Study Guide for Electrochemistry 6.Pdf

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Study Guide for CHEM 219/229 How to write a redox reaction using Electrochemical Cell notation: Terms to know: • Redox Reaction: A chemical reaction that involves the transfer of electrons between species, with one species losing electrons (being oxidized) and another species gaining electrons (being reduced). • Electrode: For the purposes of electrochemistry, a strip of metal. • Half- cell: An electrode submerged in a solution of ions of the same metal. • Electrochemical Cell: Two properly connected half cells. Two kinds: Galvanic and electrolytic o Galvanic Cell: An electrochemical cell that uses a chemical reaction to produce electrical energy. o Electrolytic Cell: An electrochemical cell that uses electrical energy to drive a chemical reaction. • Anode: The half-cell where oxidation takes place. • Cathode: The half-cell where reduction takes place. • Salt-bridge: A porous barrier (generally a U-tube) that allows ions to move between half-cells to complete the electrical circuit. General Cell Diagram: Electrode|Solution||Solution|Electrode • The single lines represent the phase boundary between each electrode and its solution • The double lines represent the salt bridge • The anode is always written on the left of the diagram, the cathode always on the right. • Example: o For the following reaction: Mg (s) + FeCl2 (aq) -> Fe (s) + MgCl2 (aq) o We start by assigning oxidation numbers: o Mg (s): . Mg O.N = 0 o FeCl2 (aq): . Fe O.N = +2 . Cl O.N = -1 o Fe(s): . Fe O.N = 0 o MgCl2 (s) . Mg O.N = +2 . Cl O.N = -1 o We can see that the Mg is being oxidized (it goes from a 0 to a +2 O.N), while the Fe is being reduced (goes from a +2 to a 0 O.N). • We can then break the redox reaction into half-reactions. One reduction half-reaction and one oxidation half-reaction: 2+ - Red: Mg (s) -> Mg (aq) + 2e 2+ - Ox: Fe (aq) + 2e -> Fe (s) • Notice that the Cl does not appear. Because it doesn’t change during the reaction, it is only a spectator ion and can be ignored. • Each of these half reactions can then be written as a half-cell, with the oxidation half-reaction being at the anode and reduction half-reaction being at the cathode: Anode|Solution||Solution|Cathode 2+ 2+ Fe(s)|Fe (aq)||Mg (aq)|Mg(s) .

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Cell Notation Practical Galvanic Cells -Batteries
Basic Redox Vocabulary • Write reactions for each of the following: • oxidation of metallic nickel by BiO+ • reduction of Zn2+ by hydroxide ion • reaction of Fe 2+ with Hg2+ 2+ • reaction of Cd with NO2 • AgI acting as an oxidizing agent toward Sn 2+ • What’s wrong with • The oxidation of Cr by Cl- • The reduction of Co 2+ by Ag+ Cell Notation • As was noted earlier, galvanic cells normally consist of two distinct regions, one housing the oxidation half and the other the reduction half. There is a simplified notation form that allows one to represent the cell easily( text p 798- 799). • The oxidation is written on the left and the reduction on the right. starting with the anode material and ending with the cathode material. • phase boundaries represented with single vertical lines “ |” • the physical separation between the two half cells is a double v ertical line “||” if it’s a salt bridge and with a single broken vertical line, “!”, if it’s a liquid junction • within each have cell, the species are written in a reactant-product order, separated by commas if they are in the same phase. Acid/base components should be included • The electrode material may be actively participating in the redox chemistry (active electrode) or merely providing surface for the electron transfer (passive or inert electrode, usually graphite or Pt) • Represent the following as galvanic cells(assume the reactions are spontaneous as written) • Tl(s) + Cd 2+ ó Tl + + Cd(s) - 2+ 2+ • Pb(s) + MnO4 ó Pb + Mn (acid) 2+ 4+ • O2(g) + Sn ó H2O + Sn Practical Galvanic Cells -batteries • Batteries represent the most common application of the electrochemical cell.
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