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Gr. 12 Electrochemistry Cheat Sheet by

Redox Reactions

Redox = "­r­edu­ction and oxid­ati­on­"
Oxida­tion: lose electrons (e¯)
(Classical: any reaction involving oxygen gas (O2) - rusting, combus­tion, etc.)
Reduc­tion: gain electrons (e¯)
(Classical: reducing a metal ore into pure metal)
Both reactions always occur together
Reducing agent: substance that causes another substa­nce to become redu­ced
Oxidizing agent: substance that causes another substa­nce to become oxid­ized
Reducing agents NEVER reduce themse­lves; is always oxid­ized to promote reduction
Oxidizing agents NEVER oxidize themse­lves; is always redu­ced to promote oxidation
Reduci­ng/­oxi­dizing agents are ALWAYS reactants
Remember: LEO the lion says GER
Lose El­ectrons = Ox­idation
Gain El­ectrons = Re­duction

Not new reacti­ons: many synthesis, decomp­osi­tion, combustion and single displa­cement reactions are often redox reactions

Oxidation States­/Ox­idation Numbers

Net charge that an atom would have if the e¯ pairs in covalent bonds belonged enti­rely to the more electr­one­gative ion
All redox reactions require a change in O.N.
O.N. = oxidation (loss of e¯)
O.N. = reduction (gain e¯)
Rules to determine oxidation numbers (O.N.)
Pure elements have O.N. = 0
C(s) = 0, O2(s) = 0, P4(s) = 0
Monoatomic ions have O.N. = their charge
Al3+ = +3, Cl¯ = -1
Hydrogen always has O.N. = +1 (except metal hydrides = -1)
HCl (H = +1); H2S (H = +1); CaH2 (H = -1)
Oxygen always has O.N. = -2 (except peroxides = -1)
Li2O (O = -2); KNO3 (O = -2); H2O2 (O = -1)
In a compound, groups I, II, and IV usually have O.N. = ionic charge
NaCl (Na = +1, Cl = -1); MgO (Mg = +2)
In a neut­ral compou­nd1, ΣO.N. = 0
CF4 (C = +4, F = -1)
In a polyatomic ion, ΣO.N. = ion's charge
NO3¯ (N = 5, O = -2; Σ O.N. = -1)
In molecular compunds with no O or H, the more electr­one­gative atom has O.N. = its usual charge
CS2 (S = -2), Li3N (N = -3)                                                                             
Note: The atoms do not actually have these charges!

[1] If a compound contains a polyatomic ion, the charge on the other ion is the oppo­site to the poly­atomic ion's charge (ex. KIO3 - K = +1 because IO3 is 1-)

Format for O.N.: "­±#" (not "­#±" - ionic charges)

Half-R­eac­tions and Dispro­por­tionate Reactions

Most often, one atom is reduced and another is oxidized, but sometimes the same atom can be oxidized and reduced in the same redox reaction
Cu2O(aq) + H2SO­4(aq) Cu(s) + CuSO4 (aq) + H2O­(l)
Cu +1 to 0 gain 1 e¯ redu­ced
Cu +1 to +2 lose 1 e¯ oxid­ized
Since both of these happen in the same reacti­on, it is disp­rop­ort­ion­ate
Half-reaction: reactions made from overall net ionic equations that focus on 1 specific atom
Zn(s) + CuSO4­(aq) ZnSO4­(aq) + Cu(s)
Zn(s) ZnSO4­(aq)
CuSO­4(aq) Cu(s)

Electr­och­emical Cells

Two types of electr­och­emical cells:
Galvanic cells
arrang­ement of 2 connected half-­cells that spon­tan­eou­sly produce an electric current; e¯ always flow from high potential low potential
Elect­rolytic cells
arrang­ement of 2 connected half-­cells that uses electrical energy to produce a non-­spo­nta­neous electric current; e¯ always flow from low potential high potential
Parts of an electr­och­emical cell
Electrolyte: solution that contains aqueous ions (cat­ions (+) and anions (-))
Elect­rode: solid metal conductor where redox reactions occur (cat­hode (oxida­tion) and anode (reduc­tion))
Salt bridge: tube that contains an electr­olyte solution and connects the 2 half-c­ells; used to maintain electrical neutra­lity
During the lifespan of the cell, the anode decr­eases in mass, while the cathode incr­eases in mass

Cell Potential

The measure of the electric potential differ­ence (voltage) between 2 half-cells
Standard cell: galvanic cell in which all entities are at SATP and all solution concen­tra­tions are 1.0 mol/L
Standard cell potential (ΔE°): the ability of each half cell to gain e¯ (red­uct­ion)
Cell potential formula:
ΔE°(cell) = E°cat­hode - E°ano­de


Cell: 2 connected electr­odes in contact with an elec­tro­lyte
Battery: 2 or more cells connected in series; voltage of battery is the sum of the voltage of all the cells
Different kinds of batteries are made for diff­erent sized devices; the bigger the battery, the more electr­olytic soluti­on, and the longer it lasts
Alkaline battery: a battery that uses an alka­line (bas­ic) electrode rather than an acid
Primary vs secondary cells
Primary cells: non-­rec­har­gea­ble cells that run until reactants are used up (galvanic cells)
Secondary cells: cells that can be rech­arged by adding an electric current (galvanic when being used, electr­olytic when being recharged)


The brea­kdo­wn/­det­eri­oration of metal by a redox reaction
Cond­itions required for corros­ion
Cond­itions that accelerate corros­ion
Oxygen (O2)
High temper­ature
Water (H2(l))
Salt and/or other electr­olytes
Decrease in pH (more acidity)
Rusting: the corrosion of iron (Fe) spec­ifi­cally

Corrosion Prevention

Method 1: Galvanize the metal
Galvanizing: process where a metal (usually steel) is coated with a thick layer of zinc (Zn) to prevent corrosion
Method 2: Cathodic protec­tion
Cathodic protec­tion: form of corrosion prevention in which e¯ are contin­uously supplied to the metal being protected, making it a cath­ode
Two forms:
Sacr­ificial anode                                          
the oxidation of a more active metal attached to the metal being protected prevents the protected metal from corrosion
Impr­essed current
e¯ from a direct current (DC) power source are contin­uously supplied to the protected metal

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