Electrochemical Cells Flashcards
Redox Reaction
Donation of electrons from one substance (the reductant) to another (the oxidant)
Displacement Reaction
A type of redox reaction between a more reactive element and a compound containing a less reactive element, resulting the in the less reactive element being displaced from the compound
Oxidant (aka oxidising agent)
A substance that causes another to lose electrons
Reductant (aka reducing agent)
A substance that causes another to gain electrons
Oxidation
The loss of electrons from a substance. Can also be observed as a gain of oxygen, increase in oxidation number or loss of hydrogen.
Reduction
The gain of electrons by a substance. Can also be observed as a loss of oxygen, decrease in oxidation number or gain of hydrogen.
Half-equation
An equation describing either oxidation or reduction occurring in a redox reaction, along with the corresponding loss or gain of electrons
Oxidation Number
A measure of the electron density around an atom, compared to its elemental form. Unlike charge, oxidation numbers have their sign before the value (the sign is always shown)
Conjugant oxidant
A substance produced when a reductant loses electrons, containing the element that has increased in oxidation number.
Conjugant reductant
A substance produced when an oxidant gains electrons, containing the element that has decreased in oxidation number.
Electrochemical Series
A list or table of reduction half-equations written in order of oxidant strength. Usually shows standard electrode potential values for each reduction half-equation.
Standard conditions (redox)
Both reactants and products are present for the relevant half-equation. All solutions present in 1M concentration. All gases present at 1atm pressure. All solids pure. Electrode made of the reductant if it is solid and conductive, otherwise an inert electrode such as graphite or platinum.
Galvanic cell
A chemical system that produces an electrical current (DC) from a spontaneous redox reaction. Made up of two half-cells, which are two separate containers or two parts of a container that are separated by a porous barrier.
Half-cell
An electrode dipped in a solution, where both oxidant and conjugate reductant are present. The electrode may be the reductant or an inert electrode.
Anode
The electrode that is the site of oxidation.
Cathode
The electrode that is the site of reduction
Salt Bridge
A structure containing free-moving unreactive ions, that provides a connection between two half-cells. Often a piece of filter paper soaked in saturated KNO3 solution.
Standard Electrode Potential
The relative strength of the oxidant under standard conditions, compared to the H+/H2 half-cell. Measured as the voltage produced when that half cell is connected as the cathode to the H+/H2 half-cell (which has a standard electrode potential of 0V)
Standard Cell Potential
The voltage (potential difference) across two half-cells under standard conditions. Calculated as E(cat) - E(an), or alternatively, E(oxidant) - E(reductant).
Standard Cell Notation
A written description of a galvanic cell set up under standard conditions, showing both half cells (anode on the left, cathode on the right) separated by a salt bridge (||). Electrodes are shown on either side and phase changes are also shown (|).
Electrolytic Cell
A cell that uses an electric current (DC) to force a chemical reaction to occur that would otherwise not happen. Composed of two electrodes in an electrolyte that is commonly an aqueous solution or molten ionic compound.
Electrolyte
A solution that conducts electricity due to ions being able to flow past each other. May be a molten ionic compound or a dissolved ionic compound.
Electroplating
Using electrolysis to cover (plate) a metal item with another metal in order to improve appearance or reduce corrosion.
Electrorefining
The use of electrolysis and an impure anode to plate a pure metal onto the cathode.
Corrosion
The disintegration of a metal as a result of a redox chemical reaction
Rusting
A specific example of corrosion involving iron.
Sacrificial Anode
A more reactive metal in contact with iron, that oxidises preferentially instead of the iron.
Storage cell
A galvanic cell in which the reactant chemicals are not replenished and will eventually run out.
Oxidation number rules
- Sum = overall charge.
- Fluorine is -1 in compounds.
- Group 1 and 2 metals are +1 and +2 respectively (in compounds)
- H bonded to non-metals is +1, bonded to metals is -1.
- O is -2 in compounds, except for peroxides where it is -1.
- Group 7 are usually -1 in compounds. If >1 halogen is present, the most electronegative is -1.
- Where no other priority exists, Group 6 NMs are usually -2, Group 15 NMs usually -3 and Group 3 metals usually +3.
