1. Oxidation-Reduction Review

2. An oxidation-reduction reaction involves the transfer of electrons (e - ). Be able to recognize oxidation/reduction reactions under the following circumstances: In chemical equations; everything except double replacement and acid base are redox reactions. If you’re not sure, look for changes in oxidation numbers of reactants and products.

3. In half reactions, oxidation always has electrons on the product side, reduction has them on the reactant side.

4. Net Ionic Equations In net ionic equations (which are always single replacement reactions on the regents), the ion on the reactant side is being reduced and the metal is being oxidized. The metal is always more active than the ion (above the ion on Table J) Zn + Cu 2+  Zn 2+ + Cu

5. Reduction is the gain of electrons. Reduction results in a reduction in the oxidation number (becomes less positive, more negative). Cu 2+ + 2e -  Cu oxidation # +2 to 0; reduction

6. A half-reaction can be written to represent reduction. Generically, M n+ + ne -  M Unbalanced half-reactions (oxidation and reduction) are often wrong answers on regents questions.

7. Oxidation is a loss of electrons. Historically, oxidation was explained as the chemical combination with oxygen. Oxidation raises the oxidation number of the element (more positive); Reduction lowers the oxidation number (more negative)

8. A half-reaction can be written to represent oxidation. M  M n+ + ne - for oxidation; M n+ + ne -  M for reduction Zn + Cu 2+  Zn 2+ + Cu Zn  Zn 2+ + 2e - oxidation Cu 2+ + 2e -  Cureduction

9. In a redox reaction, there is a conservation of charge and mass. The number of electrons lost is equal to the number of electrons gained. There is usually one question on the law of conservation of mass (and charge, and energy) on every exam.

10. Oxidation numbers can be assigned to atoms and ions. Changes in oxidation numbers indicate oxidation and reduction. Elements always have an oxidation number of zero. In spontaneous reactions, the more active metal (higher on Table J) will be oxidized, and get a (+) oxidation number (usually) based on the number of valence electrons.

11. Corrosion of metals, combustion of fuels, and spoilage of foods are examples of redox reactions. The most common corrosion reaction involves adding acid (H + ) to active metals, producing hydrogen gas. Often they will ask you for the metal that doesn’t react: it must be Cu, Ag, or Au, which are below H 2 on the activity series. Combustion reactions are usually used to balance equations. Try balancing the combustion of methane, ethane, and propane: Hydrocarbon + O 2  CO 2 + H 2 O

12. In an electrochemical cell, oxidation occurs at the anode and reduction at the cathode. The more active metal is the anode (higher on Table J); the less active metal (lower on Table J) is the cathode. The ion of that metal will be reduced. When the electrolyte in the reduction half-cell is used up, the battery “dies.” Electrons flow from the anode to the cathode. They ask questions about this nearly every year. Be on the alert for the inevitable question about salt bridges. If they ask you what they are for, write these two words: “mobile ions.”

13. A voltaic cell can operate without an outside energy source. At the heart of every battery is a spontaneous redox reaction. The electrons flow from the more active metal, whose valence electrons are in higher energy states than they will be when they reach the ion of the less active metal.

14. An electrolytic cell requires an outside energy source. The power supply makes a non-spontaneous process happen. The power supply “pulls” electrons from elements with low oxidation numbers and puts them on the element with the higher oxidation number. For example, for the non-spontaneous process 2H 2 O  2H 2 + O 2 H goes from +1 to 0; O goes from –2 to O Electrolytic cells often use two identical electrodes; there is also no separation between the electrodes (no half-cells)

15. Refining of metals and electroplating are practical applications of electrolytic cells. Group 1 and 2 metals are produced this way. Electrolysis of “fused” salts also produce pure halogens like Cl 2 and F 2. The anode is always where oxidation occurs, and the cathode is always where reduction occurs. Remember: an-oxred-cat