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Oxidation/Reduction - originally, oxidation described rxns in which oxygen was added to a reactant - reduction first meant the removal of oxygen from a compound - today’s definition: - oxidation is the process by which a substance loses one or more electrons - reduction is the process by which a substance gains one or more electrons - mnemonic devices: - OIL RIG: Oxidation Is Loss of electrons; Reduction Is Gain of electrons - LEO the lion goes GER: Loss of Electrons is Oxidation and Gain of Electrons is Reduction - Oxidation and Reduction ALWAYS occur together - the two rxns together are called oxidation-reduction or redox rxns Oxidation States - system of keeping track of e - in redox rxns by assigning charges to various atoms in a compound - also called oxidation numbers - oxidation states are NOT actual charges Rules for Assigning Oxidation States 1 – the oxidation number of an atom of an uncombined element is zero 2 – the oxidation number of any monatomic ion equals its ionic charge 3 – in compounds, the oxidation number of many elements corresponds to the elements position on the periodic table a. Alkali Metals are ALWAYS +1 b. Alkaline Earth Metals are ALWAYS +2 c. Aluminum is ALWAYS +3 d. Fluorine is ALWAYS -1 (b/c it is the most electronegative atom) e. Hydrogen has an oxidation number of +1 when combined with nonmetals f. Oxygen has an oxidation number of -2, unless it is in peroxide in which case oxygen is -1 or it is paired with Fluorine 4 – the oxidation numbers of elements in compounds are written per atom 5 – the algebraic sum of the individual oxidation numbers of all the atoms in the formula for a compound is zero 6 – the algebraic sum of the individual oxidation numbers of all the atoms in the formula for a polyatomic ion is equal to the charge of the ion 7 – in binary compounds, the element with the greatest electronegativity is assigned a negative oxidation number equal to its charge as an anion in its ionic compounds Oxidizing & Reducing Agents - in redox rxns, e- are lost and gained at the same time - the # lost must equal the # gained (law of conservation of charge) - the substance that gives up e - is the reducing agent - the reducing agent causes reduction by being oxidized - the substance that gains e - is the oxidizing agent - the oxidizing agent causes oxidation by being reduced - many strong oxidizing agents contain an atom with a high electronegativity - also many strong oxidizing agents contain an atom with a high positive oxidation # of +5, +6, or +7 b/c they easily accept electrons - common reducing agents include active free metals b/c they have low electronegativities and tend to give up their valence e- easily - other reducing agents include substances that burn readily such as hydrogen and organic substances Balancing Redox Equations Half-Reaction Method - half-rxns are equations that have e - as reactants or products - two half rxns are always devised: one to describe reduction and one to describe oxidation - two rules are used to write and balance half-rxns: - always write the formulas of molecules and ions as they actually occur - in aqueous sol’n, H +, H2O, and OH- are available to aid in balancing - in acid sol’n, H+ and H2O are available - in basic sol’n, OH - and H2O are available - each half-rxn and redox rxn can be balanced three ways: by electrons, by total charge and by atoms - any two of these methods are sufficient to give an equation overall balance - the two half-rxns must ultimately be added - the two rxns must show that the # e - lost equals the # e gained, so some multiplication may be required To balance using half rxns – acidic solution: 1 – Separate the rxn into an oxidation half rxn and a reduction half rxn 2 – Assign oxidation # to all atoms in the equation, writing them above or below the chemical symbol for each element 3 – Balance oxidized element and the reduce element numerically 4 – Balance the oxygens by adding water 5 – Balance the hydrogens by adding H + 6 – Add the two rxns together, canceling as needed To balance using half rxns – basic solution: 1 – Approach just like balancing in acidic solution 2 – Add the same # of OH- to the side with the H + ions, as H+ions, this will make H 2O 3 – Cancel H2O so that in any given half rxn, the water only appears on one side of the equation 4 – Add the two rxns together, canceling as needed Oxidation Number Method - for more complex redox rxns: - Remember: # e - lost in oxidation = # e - gained in reduction 1 – Assign oxidation # to all atoms in the equation, writing them above or below the chemical symbol for each element 2 – Identify the element oxidized and the element reduced and determine the change in oxidation # of each 3 – Connect the atoms that change oxidation # by using a bracket and write the change in oxidation # at the midpoint of each bracket 4 – Choose coefficients that make the total increase in oxidation # equal to the total decrease in oxidation # 5 – Balance the remaining elements by inspection, and then check the final equation - a rxn occurring in acidic aqueous sol’n can be balanced by adding H2O and H+ to either side of the equation - add H2O to the side of the equation that needs oxygen and H+ to the side that needs hydrogen atoms Electrochemistry - study of the interchange of chemical and electrical energy - deals with electricity-related applications of redox rxns - Remember: Redox rxns involve a transfer of e - - can be used to produce an electric current - this transfer also involves a release of heat - separating the oxidizing agent from the reducing agent is key in harnessing the energy from redox rxns - if not separated e - exchange occurs directly with no useful work obtained Electrochemical Cells - devices that use redox rxns to either produce or use electricity - voltaic cells – electrochemical cells that produces electricity as a result of a spontaneous redox rxns - also called galvanic cells - the transfer of e - takes place through an external pathway rather than directly between reactants - electrolytic cells – electrochemical cells that uses electrical energy to drive nonspontaneous redox rxns Voltaic Cells - Luigi Galvani first credited with discovering electricity in the late 1700’s when he noticed twitching in the limbs of animals he was dissecting when he touched them with tools made of different metals - In 1800, Alessandro Volta built the first batter by alternating disks of dissimilar metals separated by pieces of leather that had been soaked in salt water - voltaic cells can be visualized as consisting of two half cells - - - - one corresponding to the oxidation half-rxn and one corresponding to the reduction half-rxn electrodes – two solid metals are connected by an external circuit anode – the electrode at which oxidation occurs cathode – the electrode at which reduction occurs batteries are marked with (+) and (–) signs to indicate positive and negative electrodes - (–) indicates the anode b/c e - are released from the electrode - (+) indicates the cathode b/c e - are attracted to the electrode e- move spontaneously from anode to cathode a salt-bridge is used to keep the # of positive and negative ions equal in each half cell and is essential to complete the electrical circuit - without flow of ions within the cell, charge will begin to accumulate in each half cell - flow of e- between electrodes will stop almost instantaneously if charge builds up the salt-bridge is a tube or porous clay barrier that allows ions to move from one half cell to another but prevents the sol’n from mixing totally Voltmeter Anode (-) 1.10V Cathode (+) Salt-Bridge Cell Potential - e- move from the electrode with higher electrical potential energy to the electrode with lower electrical potential energy - cell potential – represents the difference in the electrical potential energy between two electrodes of the cell - measured in volts - also called cell voltage - by convention, the potential associated with each electrode is chosen to be the potential for reduction to occur at that electrode - standard electrode potentials are tabulated for reductions rxns - called standard reduction potentials - standard reduction potentials of all electrodes are measured against one standard electrode, the standard hydrogen electrode (SHE) - - the potential of the SHE is defined as zero the more positive the standard reduction potential, the more readily that species is reduced the more negative the standard reduction potential, the harder the species is to reduce o the cell potential, Ecell is calculated by o o o E E E cathode anode - cell to identify the anode and cathode: - the more positive standard reduction potential reduction cathode - the more negative standard reduction potential oxidation anode Batteries - different kinds of batteries are based on different redox rxns - those that cannot be recharged use redox rxns that cannot be easily reversed (primary batteries) - batteries that can be recharged use an outside power source to reverse electrode rxns (secondary batteries) Dry Cell Batteries - a voltaic cell in which the electrolyte is a paste - in common dry cell batteries, the paste consists of ZnCl 2, MnO2, NH4Cl, and H2O - a zinc container, which serves as the anode, is separated from the other chemicals by a liner of porous paper that acts as a salt-bridge - a graphite rod is in the center of the paste to serve as an inexpensive cathode - the battery will continue to produce e - until the reduction product ammonia comes out of its aqueous sol’n as a gas - advantages: low cost - disadvantages: - not rechargeable - if current drawn rapidly from cell, the voltage drops from its normal 1.5V, as reactants close to the cathode are depleted - Zn reacts directly with the slightly acidic ammonium ions, causing the cell to run down even when current is not being drawn - leads to poor shelf-life - putting batteries in the fridge will extend life by decreasing the rate of redox Alkaline Dry Cell - modification of the ordinary dry cell - KOH replaces NH4Cl in the paste - powdered zinc is used in the paste that surrounds the anode instead of zinc container as the anode - the zinc paste is separated from the alkaline paste by porous paper or a fabric barrier - pastes are contained in a steel jacket and have a brass current collector - advantages: - longer shelf-life - maintains a steady voltage under high current loads - generates about 50% more total energy than dry cells of the same size - disadvantages: higher cost Nickel-Cadmium Batteries - rechargeable voltaic cells - also called NiCad batteries - anode is cadmium, cathode is NiO 2 paste - products formed in this redox rxn are insoluble and adhere to the electrode surfaces which allows for recharging - advantages: lightweight and produces a constant voltage during discharge - disadvantages: - if discharged partially and then recharged, it develops the tendency to need recharging after only a short use - must be disposed of properly b/c of the toxicity of cadmium and its compounds Lead Acid Storage Batteries - standard automotive battery - most contain 6 cells that generate about 2V each for a total output of 12V - called lead storage battery b/c its electrodes are composed of alternating sheet so lead and lead dioxide - the electrolyte is sulfuric acid - during discharge, the lead is the anode and the lead dioxide is the cathode - the lead and lead dioxide are separated from each other by porous spacers - the lead sulfate formed during discharge is insoluble and adheres to the electrode surfaces, which makes it possible to reverse the rxns and therefore recharge the battery - advantages: inexpensive, reliable at low temps, can be recharged thousands of times, provides a large initial supply of energy, long shelf-life - disadvantages: mass and must be disposed of properly Fuel Cells - a voltaic cell in which a fuel is continuously supplied form an external reservoir to the cell - very efficient at converting chemical energy into electrical energy (~90%) - most common: hydrogen-oxygen fuel cell that powers the space shuttle - simply the oxidation of H 2 to produce H2O - hydrogen flows into the anode and is oxidized - oxygen flows into the cathode and is reduced - cell contains KOH, which supplies OH - ions that participate in the electrode rxns - advantage: b/c the fuel for the cell is provided from an outside source, fuel cells never “run down” they keep producing energy as long as fuel is fed to them Electrolytic Cells - electrical energy is applied to cause nonspontaneous redox rxns to occur - consists of a source of direct electrical current that is connected to electrodes, which are immersed in the reactants - inert electrodes are often used b/c they do not participate in the redox rxn themselves but serve as surfaces on which the redox rxns take place - e- move from anode to cathode but are “pushed” through the circuit by the external source of electrical energy - electrolysis – the process by which electricity in an electrolytic cell is used to bring about a nonspontaneous chemical change - uses: electrolysis rxns including isolation of active metals, purifications of metals, and electroplating