Unit 8 Redox Reactions Chapter 9

Oxidation-Reduction Reactions Reactions in which one substance is reduced and one is oxidized Electrons are exchanged between the oxidized substance and reduced substance (electron competition) Oxidation and reduction always occur together Can be single replacement, synthesis, or decomposition

Oxidation Loss of electrons (LEO) Gain of oxygen atom Loss of hydrogen atom Increase in oxidation state For example: Zn  Zn2+ + 2 e-

Reduction Gain of electrons (GER) Loss of an oxygen atom Gain of a hydrogen atom Decrease in oxidation state Cu2+ + 2 e-  Cu

Determining Oxidation State (number) Oxidation state: The charge an atom has, or appears to have, when the electrons of the compound are counted in accordance with a set of rules. Value is a measure of electron control/possession it has in a compound relative to the atom in the pure element (page 163) A change in oxidation state tells us that a redox reaction has occurred Two parts to oxidation number: Sign + means atom has lost electron control, - means it has gained electron control Value Refers to # of electrons over which control has changed Written with sign first followed by # (+2, -3)

Practice Find the oxidation state of Mn in KMnO4 O = -2, there are 4 of them… -2 * 4 = -8 Algebra: Mn +(+1) + (-8) = 0 Mn must be +7 in order for the molecule to have 0 overall charge Answer: Mn+7 Find the oxidation state of S in SO42- S + (-8) = -2 S = +6 Answer: S+6

Vocabulary Half Reactions: A rxn that involves only one element and shows the number of electrons gained or lost by that element Cu2+ + 2 e-  Cu Zn  Zn2+ + 2 e- Spectator Ions: Other ions in the reaction that are not oxidized or reduced. These are present in solution but not involved in rxn. SO42- Net Ionic Equations: A redox reaction that involves only the substances that are oxidized or reduced. The sum of the half reactions. Cu2+ + Zn  Cu + Zn2+ Notice that there are no electrons written on either side. Electrons must cancel out in order for the reaction to be balanced.

Balancing Redox Reactions Write half reactions first Identify species oxidized and reduced Write half reaction for each X  X2+ Y3+  Y Balance elements other than oxygen using coefficients Add same # of H2O as there is oxygen to the side w/o oxygen (If no oxygen, skip to electron step) Balance H ions by adding H+ to side that needs more Add electrons to one side of reaction so charge is conserved Do same with other half reaction Multiply half-reactions by coefficient so that when added, electrons cancel Add two half reactions to get overall equation

Oxidizing and Reducing Agents The substance that is oxidized is called the “reducing agent” Reactant that causes reduction of another compound Reactant that is oxidized Reactant that loses electrons Hydrogen is a reducing agent Typically metals (b/c they lose electrons) Which is the reducing agent? Zn + CuSO4  ZnSO4 + Cu Hint: Which substance lost electrons The substance that is reduced is called the “oxidizing agent” Reactant that causes oxidation of another compound. Reactant that is reduced. Reactant that gains electrons Oxygen is an oxidizing agent Typically nonmetals (b/c they gain electrons) Which substance is the oxidizing agent? Hint: Which substance gained electrons?

Reactivity More reactive metal will displace less reactive metal More reactive halogen will displace less reactive halogen Stronger reducing agent replaces weaker Stronger oxidizing agent replaces weaker For example: zinc replaces copper in copper (II) sulfate, so zinc is higher on the metal activity series So, which metal is oxidized? Which is reduced? (Think about the electrons, which accepted them?) You will need to be able to look at experimental data to determine activity series.