Electrochemical Cells Notes

Electrochemical cells split redox reactions into half reactions, placing the oxidant and reductant in separate compartments (cells).
This prevents direct contact and reaction but allows charge transfer through an external circuit (electrons) and a salt bridge (counter ions).
Single displacement reactions (metal displacing cation of ionic compound) are relevant.
Example: Zinc and copper reactions.
Zinc is more active than copper (from activity series), so metallic zinc spontaneously dissolves in copper (II) ion solutions.
- $Zn(s) + Cu^{+2} → Cu(s) + Zn^{+2}$ is spontaneous (negative free energy, ΔG < 0).
- $Cu(s) + Zn^{+2} → Zn(s) + Cu^{+2}$ is non-spontaneous (positive free energy, ΔG > 0).
Spontaneous reactions can be used to build a battery (Daniell cell, by John Frederic Daniell).
Electrons flow from the reductant (zinc) to the oxidant (copper ions) through an external circuit.
Salt bridge prevents copper ions from migrating to the zinc compartment.

Placing zinc metal in copper(II) solution results in a spontaneous redox reaction (\Delta G < 0).
No reaction occurs if copper metal is placed in zinc solution (\Delta G > 0).
Electrochemical cells transfer electrons spontaneously in the direction of \Delta G < 0 in galvanic or voltaic mode.
Galvanic/voltaic cells provide energy to do work (batteries).
Nonspontaneous reactions can be driven by coupling to a spontaneous reaction or by applying an external power supply (electrolytic mode).
In electrolytic mode, the electrochemical cell is driven by an external energy source.

Oxidation occurs at the anode, and reduction at the cathode - always true.
In a galvanic cell, electrons flow from the anode through a load to the cathode (+).
In an electrolytic cell, an external power supply drives the reaction; the sign of electrodes is defined by the power supply.
The electrolytic cell becomes like the load on the galvanic cell.
The external power supply produces a negative free energy in the nonspontaneous direction, greater in magnitude than the spontaneous free energy of the galvanic cell.

Electrochemical cells split oxidant and reductant, allowing electron flow through an external circuit (from reductant to oxidant) while preventing physical contact.
Two compartments: anode and cathode.
Each compartment must have an electrode to connect to the external circuit.

Oxidation occurs at the anode.
Sign switches between galvanic and electrolytic modes.
Galvanic cell: negative electrode. From the reaction perspective, the electrode gains electrons as the reductant loses them.
Electrolytic cell: positive electrode. The sign is determined by the external power supply.

Reduction occurs at the cathode.
The cathode has the opposite signs of the anode. It is [+] in a galvanic/voltaic cell and [-] in an electrolytic cell.

Connect the wire to physically connect the wire to external circuit.
Active electrode: Takes part in the half-reaction; must be a solid.
- In the Daniell cell (Zn/Cu), both electrodes are active.
- The anode loses mass (Zn becomes $Zn^{+2}$ ), while the cathode gains mass (Cu $^{+2}$ becomes Cu(s)).
Inert electrode: Does not undergo oxidation or reduction.
- Used when oxidants or reductants are gases or solutes (e.g., hydrogen half-reaction: $2H^+(aq) + 2e^- ⇋ H_2(g)$ ).
- Requires an additional substance like platinum or graphite.

A charged mobile ion functions as a conducting medium.
Typically an aqueous solution of ionic compounds ("wet cells").
For active electrodes, use a soluble salt with a common ion that also makes a soluble salt with the electrode metal.
Example: Copper(II) sulfate is soluble so sodium sulfate would be a good electrolyte for the copper half cell.
Sodium phosphate would not work because copper(II) phosphate would precipitate out.

Completes the circuit and prevents charge buildup, which would stop electron flow.
Enables an internal circuit allowing ions to complete the circuit.
Prevents the oxidant (Cu $^{+2}$ ) from migrating to the anode and being directly reduced.
Functions as a "semipermeable membrane," allowing electrolyte ions to migrate but preventing redox reaction ions from migrating.
In reality, some redox reaction ions do migrate through the salt bridge.

Shorthand for describing an electrochemical cell.
Anode on the left, cathode on the right.
Reactants and products separated by a vertical line (phase boundary).
Need to have a solid electrode on the very left of the anode and very right of the cathode.
Does not represent balanced equations, but shows what is oxidized and reduced.
Example: Zinc-copper cell operating in galvanic mode.
$Zn(s)|Zn^{+2}(aq)||Cu^{+2}(aq)|Cu(s)$
At standard one molar conditions:
$Zn(s)|Zn^{+2}(1M)||Cu^{+2}(1M)|Cu(s)$
If operating in electrolytic mode:
$Cu(s)|Cu^{+2}(aq)||Zn^{+2}(aq)|Zn(s)$

Inert electrode (Pt) added to the cathode half-reaction because no solid is involved in the reduction half reaction.
Comma separates $Fe^{+3}$ and $Fe^{+2}$ because they are in the same phase.
Cell notation does not represent a balanced equation; it identifies what is oxidized and reduced.

Electrochemical cells used to do work.
Uses spontaneous (Zn/Cu) reaction.
With 1M $Zn^{+2}$ and $Cu^{+2}$ ion concentrations, a voltmeter reads 1.10 volts (open circuit potential, no current flowing).
When current flows, [ $Zn^{+2}$ ] increases, and [ $Cu^{+2}$ ] decreases.
Zinc anode dissolves, and copper ions attach to the cathode and are reduced to neutral copper.
The electrodes are active electrodes because they take place in the Redox reaction.
Measuring the mass change of either the anode or cathode over time determines the number of electrons transferred and calculates the average current.

The word "battery" is commonly used to mean a galvanic (or voltaic) cell that can be used to power electrical devices, but that is technically incorrect and a battery can be galvanic or electrolytic.
The term battery originally came from a series of cannons, and a battery of electrochemical cells represent a series of cells that are linked in series.
A 12-volt car battery is six 2-volt cells linked in series.
Electronic Schematic Symbol of a battery implies cells connected in series.
Diagrams show current flowing through an external circuit from the left electrode to the right electrode, while in reality, electrons are flowing from the right electrode (anode, where oxidation occurs) to the left electrode (cathode, where reduction occurs).