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Oxidation-Reduction (aka Redox) Reactions

Oxidation-Reduction (aka Redox) Reactions. These are electron- transfer reactions!. http://www.calgaryacademy.com/ICT/rr/redox1.html. Oxidation Numbers = ___________________________________________________ ___________________________________________________ For example:

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Oxidation-Reduction (aka Redox) Reactions

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  1. Oxidation-Reduction (aka Redox)Reactions These are electron- transfer reactions! http://www.calgaryacademy.com/ICT/rr/redox1.html

  2. Oxidation Numbers = ___________________________________________________ ___________________________________________________ For example: Zn0 + Cu+2 Zn+2 + Cu0 These are the oxidation numbers of each atom. They keep track of where electrons are.

  3. Oxidation= the _____ of electrons Magnesium’s oxidation # is 0 when it is a solid. After oxidation, magnesium’s oxidation # increases … Mg0(s)  Mg+2 + 2e- because _______ _____________ Notice: In an oxidation reaction, the element’s oxidation number will __________________..

  4. Reduction= the ________ of electrons Fe+3 + 3e-  Fe0(s) After reduction, iron has an oxidation # of 0… An iron ion has an oxidation # of +3 because _______ _____________ Notice: In a reduction reaction, the element’s oxidation number will ______________.

  5. Hint for remembering:

  6. Another hint for remembering: LEO the lion says GER LEO Loss of electrons is oxidation GER Gain of electrons is reduction

  7. Oxidation or Reduction??? 1.) Br20 + 2e-  2Br-1 ___________ 2.) Li0 Li+1 + e- ___________

  8. Redox Reactions:Oxidation & reduction reactions can’t happen alone! Oxidation: Mg0(s)  Mg+2 + 2e- Reduction: Fe+3 + 3e-  Fe0(s) Redox Reaction: 3 Mg0(s) + 2 Fe+3  2Fe0(s) + 3 Mg+2 Mg0 is the ________ agent The oxidation # increased. Mg(s) was oxidized. Mg lost electrons. Fe3+ is the ________ agent The oxidation # decreased. Fe3+ was reduced. Fe3+ gained those electrons. http://www.calgaryacademy.com/ICT/rr/redox4.html

  9. How can we identify a redox reaction? Step 1: Assign oxidation numbers to each atom in the reaction. Rules for Assigning Oxidation Numbers 1.) Any atom that is uncombined and has a neutral charge has an oxidation number of 0. Ex: 2Na + Cl2 NaCl

  10. 2.) Ions have an oxidation number equal to their ionic charge. Ex: Chlorine ion: Cl- 3.) Group I elements always have an oxidation number of +1 in compounds, while the Group II elements always have an oxidation number of +2 in compounds. Ex: K2SO4  Each K will have an oxidation # of ____ Ex: BaCl2  Ba will have an oxidation # of ____

  11. 4.) Fluorine is always -1 in compounds. The other halogens are also -1 when they are the most electronegative element in a compound. 5.) Hydrogen is +1 in compounds… EXCEPT if it is combined with a metal. Ex: HCl  hydrogen will be ___ LiH  hydrogen will be ___ 6.) Oxygen is usually -2 in compounds… EXCEPT when it is combined with fluorine, it becomes +2….and EXCEPT when it is in the peroxide ion (O22-), it becomes -1.

  12. Rule 6: -2 7.) The sum of the oxidation numbers in all compounds must be 0. Ex: HNO3 1H  ____ 1N  ____ 3O  ____ Rule: +1

  13. Rule 6: -2 8.) The sum of oxidation numbers in polyatomic ions must be equal to the charge on the ion. Ex: Cr2O72- 2Cr  ____ 7O  ________

  14. How can we identify a redox reaction? Step 1: Assign oxidation numbers to each atom. (Use rules!) MnO2 + 4HCl  MnCl2 + Cl2 + 2H2O

  15. How can we identify a redox reaction? Step 2: Identify whether there are any changes in oxidation number for a particular atom between the reactant and product sides. **If there is a change in oxidation number for particular type of atom, the reaction is redox.*** MnO2 + 4HCl  MnCl2 + Cl2 + 2H2O +4 +1 +2 -2 -1 -1 0 +1 -2

  16. Another Example PbO2 + 4HI  I2 + PbI2 + 2H2O

  17. Half-Reactions= show either oxidation or reduction portion of the redox reaction PbO2 + 4HI  I2 + PbI2 + 2H2O Oxidation: Reduction: 0 -1 +2 -1 +1 -2 +4 -2 +1

  18. ***There must be the same number of atoms on each side of the reaction, and the net charge must be the same on both sides.*** Oxidation: I1- I20 Reduction: Pb4+ Pb2+

  19. Another Example: Cu + AgNO3 Cu(NO3)2 + Ag Oxidation: Reduction:

  20. Electrochemical Cells = _______________________________________ - 2 types: • Voltaic cell = _________________________________________________ (2) Electrolytic cell = ______________________________________________ ______________________________________________ - Electrode = ____________________________________________ ____________________________________________  anode = where _____________ occurs  cathode = where _____________ occurs

  21. RED CAT: Reduction occurs at the cathode AN OX: Anode is the site of oxidation

  22. Voltaic Cells (aka galvanic cells) **_______________________________________________________.** (2) The electrons leave the zinc anode and pass through the external circuit to the copper rod. (1) Electrons are produced at the zinc rod according to the oxidation half- reaction: B/c zinc is oxidized at the zinc rod, the zinc rod is the anode. The anode (in a voltaic cell) is a negative electrode. (3) Electrons enter the copper rod and interact with copper ions (Cu+2) in solution. There, a reduction half-reaction occurs: B/c the copper ions are reduced at the zinc rod, the zinc rod is the cathode. The cathode (in a voltaic cell) is a positive electrode. (4) To complete the circuit, both positive and negative ions move through the aqueous solutions via the salt bridge.

  23. Animation of Voltaic Cell: http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf Virtual Lab: http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/voltaicCellEMF.html

  24. Electrolytic Cells **_____________________________________________________________________________________________________________________________.** Example: Electroplating (2a)The electrons produced by the reduction of Ag are moved through the external wire by a power source. (1) Silver (Ag) is oxidized to produce silver ions (Ag+1) in the following oxidation half-reaction: Since the silver rod is the site of oxidation, it is the anode. However, in electrolytic cells, the anode is the positive electrode. (3) Electrons flowing through the spoon cause the reduction of silver ions in the following reduction half-reaction: Since reduction is occurring at the spoon, the spoon is the cathode. When silver ions become reduced, they plate (cover) the spoon. (2b) Positive silver ions migrate away from anode, towards the cathode, which, in electrolytic cells is the negative electrode.

  25. Similarities Voltaic Cells Electrolytic Cells

  26. **To determine which substance is the anode/cathode (1) Check Table J: Activity Series (2) The metal that is higher on the chart will be oxidized and thus is the anode. (3) The metal that is lower on the chart will be reduced and this is the cathode. **In a voltaic cell:  Anode = negative electrode  Cathode = positive electrode

  27. FAT CAT Electrons flow From Anode To CATthode.

  28. Salt bridge: - connects the 2 containers & provides a path for a flow of ions between the two beakers. -As electrons leave one half of a galvanic cell and flow to the other, a difference in charge is established. If no salt bridge were used, this charge difference would prevent further flow of electrons. A salt bridge allows the flow of ions to maintain a balance in charge between the oxidation and reduction vessels while keeping the contents of each separate. With the charge difference balanced, electrons can flow once again, and the reduction and oxidation reactions can proceed.

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