Chapter 17: Electrochemistry Notes PDF

Chapter 17: Electrochemistry  Electrochemistry Electrochemistry is the study of redox reactions that produce or require an electric current. The conversion between chemical energy and electrical energy is carried out in an electrochemical cell. Spontaneous redox re...

Chapter 17: Electrochemistry  Electrochemistry Electrochemistry is the study of redox reactions that produce or require an electric current. The conversion between chemical energy and electrical energy is carried out in an electrochemical cell. Spontaneous redox reactions take place in a galvanic/voltaic cell. Nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy. 2  Oxidation-Reduction Reactions in which electrons are transferred from one atom to another are called oxidation–reduction reactions. Atoms that lose electrons are being oxidized; atoms that gain electrons are being reduced. 3  Oxidation-Reduction Oxidation is the process that occurs when the oxidation number of an element increases; an element loses electrons; a compound adds oxygen; a compound loses hydrogen; or a half-reaction has electrons as products. Reduction is the process that occurs when the oxidation number of an element decreases; an element gains electrons; a compound loses oxygen; a compound gains hydrogen; or a half-reaction has electrons as reactants. 4 A Spontaneous Redox Reaction: Zn(s) + Cu2+(aq) → Zn2+ + Cu(s)  Voltaic (Galvanic) Cells: Spontaneous Redox Reactions Electrical current: The amount of electric charge that passes a point in a given period of time Whether as electrons flowing through a wire or as ions flowing through a solution Redox reactions involve the movement of electrons from one substance to another. Therefore, redox reactions have the potential to generate an electric current. A spontaneous redox reaction does not require external energy to proceed. The reaction’s ΔG is negative. Voltaic (galvanic) cells produce an electrical current from spontaneous redox reactions. To use that current, we need to separate the place where oxidation is occurring from the place where reduction is occurring.  Electrochemical Cells Oxidation and reduction half-reactions are kept as separate in half-cells in an electrochemical cell. To constitute an electrical circuit: Electron flow through a wire along with Ions (electrolyte) flowing through a solution via the salt bridge. The flow of electrons require a conductive solid electrode to allow the transfer of electrons either through: An external circuit or Metal or graphite electrode An electrochemical cell requires the exchange of ions between the two half-cells of the system via a salt bridge.  Voltaic Cell The salt bridge is required to complete the circuit and maintain charge balance.  Voltage and Current Voltage is the difference in potential Current is the number of electrons that energy between the reactants and flow through the system per second. products. It is also called the potential Unit = ampere difference. Unit = volt 1 A of current = 1 coulomb of charge flowing each second 1 V = 1 J of energy per coulomb of 1 A = 6.242 × 1018 electrons per second charge The voltage needed to drive Electrode surface area dictates the electrons through the external number of electrons that can flow. circuit Larger batteries produce larger currents. The amount of force pushing the electrons through the wire is called the electromotive force, emf.  Cell Potential The difference in potential energy between the anode and the cathode in a voltaic cell is called the cell potential. The cell potential depends on the relative ease with which the oxidizing agent is reduced at the cathode and the reducing agent is oxidized at the anode. The cell potential under standard conditions is called the standard emf, E°cell. 25 °C, 1 atm for gases, 1 M concentration of solution Sum of the cell potentials for the half-reactions  Cell Notation Shorthand description of a voltaic cell is written as follows: electrode | electrolyte || electrolyte | electrode Oxidation half-cell on the left; reduction half-cell on the right Single line | = phase barrier If multiple electrolytes in same phase, a comma is used rather than | Often use an inert electrode Double line || = salt bridge Example:  Voltaic Cell Anode = Zn(s) Cathode = Cu(s) The anode is oxidized to Cu2+ ions are reduced at Zn2+ ions. the cathode.  Electrodes When the half-reaction involves a gas, an inert electrode can be used. An inert electrode, such as platinum, is one that does not participate in the reaction but provides a surface for the transfer of electrons to take place on. Hydrogen electrode written as a cathode: H+(aq) | H2(g) | Pt Hydrogen electrode written as an anode: Pt | H2(g) | H+(aq) 13  Electrochemical Cell Notation with an Inert Electrode The half-reaction involves reducing Mn oxidation state from +7 to +2. An inert electrode (Pt) will provide a surface for the electron transfer without reacting with the MnO4−. Fe(s) | Fe2+(aq) || MnO4– (aq), Mn2+(aq), H+(aq) | Pt(s) 14  Cell potential A measure of the driving force of the cell reaction. Work is needed to move electrons in a wire or to move ions through a solution to an electrode. With electricity, an electric charge moves from a point at high electric potential to a point at low electric potential. Electrical work expended in moving a charge through a conductor: Electrical work = charge × potential difference Faraday constant (F): the magnitude of charge on one mole of electrons; 9.6485 ×104 C/mol e- 15  Cell potential Cell potential (Ecell): The maximum potential difference between the electrodes of a voltaic cell. Measured by an electronic digital voltmeter. Maximum work obtainable from a voltaic cell: 𝑤𝑤max = −𝑛𝑛𝑛𝑛𝑛𝑛cell n: number of moles of electrons transferred in the overall cell equation F: Faraday constant 16  Standard Reduction Potential The absolute tendency of a half-reaction standard reduction potential cannot be measured. Only the potentials relative to another half-reaction can be measured. To overcome this limitation, a standard half-reaction for the reduction of H+ to H2 is selected and assigned a potential difference of 0 V. Standard hydrogen electrode, SHE 17 18  Half-cell Potentials Standard reduction potentials compare the tendency for a particular reduction half-reaction to occur relative to the reduction of H+ to H2. Half-reactions with a stronger tendency toward reduction than the SHE have a positive value for E°red. Half-reactions with a stronger tendency toward oxidation than the SHE have a negative value for E°red. For an oxidation half-reaction, E°oxidation = −E°reduction. 19 Table of Standard Reduction Potentials (from Pearson Education, Inc.) A redox reaction will be spontaneous when there is a strong tendency for the oxidizing agent to be reduced and the reducing agent to be oxidized. Higher on the table of standard reduction potentials = stronger tendency for the reactant to be reduced Lower on the table of standard reduction potentials = stronger tendency for the product to be oxidized  Calculating Cell Potentials under Standard Conditions Cell potentials are intensive properties of matter. Because cell potentials are intensive physical properties, when determining the cell potential, do not multiply the half-cell E° values, even if you need to multiply the half-reactions to balance the redox equation. The cell potential of an electrochemical cell (Eocell) is the difference between the electrode potential of the cathode and that of the anode. Cell potentials can be determined using the following equation: E°cell = E°cathode – E°anode 21  Predicting Spontaneity of Redox When E is ° Reactions cell positive, the redox reaction of the cell is spontaneous. Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) spontaneous When E°cell is negative, the redox reaction of the cell is nonspontaneous. Cu(s) + Zn2+(aq) → Cu2+(aq) + Zn(s) nonspontaneous Cu2+(aq) + 2 e− → Cu(s) E°red = +0.34 V Zn2+(aq) + 2 e− → Zn(s) E°red = −0.76 V 22 Example 23  E°cell, ΔG°, and K For a spontaneous reaction, one that proceeds in the forward direction with the chemicals in their standard states, ΔG°: negative E°: positive K>1 ΔG° = −RT ln K = −nFE°cell n = the number of electrons F = Faraday’s constant = 96,485 C/mol e− 24 Dependence of Cell Potential on Concentration ΔG = ΔG° + RT ln Q (–) nFEcell = (–) nFE°cell + RT ln Q RT Ecell = E°cell – ln Q This is the Nernst equation. nF At 25 °C (T), Faraday’s constant (F) = 96,500 C/mol e– n = the number of electrons Converting from ln to log (2.303), the Nernst equation becomes 0.0592 V Ecell = E°cell – log Q n 25  Corrosion Corrosion is the spontaneous oxidation of a metal by chemicals in the environment. Mainly O2 Because many materials used are active metals, corrosion can be a very big problem. Metals are often used for their strength and malleability, but these properties are lost when the metal corrodes. For many metals, the product of corrosion does not adhere to the metal, and as it flakes off more metal can corrode. 26  Reduction of O2 O2 is very easy to reduce in moist conditions. O2(g) + 2 H2O(l) + 4 e− → 2 OH−(aq) E° = 0.40 V O2 is even easier to reduce under acidic conditions. O2(g) + 4 H+ + 4 e− → 2 H2O(l) E° = 1.23 V Because the reduction of most metal ions lies below O2 on the table of standard reduction potentials, the oxidation of those metals by O2 is spontaneous. 27  Rusting At the anodic regions, Fe(s) is oxidized to Fe2+. The electrons travel through the metal to a cathodic region where O2 is reduced. In acidic solution from gases dissolved in the moisture The Fe2+ ions migrate through the moisture to the cathodic region, where they are further oxidized to Fe3+, which combines with the oxygen and water to form rust. Rust is hydrated iron(III) oxide, Fe2O3 · nH2O. Moisture must be present. Water acts as a reactant. It is required for ion flow between cathodic and anodic regions. 28  Preventing Corrosion One way to reduce or slow corrosion is to coat the metal surface to keep it from contacting corrosive chemicals in the environment. Paint Some metals, such as Al, form an oxide that strongly attaches to the metal surface, preventing the rest from corroding. Another method to protect one metal is to attach it to a more reactive metal that is cheap. 29 Voltaic vs Electrolytic Cells 30  Electrolysis Electrolysis is the process of producing a chemical change in an electrolytic cell. Electrolytic cells can be used to separate elements from their compounds.  Electrolysis In electrolysis we use electrical energy to overcome the energy barrier of a nonspontaneous reaction, allowing it to occur. The reaction that takes place is the opposite of the spontaneous process. 2 H2(g) + O2(g) → 2 H2O(l) spontaneous 2 H2O(l) → 2 H2(g) + O2(g) electrolysis Some applications of electrolysis are the following: (1) Metal extraction from minerals and purification (2) Production of H2 for fuel cells (3) Metal plating  Electrolytic Cells The source of energy is a battery or DC power supply. The positive terminal of the source is attached to the anode. The negative terminal of the source is attached to the cathode. Electrolyte can be either an aqueous salt solution or a molten ionic salt. Cations in the electrolyte are attracted to the cathode and anions are attracted to the anode. Cations pick up electrons from the cathode and are reduced; anions release electrons to the anode and are oxidized.  Electrolysis of Aqueous Solutions Possible cathode reactions: Reduction of cation to metal Reduction of water to H2 2 H2O + 2 e− → H2 + 2 OH− E° = −0.83 V at standard conditions E° = −0.41 V at pH 7 Possible anode reactions: Oxidation of anion to element Oxidation of H2O to O2 2 H2O → O2 + 4 e− + 4H+ E° = −1.23 V at standard conditions E° = −0.82 V at pH 7 Oxidation of electrode: Particularly Cu Graphite doesn’t oxidize Half-reactions that lead to least negative Ecell will occur. Unless overvoltage changes the conditions Electrolysis of Water  Electrolysis of Pure Compounds The compound must be in molten (liquid) state. Electrodes are normally graphite. Cations are reduced at the cathode to metal element. Anions are oxidized at the anode to nonmetal element. Electrolysis of NaCl(l)  Electroplating In electroplating, the work piece is the cathode. Cations are reduced at cathode and plate to the surface of the work piece. The anode is made of the plate metal. The anode oxidizes and replaces the metal cations in the solution.  Mixtures of Ions and Electrolysis When more than one cation is present, the cation that is easiest to reduce will be reduced first at the cathode. Least negative or most positive E°red When more than one anion is present, the anion that is easiest to oxidize will be oxidized first at the anode. Least negative or most positive E°ox  Stoichiometry of Electrolysis In an electrolytic cell, the amount of product made is related to the number of electrons transferred. Essentially, the electrons are a reactant. The number of moles of electrons that flow through the electrolytic cell depends on the current and length of time. 1 amp = 1 coulomb of charge/second 1 mole of e− = 96,485 coulombs of charge Faraday’s constant Conceptual plan: time (in seconds) → coulombs → moles of electrons → moles of metal → grams of metal 39