Le Châtelier's principle and the Nernst equation.
# [[English:Electrolysis planning|Electrolysis planning]]: Determine the current or time needed to produce a specified amount of substance and evaluate why the actual yield might be lower than the ideal prediction.
# [[English:Experimental evaluation|Experimental evaluation]]: Analyze a set of measured cell voltages, identify systematic and random sources of error, and propose a revised method that would improve validity and reliability.
# [[English:Technology decision|Technology decision]]: Evaluate two electrochemical technologies for a stated application and justify a recommendation using chemical, energetic, environmental, safety, and economic evidence.

<br>
<br>

{{BR}}
= Evidence of Learning =

{| class="wikitable"
! Evidence type
! What strong evidence looks like
|-
| Knowledge
| You correctly explain redox processes, electrode roles, salt-bridge function, cell notation, standard potentials, the Nernst equation, and Faraday's law.
|-
| Quantitative skill
| You balance half-reactions, calculate E°cell, use ΔG° = −nFE°cell, apply the Nernst equation, and connect Q = It with electrolysis stoichiometry.
|-
| Practical skill
| You construct or interpret a safe cell setup, make repeatable voltage or current measurements, control variables, record uncertainty, and evaluate sources of error.
|-
| Product
| You create a diagram, model, report, spreadsheet, poster, or video that communicates electrochemical reasoning accurately.
|-
| Transfer
| You use electrochemical principles to explain an unfamiliar battery, corrosion problem, concentration effect, fuel cell, or industrial electrolysis process.
|-
| Evaluation
| You distinguish thermodynamic feasibility from reaction rate and discuss practical limitations such as resistance, overpotential, passivation, efficiency, safety, and materials impacts.
|}

<br>
<br>

{{BR}}
= OERs on the Topic =

<iframe> https://en.m.wikipedia.org/wiki/Electrochemical_cell </iframe>

<br>

For further study, use these openly accessible resources:
# [https://openstax.org/books/chemistry-2e/pages/17-2-galvanic-cells OpenStax Chemistry 2e: Galvanic Cells]: Detailed explanations, diagrams, and worked examples.
# [https://openstax.org/books/chemistry-2e/pages/17-3-electrode-and-cell-potentials OpenStax Chemistry 2e: Electrode and Cell Potentials]: Standard electrode potentials and cell-voltage calculations.
# [https://openstax.org/books/chemistry-2e/pages/17-4-potential-free-energy-and-equilibrium OpenStax Chemistry 2e: Potential, Free Energy, and Equilibrium]: Nernst equation and thermodynamic relationships.
# [https://edu.rsc.org/resources/redox-equilibria-16-18/4010292.article Royal Society of Chemistry: Redox Equilibria 16–18]: Practice material on electrode potentials and electrochemical cells.
# [https://commons.wikimedia.org/wiki/Category:Electrochemical_cells Wikimedia Commons: Electrochemical cells]: Freely licensed diagrams and images for review and projects.

{{BR}}
= Linked Learning Areas =

{| align=center
{{:D-Tab}}
'''[[English:Electrochemical Cells|Electrochemical Cells]]'''
# [[English:Redox reaction|Redox reaction]]
# [[English:Oxidation state|Oxidation state]]
# [[English:Galvanic cell|Galvanic cell]]
# [[English:Electrolytic cell|Electrolytic cell]]
# [[English:Standard electrode potential|Standard electrode potential]]
# [[English:Nernst equation|Nernst equation]]
# [[English:Electrolysis|Electrolysis]]
# [[English:Battery|Battery]]
# [[English:Fuel cell|Fuel cell]]
# [[English:Corrosion|Corrosion]]
# [[English:Thermodynamics|Thermodynamics]]
# [[English:Chemical equilibrium|Chemical equilibrium]]
|}

The topic connects chemistry with [[English:Physics|physics]] through electric potential and current, with [[English:Mathematics|mathematics]] through logarithms and quantitative modeling, with [[English:Materials science|materials science]] through electrode and electrolyte design, and with [[English:Environmental science|environmental science]] through resource use, corrosion control, battery recycling, and energy systems.

[[Category:English]]
[[Category:Chemistry]]
[[Category:Electrochemistry]]
[[Category:Science]]
[[Category:Grades 11-13]]

{{BR}}
= aiMOOC Projects =
[[Category:AI_MOOC]] [[Category:GPT aiMOOC]]
{{MT}}

English:Electrochemical Cells - Versionsgeschichte

Glanz: aiMOOC über GPT aiMOOC Action erstellt