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CHEMISTRY

CHEMISTRY. OXIDATION AND REDUCTION. OXIDATION & REDUCTION. Oxidation and reduction can be defined in two ways:. OXIDATION-REDUCTION REACTIONS. LEO SAYS GER. REDOX REACTION. What is a redox reaction?

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CHEMISTRY

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  1. CHEMISTRY OXIDATION AND REDUCTION

  2. OXIDATION & REDUCTION Oxidation and reduction can be defined in two ways:

  3. OXIDATION-REDUCTION REACTIONS LEO SAYS GER

  4. REDOX REACTION What is a redox reaction? All redox reactions exist when an element or a compound experiences a change in their oxidation state going from the left hand side of the equation to the right hand side. For example: Al + HCl AlCl3 + H2 0 +1 -1 +3 -1 0 The Al goes from 0 to +3, therefore it has been oxidized, while the H in HCl goes from +1 to 0 therefore HCl has been reduced.

  5. REDOX REACTION Oxidation and reduction ALWAYS occur at the same time in a chemical reaction. If one substance is oxidized, some other substance must be reduced. REMEMBER the substance oxidized is the REDUCING AGENT the substance reduced is the OXIDISING AGENT. Sodium is oxidized – it is the reducing agent Chlorine is reduced – it is the oxidizing agent

  6. REDOX REACTION In rare cases when an element or compound is simultaneously oxidized and reduced in the same chemical reaction, this is called disproportionation.  For example:2CuCl  Cu + CuCl2 +1-1 0 +2 -1 The Cu goes from +1 to 0 and +2, therefore it has been reducedANDoxidized, while the Cl remains with the same oxidation number.

  7. REDOX REACTION Each sodium atom loses one electron: Each chlorine atom gains one electron:

  8. REDOX REACTION Lose Electrons = Oxidation Gain Electrons = Reduction Sodium is oxidized Chlorine is reduced

  9. REDOX REACTION NOT ALL REACTIONS ARE REDOX REACTIONS Reactions in which there has been no change in oxidation number are not redoxrxns.

  10. REDOX REACTIONS Spontaneous redox reactions can transfer energy. • Electrons (electricity) • Heat Non-spontaneous redox reactions can be made to happen with electricity.

  11. RULES FOR ASSIGNING OXIDATION NUMBERS Rules 1 & 2 • The oxidation number of any uncombined element is zero 2. The oxidation number of a monatomic ion equals its charge

  12. RULES FOR ASSIGNING OXIDATION NUMBERS Rules 3 & 4 3.The oxidation number of oxygen in compounds is -2 4. The oxidation number of hydrogen in compounds is +1

  13. RULES FOR ASSIGNING OXIDATION NUMBERS Rule 5 5. The sum of the oxidation numbers in the formula of a compound is 0 2(+1) + (-2) = 0 H O (+2) + 2(-2) + 2(+1) = 0 Ca O H

  14. RULES FOR ASSIGNING OXIDATION NUMBERS Rule 6 6. The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its charge X + 4(-2) = -2 S O X + 3(-2) = -1 N O  X = +6  X = +5

  15. WRITING BALANCED EQUATIONSUSING TWO HALF EQUATIONS • In ½ equations, electrons are present. • When combining two ½ equations, one must ensure the number of electrons involved are the same for both equations. • This is done by multiplying one or both equations by an integer to make the number of electrons equal in both equations. • When this is done, all the species in the equation would be multiplied by that integer chosen. • Combine the two half equations withoutadding the electrons and you are done. • Remember the full equation must have the same total charge on both sides.

  16. EXAMPLE EQUATION A: Cr2O72-+ 14H+ + 6e- 2Cr3+ + 7H2O EQUATION B: Fe2+ - e- Fe3+ Equation A has 6 electrons involved and equation B has 1 electron involved. Therefore equation B is multiplied by 6 to have 6 electrons involved. Therefore equation B changes to: 6Fe2+ - 6e- 6Fe3+. Combining both equations and removing the electrons would give: 6Fe2+ + Cr2O72- + 14H+ 6Fe3+ + 2Cr3+ + 7H2O

  17. TRENDS IN OXIDATION AND REDUCTION • Active metals: • Lose electrons easily • Are easily oxidized • Are strong reducing agents • Active nonmetals: • Gain electrons easily • Are easily reduced • Are strong oxidizing agents

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