Redox ReactionsRedox = "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 NumbersNet 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 ReactionsMost 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 CellsTwo 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 PotentialThe 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 |
| | BatteriesCell: 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) |
CorrosionThe breakdown/deterioration of metal by a redox reaction | Conditions required for corrosion | Conditions that accelerate corrosion | Oxygen (O2) | High temperature | Water (H2O(l)) | Salt and/or other electrolytes | | | Decrease in pH (more acidity) | Rusting: the corrosion of iron (Fe) specifically |
Corrosion PreventionMethod 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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