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Electrochemistry Electron Transfer Reactions Reactions where electrons flow between atoms Only reactions which produce electricity Only valence electrons are transferred Oxidation State Used to determine if electrons are transferred in a reaction Artificial numbers which DO NOT INDICATE the atom’s ionic CHARGE ELECTRON BOOKKEEPING Oxidation Numbers charge an atom would have if electrons in its bonds belonged completely to the more electronegative atom oxidation numbers are written as +1 vs. 1+ to distinguish them from ionic charges. different than formal charge (charge an atom would possess if all atoms had same electronegativity) Assigning Oxidation States Rules on p.137 Differ some depending on resource Oxidation States Shortcut Method RULE 1 The oxidation state of an atom in the most abundant, naturally occurring form of an element is typically zero. EX: diatomic noble gas metal others H2 Ne Fe P4, S8 Oxidation States Shortcut Method Rule 2 The oxidation state of any Group 1A atom in a compound is always +1. Rule 2 The oxidation state of any Group 2A atom in a compound is always +2. Oxidation States Shortcut Method Rule 2 The oxidation state of a monatomic ion (one atom) is equal to the charge of the ion. Ex: Al3+, N3-, S2- Oxidation States Shortcut Method Rule 3a The oxidation state of oxygen in a compound is almost always -2. Exceptions: peroxides O22- are -1 when oxygen is bound to a more electronegative atom (fluorine) Oxidation States Shortcut Method Rule 3b The oxidation state of hydrogen in a compound is almost always +1. Exceptions: when hydrogen is bound to a less electronegative atom (hydrides) Ex: AsH3, SiH4, GeH4 Oxidation States Shortcut Method Rule 3c If a halogen is in a compound, it will most likely have a -1 oxidation number. The oxidation number of fluorine is always -1. Exception: Br, Cl or I is bonded to oxygen or fluorine Oxidation States Shortcut Method Rule 4 The sum of the oxidation states for every atom in a formula must add up to the overall charge of the formula. Ex: H2O, (SO4)2-, (NH4)1+ Another Version of Rules 1. 2. 3. 4. 5. 6. 7. Any element by itself is 0 (O in O2 is zero) H is +1 except in metal hydrides (-1) O is –2 except in peroxides (-1) or combined with F (+2) Metal in ionic compound equals its positive ionic charge In covalent compound the most electronegative atom gets the negative charge. Sum of compound’s oxidation numbers must equal zero. Sum of polyatomic ion’s oxidation numbers must equal the charge of the ion. What is REDOX? REDOX stands for REDuction/OXidation Oxidation is often thought of as a combination of a substance with oxygen (rusting, burning) Oxidation refers to a loss of electrons Reduction refers to a gain of electrons OIL RIG Oxidation Is Loss of electrons Reduction Is Gain of electrons Increase in oxidation number Oxidizing & Reducing Agents REDUCING AGENT substance in a REDOX reaction that donates an electron Contains element that is oxidized organic chemistry: oxidations remove protons or add oxygen to the organic molecule Oxidizing & Reducing Agents OXIDIZING AGENT Substance that readily transfers oxygen atoms or that gains electrons in a redox reaction Contains element that is reduced organic chemistry: reductions adds protons or remove oxygen from an organic molecule. Applications of Redox Batteries Electroplating (mirrors, jewelry) Fuel cell Bleaching Hardening of arteries Cancer arthritis Using Organic Redox majority of the energy generated in our bodies comes from aerobic metabolism C6H12O6(aq) + 6O2(g) 6CO2(g) + 6H2O(l) oxidation of compounds such as carbohydrates (using the O2 that we breathe) in redox reactions produces ATP, our “energy currency” Pyruvate + NADH + H+ lactate + NAD+ Using Organic Redox BREATHALYZER Ethanol is oxidized to acetic acid Oxygen is reduced to water CH3CH2OH(l) + O2(g) → CH3COOH(l) + H2O(l) Electric current produced is measured and displayed as blood alcohol content Practice For each reaction determine the: 1. oxidation numbers of each atom 2. the oxidized element 3. the reduced element 4. Compound containing the reduced element (oxidizing agent) 5. Compound containing the oxidized element (reducing agent) 1. 2. 3. 4. 2Na + Cl2 2NaCl 2H2 + O2 H2O 4Fe + 3O2 → 2Fe2O3 C6H12O6 + 6O2 6CO2 + 6H2O Balancing Redox Reactions 1. Write separate equations (halfreactions) for oxidation and reduction 2. For each half-reaction: Balance elements involved in e- transfer Balance number e- lost and gained by multiplying half reactions by whole numbers 3. Add half-reactions/cancel like terms (e-) Acidic and Basic 4. Acidic conditions: Balance oxygen using H2O Balance hydrogen using H+ 5. Basic conditions: Balance oxygen using OH Balance hydrogen using H2O Examples Acidic conditions: MnO CN 4(aq) 2 (aq) Fe Mn acid 2 (aq) 3 (aq) Fe (aq) MnO 2(s) Basic conditions: (aq) MnO 4 (aq) CNO base Voltaic Cells converts chemical energy into electrical energy perform electrical work battery - two or more voltaic cells connected in series to produce a steady flow of current Voltaic Cells Cathode – reduction half reaction; positive Anode – oxidation half reaction; negative Electrons flow from anode cathode (ac) Voltaic Cells Half cell – one compartment of voltaic cell Salt bridge – electrolyte solution or membrane which allows ions to flow to balance overall charge of half cell Voltaic Cells Zn Zn2+ + 2eoxidation anode (-) Cu2+ + 2e- Cu reduction cathode (+) Cell EMF Electromotive force (EMF) Difference in potential that pushes ethrough the external circuit Volt (V) = 1 Joule (J) 1 Coulomb (C) Standard Cell Each half cell is at standard conditions: 1atm, 1M, 25°C Cell EMF Standard Reduction Potential (E0r) Standard half cell potential of a reduction reaction Measured in Volts relative to standard hydrogen electrode (0 V) Appendix E (p.1117) Cell EMF Standard Cell Potential (EMF) E0 cell = E0r cathode – E0r anode Intensive property – doesn’t change with amount (like density, boiling point) + value spontaneous Cell EMF Standard Cell Potential (EMF) Cathode is half reaction with more positive E0r More positive E0r means more likely to be reduced (oxidizing agent) Halogens, oxygen Very negative E0r (reducing agent) Cell EMF Al Ag1+ NO31- E0r = -1.66 V E0r = +0.80 V E0r = +0.96 V NO31- is strongest oxidizing agent Al is strongest reducing agent Cell EMF – calculate E0cell Example Cd2+ + 2e- Cd Sn2+ + 2e- Sn E0 cell E0r = -0.403 V E0r = -0.136 V (cathode) = E0r cathode – E0r anode = (-0.136) – (-0.403) = 0.267 V (spontaneous) Cell EMF - calculate E0r Example In1+ 2e- + In3+ Br2 + 2e- 2 Br1E0 cell = 1.46 V E0r = ? E0r = +1.06 V