2Oxidation-Reduction (Redox) Reactions “redox” reactions: rxns in which electrons are transferred from one species to anotheroxidation & reduction always occur simultaneouslywe use OXIDATION NUMBERS to keep track of electron transfers
3Rules for Assigning Oxidation Numbers: 1) the ox. state of any free (uncombined) element is zero.Ex: Na, S, O2, H2, Cl2, O3
4Rules for Assigning Oxidation Numbers: 2) The ox. state of an element in a simple ion is the charge of the ion.Mg2+ oxidation of Mg is +2
5Rules for Assigning Oxidation Numbers: 3) the ox. # for hydrogen is +1(unless combined with a metal, then it has an ox. # of –1)Ex: NaOH (H bonded to O) v. NaH (H bonded to Na)
6Rules for Assigning Oxidation Numbers: 4) the ox. # of fluorine is always –1.
7Rules for Assigning Oxidation Numbers: 5) the ox. # of oxygen is usually –2.Why USUALLY? Not -2 when it’s in a peroxide, such as hydrogen peroxide:H2O2
8Rules for Assigning Oxidation Numbers: 6) in any neutral compound, the sum of the oxidation #’s = zero.
9Rules for Assigning Oxidation Numbers: 7) in a polyatomic ion, the sum of the oxidation #’s = the overall charge of the ion.
10Rules for Assigning Oxidation Numbers: **use these rules to assign oxidation #’s; assign known #’s first, then fill in the #’s for the remaining elements:
11Examples: Assign oxidation #’s to each element: a) NaNO3
12Examples: Assign oxidation #’s to each element: b) SO32-
13Examples: Assign oxidation #’s to each element: c) HCO3-
14Examples: Assign oxidation #’s to each element: d) H3PO4
15Examples: Assign oxidation #’s to each element: e) Cr2O72-
16Examples: Assign oxidation #’s to each element: f) K2Sn(OH)6
17DefinitionsOxidation: the process of losing electrons (ox # increases)Reduction: the process of gaining electrons (ox # decreases)Oxidizing agents: species that cause oxidation (they are reduced)Reducing agents: species that cause reduction (they are oxidized)
18To help you remember…OIL RIGOxidation Is LossReduction Is Gain
19Are all rxns REDOX rxns?a reaction is “redox” if a change in oxidation # happens; if no change in oxidation # occurs, the reaction is nonredox.
25Balancing Redox Equations In balancing redox equations, the # of electrons lost in oxidation (the increase in ox. #) must equal the # of electrons gained in reduction (the decrease in ox. #)There are 2 methods for balancing redox equations:Change in Oxidation-Number MethodThe Half-Reaction Method
261. Change in Oxidation-Number Method: based on equal total increases and decreases in oxidation #’sSteps:Write equation and assign oxidation #’s.Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each.Connect the atoms that change ox. #’s using a bracket; write the change in ox. # at the midpoint of each bracket.Choose coefficients that make the total increase in ox. # = the total decrease in ox. #.Balance the remaining elements by inspection.
28If needed, reactions that take place in acidic or basic solutions can be balanced as follows: add H2O to the side needing oxygenbalance as if in acidic sol’nthen add H+ to balance the hydrogenadd enough OH- to both sides to cancel out each H+ (making H2O) & then cancel out water as appropriate
29Example: Balance the following equation, assuming it takes place in acidic solution. ClO4- + I- Cl- + I2
302. The Half-Reaction Method: separate and balance the oxidation and reduction half-reactions.Steps:Write equation and assign oxidation #’s.Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each.Construct unbalanced oxidation and reduction half reactions.Balance the elements and the charges (by adding electrons as reactants or products) in each half-reaction.Balance the electron transfer by multiplying the balanced half-reaction by appropriate integers.Add the resulting half-reaction and eliminate any common terms to obtain the balanced equation.
31Example: Balance the following using the half-reaction method: HNO3 + H2S NO + S + H2O