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Redox = "reduction and oxidation"
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Oxidation: lose electrons (e¯)
(Classical: any reaction involving oxygen gas (O 2) - rusting, combustion, etc.) |
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Reduction: gain electrons (e¯)
(Classical: reducing a metal ore into pure metal) |
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Both reactions always occur together
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Reducing agent: substance that causes another substance to become reduced
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Oxidizing agent: substance that causes another substance to become oxidized
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Reducing agents NEVER reduce themselves; is always oxidized to promote reduction
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Oxidizing agents NEVER oxidize themselves; is always reduced to promote oxidation
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Reducing/oxidizing agents are ALWAYS reactants
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Net charge that an atom would have if the e¯ pairs in covalent bonds belonged entirely to the more electronegative ion
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All redox reactions require a change in O.N.
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O.N. = oxidation (loss of e¯)
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O.N. = reduction (gain e¯)
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Rules to determine oxidation numbers (O.N.)
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Pure elements have O.N. = 0
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C
(s) = 0, O2(s) = 0, P4(s) = 0 |
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Monoatomic ions have O.N. = their charge
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Al3+ = +3, Cl¯ = -1
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Hydrogen always has O.N. = +1 (except metal hydrides = -1)
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HCl (H = +1); H
2S (H = +1); CaH2 (H = -1) |
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Oxygen always has O.N. = -2 (except peroxides = -1)
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Li
2O (O = -2); KNO3 (O = -2); H2O2 (O = -1) |
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In a compound, groups I, II, and IV usually have O.N. = ionic charge
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NaCl (Na = +1, Cl = -1); MgO (Mg = +2)
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In a neutral compound1, ΣO.N. = 0
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CF
4 (C = +4, F = -1) |
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In a polyatomic ion, ΣO.N. = ion's charge
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NO
3¯ (N = 5, O = -2; Σ O.N. = -1) |
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In molecular compunds with no O or H, the more electronegative atom has O.N. = its usual charge
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CS
2 (S = -2), Li3N (N = -3) |
3 - K = +1 because IO3 is 1-)|
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
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Cu
2O(aq) + H2SO4(aq) Cu(s) + CuSO4 (aq) + H2O(l) |
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Cu +1 to 0 gain 1 e¯ reduced
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Cu +1 to +2 lose 1 e¯ oxidized
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Since both of these happen in the same reaction, it is disproportionate
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Half-reaction: reactions made from overall net ionic equations that focus on 1 specific atom
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Zn
(s) + CuSO4(aq) ZnSO4(aq) + Cu(s) |
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Zn
(s) ZnSO4(aq) |
CuSO
4(aq) Cu(s) |
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Two types of electrochemical cells:
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Galvanic cells
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arrangement of 2 connected half-cells that spontaneously produce an electric current; e¯ always flow from high potential low potential
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Electrolytic cells
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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
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Parts of an electrochemical cell
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Electrolyte: solution that contains aqueous ions (cations (+) and anions (-))
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Electrode: solid metal conductor where redox reactions occur (cathode (oxidation) and anode (reduction))
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Salt bridge: tube that contains an electrolyte solution and connects the 2 half-cells; used to maintain electrical neutrality
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The measure of the electric potential difference (voltage) between 2 half-cells
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Standard cell: galvanic cell in which all entities are at SATP and all solution concentrations are 1.0 mol/L
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Standard cell potential (ΔE°): the ability of each half cell to gain e¯ (reduction)
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Cell potential formula:
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ΔE°
(cell) = E°cathode - E°anode |
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Cell: 2 connected electrodes in contact with an electrolyte
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Battery: 2 or more cells connected in series; voltage of battery is the sum of the voltage of all the cells
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Different kinds of batteries are made for different sized devices; the bigger the battery, the more electrolytic solution, and the longer it lasts
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Alkaline battery: a battery that uses an alkaline (basic) electrode rather than an acid
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Primary vs secondary cells
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Primary cells: non-rechargeable cells that run until reactants are used up (galvanic cells)
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Secondary cells: cells that can be recharged by adding an electric current (galvanic when being used, electrolytic when being recharged)
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The breakdown/deterioration of metal by a redox reaction
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Conditions required for corrosion
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Conditions that accelerate corrosion
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Oxygen (O
2) |
High temperature
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Water (H
2O(l)) |
Salt and/or other electrolytes
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Decrease in pH (more acidity)
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Rusting: the corrosion of iron (Fe) specifically
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Method 1: Galvanize the metal
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Galvanizing: process where a metal (usually steel) is coated with a thick layer of zinc (Zn) to prevent corrosion
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Method 2: Cathodic protection
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Cathodic protection: form of corrosion prevention in which e¯ are continuously supplied to the metal being protected, making it a cathode
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Two forms:
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Sacrificial anode
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the oxidation of a more active metal attached to the metal being protected prevents the protected metal from corrosion
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Impressed current
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e¯ from a direct current (DC) power source are continuously supplied to the protected metal
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