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Powerpoint part one


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Powerpoint part one

  1. 1. Matt Berlin<br />Balancing Half Reactions and Predicting Single Replacement Reactions<br />
  2. 2. <ul><li>A half-reaction is one which shows either reduction OR oxidation, but not both
  3. 3. A half reaction does not occur by itself
  4. 4. At least two such reactions must be coupled for a Redox Reaction
  5. 5. Here are some examples of a half reaction</li></ul>2H+ + 2e-  H2 Cu  Cu2+ + 2e-<br /> Cl2 + 2e-  2Cl-<br />What is a Half Reaction<br />
  6. 6. <ul><li>When an element loses an electron it is oxidation
  7. 7. When an element gains an electron it is reduction
  8. 8. Remember the acronym O.i.l.r.i.g.</li></ul>Oxidation is loss<br />Reduction is Gain<br />What is Oxidized and What is Reduced?<br />
  9. 9. Steps to Balance A Half-Reaction<br />Identify key element that undergoes an oxidation state change<br />Balance the number of atoms for the key element<br />Add electrons to compensate for oxidation state change<br />Add H+ or OH- to balance charge<br />Add H2O as needed to balance equation<br />
  10. 10. Example Of Balancing a Half-Reaction<br />Balance the two half reactions for the reaction in an acid solution<br />H2O2 + I  I2 + H2O<br />First Identify what is being oxidized/reduced<br />The Half-reaction 2I-  I2 Shows Oxidation<br />The Half-reaction H2O2  2H2O<br />
  11. 11. Example Cont.<br />Add Electrons to compensate for the changes in oxidation state<br />For the Iodine reaction you must add two electrons to the products<br />2I- I2 + 2e-<br />For the Oxygen reaction you must add two electrons to the reactants<br />H2O2 + 2e- 2H2O<br />
  12. 12. Example Cont. <br />Add H+ or OH- to balance the charge<br />In this case H+ should be added to the reduction half reaction, the balanced equations are as follows<br />2I- I2 + 2e- (oxidation)<br />H2O2 + H+ + 2e-  2H2O (reduction)<br />
  13. 13. How to Become Better!!!<br /><ul><li> Practice, Practice, Practice
  14. 14. Remember O.i.l.r.i.g.</li></ul>Oxidation is loss<br />Reduction is gain<br />
  15. 15. Predicting Single Replacement Reactions<br />Activity Series<br /><ul><li>In a single replacement reaction a metal replaces a metal
  16. 16. The more active metal replaces the less active metal
  17. 17. The activity series determines which metal is more active</li></li></ul><li>Activity Series<br /><ul><li>The most active metal is at the top and as you move down the metals become less active
  18. 18. Same for nonmetals</li></li></ul><li>Rules for Single Replacement<br /><ul><li>If a metal is higher on the activity series then it will replace the lower metal</li></ul>Same Goes for nonmetals<br />Metals Replace metals <br />Nonmetals replace nonmetals<br />If a metal is lower on the activity series no reaction will occur<br />Same goes for nonmetals<br />The Basic formula for a single replacement is<br />A + BC  AC + B<br />
  19. 19. Steps for Doing a Single Replacement Reaction<br />Figure out if the reaction will occur by using the activity series<br />Write reactions<br />Write products<br />Balance<br />
  20. 20. Example of a Single Replacement Reaction<br />Magnesium Metal + Aqueous Aluminum Chloride<br />Figure out if the reaction will occur by looking at the activity series<br />Magnesium is higher on the activity series than aluminum<br />Thus the reaction will<br /> occur<br />Amy, “Beakers” January 28, 2010 via Flickr, Creative Commons Attribution. <br />
  21. 21. Example Cont.<br />2. Write the Reactants<br /><ul><li>We know Al has a 3+ charger and Cl has a 1- charge
  22. 22. Aluminum Chloride would be come AlCl3</li></ul>The reactants would look as follows:<br />Mg + AlCl3<br />
  23. 23. Example Cont.<br />Carry out the single replacement and write the products<br />Magnesium would replace aluminum in the reaction as follows<br />Mg + AlCl3 MgCl2 + Al<br />Al has a 1- charge and Mg has a 2+ charge so they combine to make MgCl2<br />
  24. 24. Example Cont.<br />Balance the Equation<br />The Cl’s in the equation are not balance<br />The least common multiple is 6 for the Cl’s<br />Multiply the AlCl3 by 2 and MgCl2 by 3<br />Now the Cl’s are balanced and the equations looks like this:<br />Mg + 2 AlCl3  3 MgCl2 + Al<br />
  25. 25. Example Cont.<br />Balance the equation<br />The Al’s and the Mg’s are now unbalanced<br />Multiply the Mg by 2 and the Al by 3 as follows<br />3Mg + 2AlCl3  3MgCl2 + 2Al<br />This is the final equation<br />
  26. 26. What was learned?<br /><ul><li>In this powerpoint you learned the following</li></ul>What is being reduced in a half-reaction<br />What is being oxidized in a half-reaction<br />How to balance half reactions<br />How to Predict single replacement reactions<br />Remember O.i.l.r.ig.<br />Practice makes Perfect<br />