Redox Reactions
Redox = "reduction and oxidation" |
Oxidation: lose electrons (e¯) (Classical: any reaction involving oxygen gas (O 2
) - 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, O 2(s)
= 0, P 4(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); H 2
S (H = +1); CaH 2
(H = -1) |
Oxygen always has O.N. = -2 (except peroxides = -1) |
Li 2
O (O = -2); KNO 3
(O = -2); H 2
O 2
(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 |
|
In a polyatomic ion, ΣO.N. = ion's charge |
NO 3
¯ (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 |
CS 2
(S = -2), Li 3
N (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. KIO 3
- K = +1 because IO 3
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 |
Cu2
O(aq)
+ H2
SO4(aq)
Cu(s)
+ CuSO4 (aq)
+ H2
O(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)
|
|
|
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
|
|
|
Batteries
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) |
Corrosion
The breakdown/deterioration of metal by a redox reaction |
Conditions required for corrosion |
Conditions that accelerate corrosion |
|
High temperature |
|
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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