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Vocabulary flashcards covering key electrochemistry concepts from Chapter 12 DAT Booster.
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Oxidation
Loss of electrons in a chemical process; oxidation state increases.
Reduction
Gain of electrons in a chemical process; oxidation state decreases.
Redox (oxidation–reduction) reaction
A reaction that involves transfer of electrons between species (one is oxidized and the other reduced).
Oxidizing agent
Substance that accepts electrons from another species; it is reduced in the process.
Reducing agent
Substance that donates electrons to another species; it is oxidized in the process.
Oxidation state (oxidation number)
A bookkeeping value assigned to atoms to track electron transfer in reactions.
Rule 1: Elements in elemental form
Oxidation number is 0 for atoms in their elemental state.
Rule 2: Monatomic ions
Oxidation numbers of atoms in monatomic ions equal their charge.
Rule 3a: Fluorine
Fluorine is always assigned an oxidation number of -1 in compounds.
Rule 3b: Oxygen
Oxygen is usually -2; exceptions include peroxides (-1) and oxygen bonded to fluorine (+1).
Rule 3c: Hydrogen
Hydrogen is usually +1; exception: in metal hydrides, hydrogen is -1.
Rule 4: Alkali metals
Alkali metals (Group 1) have oxidation number +1 in compounds.
Rule 4: Alkaline earth metals
Alkaline earth metals (Group 2) have oxidation number +2 in compounds.
Rule 5: Sum of oxidation numbers
The sum of all oxidation numbers in a molecule equals its overall charge; use to solve for unknowns.
Algebraic determination of oxidation numbers
If some oxidation numbers are unknown, set up an equation using the sum rule and solve.
SO2 oxidation state (example)
In SO2, the sulfur atom has an oxidation state of +4.
SO4^2− oxidation state (example)
In sulfate (SO4^2−), sulfur has an oxidation state of +6.
Half-reaction
One component of a redox reaction that shows either oxidation or reduction (electrons exchanged).
Oxidation half-reaction
The half-reaction in which oxidation occurs (electrons are produced).
Reduction half-reaction
The half-reaction in which reduction occurs (electrons are consumed).
Balancing Redox Reactions in acidic conditions
Seven-step method: write half-reactions, balance atoms, balance O with H2O, balance H with H+, balance charges with e−, equalize electrons, add half-reactions.
Balancing Redox Reactions in basic conditions
Seven-step method using OH− to neutralize H+ and form water, then balance as in acidic conditions.
Anode
Electrode where oxidation occurs; electrons flow from anode to cathode; in galvanic cells the anode is negative, in electrolytic cells it is positive.
Cathode
Electrode where reduction occurs; electrons flow into the cathode; in galvanic cells the cathode is positive, in electrolytic cells it is negative.
Galvanic (Voltaic) Cell
Spontaneous redox cell that generates electricity; E°cell > 0; chemical energy converted to electrical energy.
Electrolytic Cell
Nonspontaneous redox cell driven by an external power source; E°cell < 0; electrical energy converted to chemical energy.
Salt bridge
Tube filled with nonreactive electrolytes that balances charge and allows ion flow between half-cells.
Electrochemical cell notation
Shorthand cell representation with anode on the left and cathode on the right; a single line separates phases, a double line separates half-cells.
Standard cell potential (E°cell)
Overall cell potential under standard conditions; E°cell > 0 for spontaneous galvanic cells.
E°reduction
Standard reduction potential of the reduction half-reaction.
E°oxidation
Standard oxidation potential; equal in magnitude and opposite sign to E°reduction for the corresponding oxidation half-reaction.
E°cell calculation
E°cell = E°reduction + E°oxidation; when balancing, E°reduction values do not change with coefficients.
ΔG° = −nFE°cell
Relation between standard cell potential and Gibbs free energy; negative ΔG° indicates spontaneity.
Faraday’s constant (F)
F = 96,500 C per mole of electrons.
I × t = nF (electrolysis)
Charge (Q) equals I·t; number of moles of electrons n = Q/F = (I·t)/F; used to compute deposition or electrolysis quantities.
Electroplating
Depositing a metal onto a conductive object using an electrolytic cell.
Molten electrolysis
Electrolysis of a molten salt at high temperature; ions move to electrodes to form elements.