Chem 132- General Chemistry II Unit 10 :Electrochemistry

Electrochemistry The study of the relationship between chemical change and electrical work It is typically investigated through use of electrochemical cells Two types of electrochemical cells: Those that release free energy from a spontaneous rxn to produce electricity (voltaic cells) Those that absorb free energy from a source of electricity to drive a non-spontaneous reaction (electrolytic cells) Both instances involvement movement of electrons from one chemical species to another in a REDOX reaction

Oxidation-Reduction Processes Oxidation-reduction processes are responsible for many types of chemical changes Oxidation - defined by one of the following: loss of electrons loss of hydrogen atoms gain of oxygen atoms Example: NaNa+ + e- Oxidation half reaction

Oxidation-Reduction Processes Reduction - defined by one of the following: gain of electrons gain of hydrogen loss of oxygen Example: Cl + e-  Cl- Reduction half reaction Cannot have oxidation without reduction

Oxidation and Reduction as Complementary Processes Na + Cl Na+ + Cl- Na  Na+ + e- Cl + e-  Cl- Oxidation half reaction Reduction half reaction Overall reaction One way to balance REDOX reactions is by the half reaction method: Divides the overall reaction into oxidation and reductions reactions

Oxidation numbers (O.N.) Another way to define a redox reaction is one in which the oxidation number of the species changes Oxidation is represented by an increase in O.N. Reduction is represented by a decrease (reduction) in O.N. Na  Na+ + e- Oxidation: 0  +1 Cl + e-  Cl- Reduction: 0  -1

Oxidizing Agent Reducing Agent Is reduced Gains electrons Causes oxidation Reducing Agent Is oxidized Loses electrons Causes reduction

Voltaic Cells: Zinc-Copper Cell Voltaic cell - electrochemical cell that converts stored chemical energy into electrical energy Let’s consider the following reaction: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) Is Zn oxidized or reduced? Oxidized Copper is reduced

Voltaic Cell Generating Electrical Current Zn  Zn2+ + 2e- Oxidation Cu2+ + 2e-  Cu Reduction

Voltaic Cells Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) If the two reactants are placed in the same flask they cannot produce electrical current A voltaic cell separates the two half reactions This makes the electrons flow through a wire to allow the oxidation and reduction to occur  Video

Voltaic Cell Generating Electrical Current Each half cell is originally neutral but as Zn metal anode losses e-s & Zn2+ enters solution net +ve charge would develop Similarly, as copper solution gains e-s  Cu metal deposited on cathode  solution would develop a net –ve charge Charge imbalance would stop cell operation. This avoided by a salt bridge. Voltaic Cell Generating Electrical Current

Things to note about the salt bridge completes the circuit & allows the flow of ions prevents mixing of the 2 solutions is necessary for the voltaic cell to work allows electrical contact between the two solutions in the half-cells. the ions in the salt bridge DO NOT participate in the half reactions

Zinc & copper are active electrodes  metal bars themselves are components of the half reactions Inactive electrodes e.g. platinum & graphite  conduct e-s into or out of half cells but cannot participate in half reactions

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 There is a commonly used short hand notation for voltaic cells. The Zn-Cu cell provides a good example.

Silver Battery Batteries use the concept of the voltaic cell Modern batteries are: Smaller Safer More dependable

The Copper - Silver Cell Cell components: A Cu strip immersed in 1.0 M copper (II) sulfate. A Ag strip immersed in 1.0 M silver (I) nitrate. A wire and a salt bridge to complete the circuit. The initial cell voltage is 0.46 volts.

The Copper - Silver Cell

The Copper - Silver Cell Compare the Zn-Cu cell to the Cu-Ag cell The Cu electrode is the cathode in the Zn-Cu cell. The Cu electrode is the anode in the 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 These experimental facts demonstrate that Cu2+ is a stronger oxidizing agent than Zn2+. In other words Cu2+ oxidizes metallic Zn to Zn2+. Similarly, Ag+ is is a stronger oxidizing agent than Cu2+. Because Ag+ oxidizes metallic Cu to Cu 2+. If we arrange these species in order of increasing strengths, we see that:

Why does a voltaic cell work? The spontaneous reactions occurs as a result of the different abilities of metals to give up their e-s AND the ability of e-s to flow through the circuit Example: Zn/Cu cell Zn is a stronger reducing agent (oxidized easily) gives up its e-s more readily than Cu It produces e-s more readily The electrons in Zn electrode exerts a greater electron “pressure” than those of Cu electrode  greater potential energy (referred to as electrical potential) to “push” electrons through the circuit

Cell Potential: Output of a Voltaic Cell Purpose of a voltaic cell is to convert the free energy change of spontaneous rxn into electrical energy (e-s moving through an external circuit) This electrical energy can do work & is proportional to the difference in electrical potential b/w the two electrodes This difference in electrical potential of the electrodes is the cell potential (Ecell) OR voltage OR electromotive force (emf)

Cell Potential Electrons flow spontaneously from –ve  +ve electrode i.e. towards the electrode with the more +ve electrical potential So, for a spontaneous cell reaction, the difference in electrical potential of the electrodes is positive: Ecell > 0 for a spontaneous reaction The more +ve Ecell is, the more work the cell can do Ecell < 0 is associated with a non-spontaneous reaction Ecell = 0 means the reaction has reached equilibrium and the cell can do no more work

Standard Electrode Potential The measured potential of a voltaic cell is affected by: changes in concentration as the reaction proceeds energy losses due to heating of the cell and external circuit To compare the output of different cells we obtain a standard cell potential (E0cell) Cell potential measured at temperature of 298K All components must be in their standard states: 1atm for gases 1M solutions pure solids for electrodes

Standard Electrode Potential Example: Zn/Cu2+ cell [Zn2+] = [Cu2+] = 1.0M Temp. = 298K Pure copper and zinc electrodes Cell potential = 1.10V

Standard Half-cell Potentials Just as each half-reaction makes up the overall reaction, the potential of each half cell makes up part of the overall cell potential Standard electrode potential (E0half-cell) By convention a standard electrode potential always refers to the half reaction written as a reduction

Standard Half-cell Potentials Voltaic cell half reactions are: But electrode potentials given for: Zn 2+ (aq)  + 2e-  Zn(s) E0zinc = -0.76V Cu 2+(aq) + 2e-  Cu(s) E0copper = + 0.34 For any voltaic cell the standard electrode potential is given by: E0cell = E0cathode(reduction) - E0anode (oxidation) Cu2+ + 2e-  Cu Reduction Zn  Zn2+ + 2e- Oxidation

The hydrogen electrode The half cell potentials such as E0zinc and E0copper are values relative to that of a standard hydrogen electrode (SHE): 2H+(aq:1M) + 2e-  H2 (g:1atm) E0reference= 0.00V The SHE is assigned an arbitrary voltage of 0.000000… V Can therefore construct a voltaic cell consisting of this reference cell and & another half- cell whose potential we want to find

The Zinc-SHE Cell The initial cell voltage is 0.763 volts. A Zn strip immersed in 1.0 M zinc (II) sulfate. The other electrode is the Standard Hydrogen Electrode. A wire and a salt bridge to complete the circuit. The initial cell voltage is 0.763 volts.

The Zinc-SHE Cell The cathode is the Standard Hydrogen Electrode. In other words Zn reduces H+ to H2. The anode is Zn metal. Zn metal is oxidized to Zn2+ ions.

The Copper-SHE Cell The initial cell voltage is 0.337 volts. A Cu strip immersed in 1.0 M copper (II) sulfate. The other electrode is a Standard Hydrogen Electrode. A wire and a salt bridge to complete the circuit. The initial cell voltage is 0.337 volts.

The Copper-SHE Cell In this cell the SHE is the anode The Cu2+ ions oxidize H2 to H+. The Cu is the cathode. The Cu2+ ions are reduced to Cu metal.

Uses of Standard Electrode Potentials 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. 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 Standard Electrode Potentials 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 denote that the reaction tends to occur to the right. The larger the value, the greater the tendency to occur to the right. It is the opposite for negative values of Eo.

Uses of Standard Electrode Potentials Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously. Example 21-3: Will silver ions, Ag+, oxidize metallic zinc to Zn2+ ions, or will Zn oxidize metallic Ag+ ions toAg? Steps for obtaining the equation for the spontaneous reaction.

Uses of Standard Electrode Potentials 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. 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.

Uses of Standard Electrode Potentials We do not multiply the potentials by the numbers used to balance electron transfer. The reason is that each potential represents a tendency for a reaction to occur relation to the SHE and does not depend on how many times it occurs . An electrical potential is an intensive property

Electrode Potentials for Other Half-Reactions Example 21-4: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution?

Be sure know how to draw and label a voltaic cell Label anode and cathode Direction of flow of electrons ( always towards the cathode) Direction of flow of ions in the salt bridge Write half reactions that occur at each electrode You should also know the functions of the salt bridge Calculate the Eocell from reduction potential data The trend: