Electrochemistry Unit Test Review

Electrochemistry Unit Test Review Definitions Half-Reaction: Either the oxidation or reduction reaction component to a redox reaction. It’s found by finding the change in oxidation states of the chemicals in the redox reaction. Corrosion: An electrochemical process in which a metal reacts with substances in the environment, returning the metal to an ore like state. Spontaneous oxidation. Electrochemical Cell: A device that converts chemical energy into electrical energy, usually built with 2 different metals separated by an electrolyte (galvanic/voltaic cell Reducing Agent (Chemical): The chemical that donates electrons to cause the reduction. Always a reactant. Oxidizing Agent (Chemical): The chemical that accepts electrons to cause the oxidation. Always a reactant. Redox Reaction (Electrochemistry): The reactions where electrons are transferred. All single displacement reactions are redox reactions. Standard Reduction Potentials: The likelihood of an ion being reduced; an ion gaining electrons. Volta: Realized that if you separate the reactants from the products you can create an electrical potential (voltage). Balancing Redox Reactions 1. Write out the net ionic equation. 2. Separate the net ionic equation into an oxidation and reduction half reaction. 3. Balance each half reaction. a. Balance all atoms, leave O and H until end. b. Balance O with H2O. c. Balance H with H+. d. Balance the charge with electrons. 4. Multiply each half-reactions by a whole number so the total electrons lost and gained is equal. 5. Combine the half-reactions to remove the electrons. 6. If needed, add the right number of OH- to each side of the reaction, to remove any H+, if it occurs in a basic solution. Oxidation-Reduction Reactions • During chemical reactions, there’s usually a transfer of electrons from one particle to another. These reactions involving the transfer of electrons can be split into 2 half reactions. • Oxidation: The half reaction where a particle loses electrons. • Reduction: The half reaction where a particle gains electrons. Oxidation Number • The apparent charge of an atom if all electrons are transferred to the most electronegative atom. • For ions like copper or iron, the oxidation number is equal to its charge. • For all elements the oxidation number is equal to 0. • For complex compounds the oxidation number must be calculated. • The total oxidation state for an uncharged particle must be 0. • Many atoms have more than 1 oxidation number. Standard Reduction Potentials • Measure the potential of a cell. • The reduction reaction always occurs at the cathode which is the positive terminal of a galvanic cell (the terminal which electrons travel towards). • The metal with the highest reduction potential is the cathode because the reaction is spontaneous. The voltage for the cell can then be calculated. ΔEo = Eor CATHODE - Eor ANODE Reaction Spontaneity • If the ΔEo is positive, the reaction is spontaneous. • If the ΔEo is negative, the reaction isn’t spontaneous (a quantitative way to prove is a reaction occurs). Batteries • Refers to a number of cells connected in series with one another. • Frequently used interchangeably with the term cell as most common batteries are actually cells (AA, AAA, C, D). • Examples of actual batteries are 6V and 9V batteries, as well as car batteries (composed of multiple cells). Types of Cells • Primary Cells: Can only be used once and not recharged. • Secondary Cells: Can be reused and are rechargeable. • Dry Cells: Thick electrolyte paste (household batteries). • Wet Cells: Aqueous solutions (sulphuric acid in car batteries). • M stands for an alloy including a number of rare earth metals. Predicting Reactions • In a galvanic/voltaic cell, the reaction is always spontaneous because all the reactions for both the all possible reactions are present. • A galvanic cell will go in the direction that produces the most stable products. • If all the reactants are not available (less stable option), the ΔEo can be calculated for the reaction and used to predict if it’s spontaneous or not. • Even if a reaction is not spontaneous, it can still occur, just not spontaneously. Electrolysis • An electrolytic cell is a different type of cell which an electrical energy source replaces the voltage meter. If it’s power can create a larger potential difference in the opposite direction, the reaction can go in the opposite, not spontaneous direction. • By overcoming the potential difference/cell potential, you can force the reaction to happen. In calculations when negative, it’s not both reactions, only one, which is forced to happen using a power source. • In some cases, for the reaction to happen, you must had H+ ions (strong acid) and you need to increase the voltage. • Like galvanic cells, inert electrodes can be used and the electrolyte can be mixed. • Ex., electrolysis (decomposition) of water.