01 / 52Understanding Electrochemistry: Reactions, Cells, and Conduction
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Understanding Electrochemistry: Reactions, Cells, and Conduction
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Electrochemistry • Electrochemical reactions are oxidation-reduction reactions. • The two parts of the reaction are physically separated. • oxidation occurs at one cell • reduction occurs in the other cell • There are two kinds electrochemical cells. • Electrolytic cells - nonspontaneous chemical reactions • Voltaic or galvanic cells - spontaneous chemical reactions
Electrical Conduction • Ionic or electrolytic conduction • Ionic motion transports the electrons • Positively charged ions, cations, move toward the negative electrode. • Negatively charged ions, anions, move toward the positive electrode.
Electrodes • Conventions for electrodes: • Cathode - electrode at which reduction occurs • Anode - electrode at which oxidation occurs • Inert electrodes do not react with the liquids or products of the electrochemical reaction. • Graphite and Platinum are common inert electrodes.
Voltaic or Galvanic Cells • Electrochemical cells in which a spontaneous chemical reaction produces electrical energy. • Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference. • Examples include:
The Construction of Simple Voltaic Cells • Half-cell contains the oxidized and reduced forms of an element (or other chemical species) in contact with each other. • Simple cells consist of: • two pieces of metal immersed in solutions of their ions • wire to connect the two half-cells • salt bridge to • complete circuit • maintain neutrality • prevent solution mixing
The Zinc-Copper Cell • Cell components: Cu strip immersed in 1.0 M copper (II) sulfate Zn strip immersed in 1.0 M zinc (II) sulfate wire and a salt bridge to complete circuit • Initial voltage is 1.10 volts
The Zinc-Copper Cell • In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).
The Zinc-Copper Cell • Short hand notation for voltaic cells • Zn-Cu cell example
The Copper - Silver Cell • Cell components: Cu strip immersed in 1.0 M copper (II) sulfate Ag strip immersed in 1.0 M silver (I) nitrate wire and a salt bridge to complete circuit • Initial voltage is 0.46 volts
The Copper - Silver Cell • Compare the Zn-Cu cell to the Cu-Ag cell Cu electrode is cathode in Zn-Cu cell Cu electrode is anode in Cu-Ag cell • Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.
The Copper - Silver Cell • Demonstrates that Cu2+ is a stronger oxidizing agent than Zn2+ Cu2+ oxidizes metallic Zn to Zn2+ • Ag+ is is a stronger oxidizing agent than Cu2+ Ag+ oxidizes metallic Cu to Cu2+ • Arrange these species in order of increasing strengths
Standard Electrode Potential • Establish an arbitrary standard to measure potentials of a variety of electrodes • Standard Hydrogen Electrode (SHE) • assigned an arbitrary voltage of 0.000000… V
The Zinc-SHE Cell • Cell components: Zn strip immersed in 1.0 M zinc (II) sulfate other electrode is a SHE wire and a salt bridge to complete circuit • Initial voltage is 0.763 volts
The Zinc-SHE Cell • SHE is the cathode Zn reduces H+ to H2 • Zn is the anode
The Copper-SHE Cell • Cell components: Cu strip immersed in 1.0 M copper (II) sulfate other electrode is a SHE wire and a salt bridge to complete circuit • Initial voltage is 0.337 volts
The Copper-SHE Cell • SHE is the anode Cu2+ ions oxidize hydrogen to H+ • Cu is the cathode
The Electromotive (Activity) Series of the Elements • Electrodes that force the SHE to act as an anode are assigned positive standard reduction potentials. • Electrodes that force the SHE to act as the cathode are assigned negative standard reduction potentials. • Table 20.1 in Text lists the Standard Reduction Potentials
Uses of the Electromotive Series • Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written. • For example, the half-reaction for the standard potassium electrode is: • The large negative value tells us that this reaction will occur only under extreme conditions.
Uses of the Electromotive Series • Compare the potassium half-reaction to fluorine’s half-reaction: • The large positive value denotes that this reaction occurs readily as written. • Positive E0 values tell us that the reaction tends to occur to the right • larger the value, greater tendency to occur to the right • Opposite for negative values
Uses of the Electromotive Series • Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously. • Steps for obtaining the equation for the spontaneous reaction.
Uses of the Electromotive Series • Choose the appropriate half-reactions from a table of standard reduction potentials. • Write the equation for the half-reaction with the more positive E0 value first, along with its E0 value. • Write the equation for the other half-reaction as an oxidation with its oxidation potential, i.e. reverse the tabulated reduction half-reaction and change the sign of the tabulated E0.
Uses of the Electromotive Series • Balance the electron transfer. • Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E0cell is positive, which indicates that the forward reaction is spontaneous.
Electrode Potentials for Other Half-Reactions • Example: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese ions to permanganate ions in acidic solution? • Follow the steps outlined in the previous slides. E0 values are not multiplied by any stoichiometric relationships in this procedure.
Electrode Potentials for Other Half-Reactions • Example: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese ions to permanganate ions in acidic solution?
Electrode Potentials for Other Half-Reactions • Example: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese ions to permanganate ions in acidic solution?
Electrode Potentials for Other Half-Reactions • Example: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese ions to permanganate ions in acidic solution?
Electrode Potentials for Other Half-Reactions • Example: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese ions to permanganate ions in acidic solution? • Thus permanganate ions will oxidize iron (II) ions to iron (III) and are reduced to manganese (II) ions in acidic solution.
Effect of Concentrations (or Partial Pressures) on Electrode Potentials - Nernst Equation • Standard electrode potentials are determined at thermodynamic standard conditions. 1 M solutions 1 atm of pressure for gases liquids and solids in their standard states temperature of 250 C
Effect of Concentrations (or Partial Pressures) on Electrode Potentials - Nernst Equation • The Nernst equation can also be used to calculate the potential for a cell that consists of two nonstandard electrodes. • Example: Calculate the initial potential of a cell that consists of an Fe3+/Fe2+ electrode in which [Fe3+]=1.0 x 10-2M and [Fe2+]=0.1 M connected to a Sn4+/Sn2+ electrode in which [Sn4+]=1.0 M and [Sn2+]=0.10 M . A wire and salt bridge complete the circuit.
