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JUST REMEMBER . . .

Learn about REDOX reactions and oxidation states in chemistry. Discover the rules for assigning oxidation numbers and how to balance REDOX reactions. Explore the activity series and understand electrochemical cells and electroplating.

kennethrussell
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JUST REMEMBER . . .

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  1. JUST REMEMBER . . . “OIL RIG” (oxidation is losing, reduction is gaining)

  2. REDOX • In a REDOX reaction, when something is being oxidized, it is called the reducing agent, and when something is being reduced it is called the oxidizing agent

  3. Oxidation States • Oxidation States are assigned to atoms to identify how many electrons are either gained or lost by an atom • For example metals in Group 2 (like Ca) have a +2 oxidation state. Therefore Metals in Group 2 lose two electrons when they form compounds • Changes in oxidation numbers indicate that a redox reactionhas occurred. • It is important to learn the rules for assigning oxidation states to atoms in order to determine whether oxidation or reduction has occurred.

  4. Oxidation State Rules 1) Free elements (not combined with any other element) have an oxidation number of zero. Ex: Na, O2, H2 2) All metals in Group 1 have an oxidation number of +1. 3) All metals in Group 2 have an oxidation number of +2. 4) F (fluorine) always has an oxidation of -1 5) The oxidation of simple ions is equal to the charge on the ion. Ex: Mg+2 has an oxidation number of +2. 6) The sum of the oxidation numbers must equal 0 Examples:NaCl & MgCl2 7) In polyatomic ions, the sum of the oxidation numbers of all the atoms must equal the charge of the ion. Example: sulfate ion SO4-2 . 8) In general, oxygen has an oxidation number of -2. Oxygen has an oxidation number of -1 in peroxides (O2-2) Example: H2O2. Oxygen has an oxidation number of +2 in compounds with fluorine Example: OF2 9)Hydrogen has an oxidation number of +1 in all compounds combined with a non-metal. Hydrogen has an oxidation number of -1 when it is a metal hydrides (metal and hydrogen. Example:LiH, and CaH2

  5. Assigning Oxidation Numbers Dr. McGuiness’ “bookend” technique to assigning oxidation numbers to compounds with more than two elements. • Identify the oxidation # of the last element (overall charge) • Identify the oxidation # of the first element (overall charge) • If there is no charge to the compound, then the overall charge must be 0, therefore you can determine the oxidation # of the element in the middle. Li(MnO4) (+1) + (?) + (-8) = 0 -2 x 4 = -8 +1 x 1 = +1

  6. Half-Reactions Example: Cu + Ag(NO3) → Cu(NO3)2 + Ag • Assign oxidation numbers to everything • See which oxidation numbers change from reactant side to product side. • Determine which half-reaction is oxidation (loses e-), and which is reduction (gains e-) • OX: Cu0 → Cu2+ + 2e- • RED: Ag+ + 1e- → Ag0

  7. Balancing REDOX reactions Cu + 2Ag(NO3) → Cu(NO3)2 + 2Ag • OX: Cu0 → Cu2+ + 2e- • RED: 2(Ag+ + 1e- → Ag)  2Ag+ + 2e- 2Ag0 • When copper is oxidized, it loses 2 electrons, which are gained by the silver ion. • Copper = reducing agent • Silver = oxidizing agent

  8. Activity Series – Table J • Table J compares how active each metal and nonmetal is. • Metals higher up are more active, and replace metals from below them from compounds (remember single replacement???) • In a single replacement, the free element has to be more reactive than the element in compound in order for the reaction to be spontaneous. • If it isn’t – the reaction does NOT GO!!

  9. Electrochemical Cells • There are two types of electrochemical cells, Voltaic & Electrolytic • These cells rely on REDOX reaction in different ways to either generate energy or to separate elements in compound that would normally not exist on their own in nature.

  10. Voltaic Electrolytic • Spontaneous • Chemical  Electrical Energy • Anode – oxidation (-) • Cathode – reduction (+) • Example: BATTERY • Electrons travel from anode to cathode (wire) • More reactive metal ALWAYS the site of oxidation • Salt bridge is for the flow of ions from one half-cell to another • Non-spontaneous • Electrical  Chemical • Requires a battery as a source of energy • Anode – oxidation (+) • Cathode – reduction (-) • (+) charged ion moves toward the cathode • (-) charged ion moves toward the anode

  11. “RED CAT” “AN OX”

  12. Electroplating • Another example of an electrolytic cell. • The process of electroplating requires a layer of metal such as silver or copper, coating or covering any object to be plated (spoon or fork) • The item being plated is the cathode - reduction/(-) • The electrode must be the same metal that you are plating the object in. Cathode (-)

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