<p> Redox Reactions: Transferring Electrons</p><p>When we think about chemical reactions, in reality they are all about moving around electrons. </p><p>Let’s think about something that most people don’t think about as “chemistry”</p><p>A battery is basically a chemical reaction that’s producing electron flow for whatever gizmo you have it hooked up to</p><p>When a reaction transfers an electron from one substance to another we call it an “Oxidation-Reduction” reaction or Redox for short.</p><p>If a substance is losing the electron it is being oxidized If a substance is gaining, it is being reduced.</p><p>LEO says GER</p><p>How are Redox reactions different from other reactions we have studied so far?</p><p>The acceptance of an electron (reduction) by a material is matched with the donation of an electron (oxidation).</p><p>Each part is called a half-reaction, but they occur at the same time. </p><p>Because electrons are being transferred, the number of electrons gained must equal the number of electrons lost.</p><p>Important Terminology</p><p>If a substance is losing the electron it is being oxidized If a substance is gaining, it is being reduced.</p><p>LEO says GER</p><p>+ - Ag (aq) + e  Ag(s) (reduction) 2+ - Cu (s)  Cu + 2e (oxidation)</p><p>Because electrons are being transferred, the number of electrons gained must equal the number of electrons lost.</p><p>Reducing Agent: the one that does the reducing. Oxidizing Agent: the one that does the oxidizing.</p><p>Be Careful Here!</p><p>Oxidation States: Electron Bookkeeping</p><p>How do we know when a redox reaction is occurring?</p><p>Oxidation Numbers: positive or negative numbers assigned to atoms depending on their electronegativity.</p><p>Note: oxidation numbers are different from charges! We put the sign before the number for oxidation states.</p><p>Rules for Assigning Oxidation Numbers</p><p>1. Free elements (ie. Diatomic, solids) are zero. 2. The sum of oxidation numbers in an uncharged molecule must be zero or equal to the molecule’s charge. 3. Hydrogen always +1 (except -1 in hydrides) 4. Oxygen always -2 (except -1 in peroxides) 5. Group 1 = +1 6. Halogens = -1 7. Group 6 = -2 8. group 2 = +2</p><p>Remember the total oxidation numbers must equal the total charge.</p><p>Using Oxidation Numbers</p><p>By assigning oxidation numbers we can identify oxidation and reduction reactions. Basically we assign oxidation numbers to all atoms or groups in an equation. If we notice one atom or group has increased in oxidation state (more positive) it has been oxidized. If it has a lower oxidation state (more negative) it has been reduced. </p><p>This makes sense logically because gaining an electron makes something more negative and losing one makes something more positive.</p><p>Example:</p><p>2 KCl + MnO2 + 2H2O4  K2SO4 +MnSO4 + Cl2 + 2H2O</p><p>Let’s assign oxidation numbers:</p><p>+1 -1 +4 –2 +1 +6 –2 +1 +6 –2 +2 +6 –2 0 K Cl + Mn O2 H2 S O4  K2 S O4 Mn S O4 Cl2 –2 H2 +1 O </p><p>Notice that Cl goes from –1 to 0. It has been oxidized.</p><p>Mn goes from +4 to +2. It has been reduced.</p>