Introduction to electrochemistry Oxidation reduction reactions involve energy changes. Because these reactions involve electronic transfer, the net release or net absorption of energy can occur in the form of electrical energy rather than as heat. This property allows for great many practical applications of redox reactions. The branch of chemistry that deals with electricity related applications of oxidation reduction reactions is called electrochemistry.
Electrochemical Cells Redox reactions involve transfer of electrons from oxidized to reduced substance If two substances in contact, a transfer of energy as heat happens as well
Zinc strip in contact with copper (II) sulfate solution Zinc strip loses e- to copper (II) ions in solution Copper (II) accept electrons and fall out of solution as copper atoms As electrons are transferred, energy released as heat (rise in temperature)
If we separate substances the e- transfer comes with transfer of electricity instead of heat One way to separate is with porous barrier Prevents metal atoms of one half-reaction from mixing with metal atoms of the other
Ions in two solutions can move through porous barrier e- can be transferred from one side to the other through connecting wire Current moves in circuit so movement of e- balanced by movement of ions in solution
Altering system in Figure 19-7 so electrical current is produced involves separating copper and zinc (Figure 19-8) Zn strip is in aqueous solution of ZnSO 4 Cu strip is in solution of CuSO 4
Both solutions are electrolytes Electrode à conductor used to establish electrical contact with nonmetallic part of a circuit, such as an electrolyte Zn and Cu strips are electrodes
Half-cell à single electrode immersed in a solution of its ions Zn strip in ZnSO 4 (aq) is anode à where oxidation takes place Cu strip in CuSO 4 (aq) is cathode à where reduction takes place
The complete cell Copper half-cell written as Cu 2+ /Cu Zinc half-cell written as Zn 2+ /Zn Two half-cells together make electrochemical cell à system of electrodes and electrolytes in which either chemical reactions produce electrical energy or an electric current produces chemical change
Electrochemical cell can be represented by following notation: Anode І Cathode Cell made of zinc and copper would be written Zn І Cu 2 types of electrochemical cells Voltaic (also called galvanic) electrolytic
Voltaic cells Voltaic cells use spontaneous oxidation reduction reactions to convert Chemical Energy into electrical energy. Voltaic cells are also called galvanic cells. The most common application of voltaic cells is in batteries.
Voltaic Cells Voltaic cell à if redox reaction in electrochemical cell happens spontaneously and produces electrical energy Cations in solution reduced when they gain e- at surface of cathode to become metal atoms
Half reaction for cathode: Half reaction for the anode:
e- given up at anode pass along external wire to cathode
Movement of e- through wire must be balanced by movement of ions in solution Anions move to anode to replace negative e- moving away
Cations move to cathode as positive charge lost through reduction
Dry cells are common sources of electrical energy They are voltaic cells Three most common types: Zinc-carbon battery Alkaline battery Mercury battery
Zinc-Carbon Dry Cells Consist of zinc container (anode) Filled with moist paste of MnO 2, graphite, and NH 4 Cl When external circuit closed, Zn oxidized at negative electrode (anode)
e- move across circuit and reenter cell through carbon rod (cathode) MnO 2 reduced in presence of water
Alkaline Batteries Do not have carbon rod cathode Allows them to be smaller Uses paste of Zn metal and KOH instead of solid metal anode Half-reaction at anode:
Reduction half-reaction (at cathode) same as Zn-C dry cell
Mercury Batteries Tiny batteries in hearing aids, calculators, etc. Anode half-reaction same as alkaline dry cell
Cathode half-reaction is as follows:
Fuel cells Voltaic cell where reactants continuously supplied and products continuously removed Unlike battery, could in principle work forever Very efficient and have very low emissions
Corrosion and its prevention Corrosion is electrochemical process that has large economic impact About 20% all iron and steel produced used to repair or replace corroded structures Rust, hydrated iron (III) oxide, forms by the following reaction 4Fe(s) + 3O 2 (g) + xh 2 O(l) à 2Fe 2 O 3 xh 2 O(s)
4Fe(s) + 3O 2 (g) + xh 2 O(l) à Amount of hydration varies Affects the color of rust formed Electrochemical reactions: Anode: Fe(s) à Fe 2+ (aq) + 2e - Cathode: O 2 (g) + 2H 2 O(l) + 4e - à 4OH - (aq) Anode and cathode reactions occur in different areas of metal surface
Circuit completed by electronic flow through metal Acts like wire in electrochemical cell Waters serves as salt bridge For corrosion to occur, water and oxygen must be present
When iron exposed to water and oxygen, metal at anode oxidized to Fe 2+ ions Electrons released travel along metal to cathode Fe 2+ ions travel along moisture to cathode At cathode, Fe 2+ further oxidized to Fe 3+
Corrosion prevention Coat steel with zinc Galvanizing Zinc more easily oxidized than iron Will react before the iron is oxidized Cathodic protection More easily oxidize metal = Sacrificial anode
Alaskan Oil pipeline Zinc connected to pipe by wire As zinc (anode) corrodes it gives electrons to the cathode (steel) As a dissolves, zinc needs to be replaced
Electrode Potential There are 2 electrodes, Zn and Cu These 2 metals each have different tendencies for accepting electrons Reduction potential à tendency for halfreaction of a metal to happen as a reduction half-reaction in an electrochemical cell
There are 2 half-cells 1. Strip of Zn in solution of ZnSO 4 2. Strip of Cu in CuSO 4 Electrode potential à difference in potential between an electrode and its solution
When these 2 half-cells are connected and reaction begins, difference in potential is observed between the electrodes This voltage is a measure of energy required to move certain electric charge between electrodes Potential difference measured in volts
Voltmeter connected across the Cu І Zn voltaic cell measures potential difference of about 1.10 V when solution concentrations are each 1 M
Potential difference measured across voltaic cell roughly equals sum of electrode potentials for the 2 half-reactions Easy to measure voltage across voltaic cell No way to measure individual electrode potential directly b/c there is no transfer of e- unless both anode and cathode are connected in complete circuit
Relative value for potential of half-reaction can be determined by connecting it to standard half-cell as reference This is called a standard hydrogen electrode, or SHE
Platinum electrode dipped in 1.00 M acid solution Surrounded by hydrogen gas at 1 atm pressure and 25 C Other electrodes ranked according to ability to reduce H under these conditions
Anodic reaction for SHE is described by forward half-reaction in following equation Cathodic half-reaction is reverse Random potential of 0.00 V assigned to both of these reactions Any voltage measurement is credited to the half-cell connected to SHE Standard electrode potential, E⁰ à a half-cell potential measured relative to a potential of zero for the SHE
Electrode potentials expressed as potentials for reduction Provide reliable indication of tendency of a substance to be reduced Positive E⁰ values indicate H more willing to give up e- than other electrode Half-reactions with positive reduction potentials are favored
Half-reactions with negative reduction potentials are not favored These half-reactions prefer oxidation over reduction Negative E⁰ values indicate that the metal or other electrode is more willing to give up e- than hydrogen
When half-reaction is written as oxidation rxn, sign of electrode potential changed as shown for redox half-rxns for Zn
To measure reduction potential of Zn half-cell, it is connected to SHE Potential difference is -0.76 V Negative number indicates e- flow through external circuit from Zn electrode (Zn oxidized) to H electrode (H ions reduced)
Copper half-cell paired with SHE gives E⁰ of +0.34 V This positive number indicates Cu 2+ (aq) more readily reduced than H + (aq)
Standard electrode potentials can be used to predict if redox reactions will happen naturally A naturally occurring rxn will have positive value for E⁰ cell according to the following equation
When evaluating the 2 half-rxns of a cell, the half-rxn with more negative E⁰ is the anode Oxidation happens at the anode, so half-cell rxn is reverse of reduction rxn in Table 19-4 When rxn reversed, actual half-cell potential is negative of E⁰ The total potential of a cell is calculated by subtracting standard reduction potential for rxn at anode (E⁰ anode ) from standard reduction potential for rxn at cathode (E⁰ cathode )
Example Cells made from Fe in solution of Fe(NO 3 ) 3 and Ag in solution of AgNO 3 Table 19-4 gives
Fe in Fe(NO 3 ) 3 is anode b/c it has lower reduction potential Ag in AgNO 3 is cathode Overall cell reaction is
Reduction of Ag ions multiplied by 3 so number of e- lost in that half-reaction equals number of e- gained in oxidation of iron Standard reduction potentials for anode and cathode are as follows NOTE When a half-reaction is multiplied by a constant, the E⁰ value is NOT multiplied.
So Potential for this cell can be calculated as follows If calculated value for E⁰ cell were negative, the reaction would NOT happen naturally It would not happen in a voltaic cell It could be made to happen in an electrolytic cell
Some oxidation reduction reactions do not occur spontaneously but can be driven by electrical energy. If electrical energy is required to produce a redox reaction and bring about a chemical change an electrochemical cell, it is an electrolytic cell. Most commercial uses of redox reactions make use of electrolytic cells. Electrolytic cells
How electrolytic cells work Comparison of electrolytic in voltaic cells seen in figure 13
Electrode of cell connected to negative terminal of battery gets an excess of e- and becomes the cathode of electrolytic cell Electrode of cell connected to positive terminal of battery loses e- to battery and becomes the anode e- pump simultaneously supplying e- to cathode and recovering from anode This energy input from battery drives electrode reactions in electrolytic cell
Voltaic cell has copper cathode and zinc anode If battery connected to positive terminal contacts copper electrode and negative terminal contacts zinc electrode, e- flow in opposite directions
Battery forces cell to reverse its reaction Zn becomes cathode Cu becomes anode
Half-reaction at anode (Cu oxidized): Half-reaction at cathode (Zn reduced):
Important Differences Between Voltaic and Electrolytic Cells 1. The anode and cathode of an electrolytic cell are connected to a battery or other direct-current source, whereas a voltaic cell serves as a source of electrical energy. 2. Electrolytic cells have electrical energy from external sources creating nonspontaneous redox reactions. Voltaic cells have spontaneous redox reactions occurring. 3. Electrolytic cell have electrical energy converted to chemical energy. Voltaic cells have chemical energy converted to electrical energy
Electroplating Metals like Cu, Ag, Au are difficult to oxidize In electrolytic cell, these metals form ions at anode that are easily reduced at a cathode This allows solid metal from one electrode to be deposited on the other electrode Electroplating à an electrolytic process in which a metal ion is reduced and a solid metal is deposited on a surface
Electroplating cell has solution of a salt of plating metal Has object to be plated (cathode) Has piece of plating metal (anode)
Silver-plating cell has solution of soluble silver salt and silver anode Cathode is object to be plated Silver anode connected to positive electrode of battery or other source of direct current Object to be plated connected to negative electrode
Silver ions reduced at cathode according to following equation Silver ions deposited as metallic silver when electrons flow through circuit
Metallic silver is removed from anode as ions Silver atoms oxidized at anode as follows This maintains Ag+ concentration of solution
Rechargeable Cells Rechargeable cell combines redox chemistry of both voltaic and electrolytic cells When cell converts chemical energy to electrical energy it is a voltaic cell When recharged, it operates as an electrolytic cell converting electrical to chemical energy
The standard 12V car battery is set of 6 rechargeable cells Anode in each cell is lead submerged in solution of H 2 SO 4
Anode half-reaction is as follows
e- move through circuit to cathode where PbO 2 is reduced as follows
Next redox reaction for discharge cycle is as follows
A car s battery produces electric energy needed to start engine Sulfuric acid (present as ions) is consumed Lead (II) sulfate builds up as white powder on electrodes Once car is running, half-reactions are reversed by voltage produced by alternator Pb, PbO 2 and H 2 SO 4 are regenerated Battery can be recharged as long as all reactants are present and all reactions are reversible