Chapter 12 DAT Booster | Booster Prep™ Electrochemistry - Vocabulary Flashcards

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Vocabulary flashcards covering key electrochemistry concepts from Chapter 12 DAT Booster.

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37 Terms

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Oxidation

Loss of electrons in a chemical process; oxidation state increases.

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Reduction

Gain of electrons in a chemical process; oxidation state decreases.

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Redox (oxidation–reduction) reaction

A reaction that involves transfer of electrons between species (one is oxidized and the other reduced).

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Oxidizing agent

Substance that accepts electrons from another species; it is reduced in the process.

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Reducing agent

Substance that donates electrons to another species; it is oxidized in the process.

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Oxidation state (oxidation number)

A bookkeeping value assigned to atoms to track electron transfer in reactions.

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Rule 1: Elements in elemental form

Oxidation number is 0 for atoms in their elemental state.

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Rule 2: Monatomic ions

Oxidation numbers of atoms in monatomic ions equal their charge.

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Rule 3a: Fluorine

Fluorine is always assigned an oxidation number of -1 in compounds.

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Rule 3b: Oxygen

Oxygen is usually -2; exceptions include peroxides (-1) and oxygen bonded to fluorine (+1).

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Rule 3c: Hydrogen

Hydrogen is usually +1; exception: in metal hydrides, hydrogen is -1.

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Rule 4: Alkali metals

Alkali metals (Group 1) have oxidation number +1 in compounds.

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Rule 4: Alkaline earth metals

Alkaline earth metals (Group 2) have oxidation number +2 in compounds.

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Rule 5: Sum of oxidation numbers

The sum of all oxidation numbers in a molecule equals its overall charge; use to solve for unknowns.

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Algebraic determination of oxidation numbers

If some oxidation numbers are unknown, set up an equation using the sum rule and solve.

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SO2 oxidation state (example)

In SO2, the sulfur atom has an oxidation state of +4.

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SO4^2− oxidation state (example)

In sulfate (SO4^2−), sulfur has an oxidation state of +6.

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Half-reaction

One component of a redox reaction that shows either oxidation or reduction (electrons exchanged).

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Oxidation half-reaction

The half-reaction in which oxidation occurs (electrons are produced).

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Reduction half-reaction

The half-reaction in which reduction occurs (electrons are consumed).

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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.

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Balancing Redox Reactions in basic conditions

Seven-step method using OH− to neutralize H+ and form water, then balance as in acidic conditions.

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Anode

Electrode where oxidation occurs; electrons flow from anode to cathode; in galvanic cells the anode is negative, in electrolytic cells it is positive.

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Cathode

Electrode where reduction occurs; electrons flow into the cathode; in galvanic cells the cathode is positive, in electrolytic cells it is negative.

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Galvanic (Voltaic) Cell

Spontaneous redox cell that generates electricity; E°cell > 0; chemical energy converted to electrical energy.

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Electrolytic Cell

Nonspontaneous redox cell driven by an external power source; E°cell < 0; electrical energy converted to chemical energy.

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Salt bridge

Tube filled with nonreactive electrolytes that balances charge and allows ion flow between half-cells.

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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.

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Standard cell potential (E°cell)

Overall cell potential under standard conditions; E°cell > 0 for spontaneous galvanic cells.

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E°reduction

Standard reduction potential of the reduction half-reaction.

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E°oxidation

Standard oxidation potential; equal in magnitude and opposite sign to E°reduction for the corresponding oxidation half-reaction.

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E°cell calculation

E°cell = E°reduction + E°oxidation; when balancing, E°reduction values do not change with coefficients.

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ΔG° = −nFE°cell

Relation between standard cell potential and Gibbs free energy; negative ΔG° indicates spontaneity.

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Faraday’s constant (F)

F = 96,500 C per mole of electrons.

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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.

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Electroplating

Depositing a metal onto a conductive object using an electrolytic cell.

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Molten electrolysis

Electrolysis of a molten salt at high temperature; ions move to electrodes to form elements.