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 ZnSO4

Cu strip is in solution of CuSO4

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 ZnSO4 (aq) is anode à where oxidation takes place

Cu strip in CuSO4 (aq) is cathode à where reduction takes place

The complete cell Copper half-cell written as Cu2+/Cu

Zinc half-cell written as Zn2+/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 CellsConsist of zinc container (anode)

Filled with moist paste of MnO2, graphite, and NH4Cl

When external circuit closed, Zn oxidized at negative electrode (anode)

e- move across circuit and reenter cell through carbon rod (cathode)

!MnO2 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) + 3O2(g) + xH2O(l) à 2Fe2O3·xH2O(s)

4Fe(s) + 3O2(g) + xH2O(l) à

Amount of hydration varies

Affects the color of rust formed

Electrochemical reactions:

Anode: Fe(s) à Fe2+(aq) + 2e-

Cathode: O2(g) + 2H2O(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 Fe2+ ions

Electrons released travel along metal to cathode

Fe2+ ions travel along moisture to cathode

At cathode, Fe2+ further oxidized to Fe3+

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 half-reaction 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 ZnSO4

2. Strip of Cu in CuSO4

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 Cu2+(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 E⁰cell 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 (E⁰anode) from standard reduction potential for rxn at cathode (E⁰cathode)

Example

Cells made from Fe in solution of Fe(NO3)3 and Ag in solution of AgNO3

Table 19-4 gives

Fe in Fe(NO3)3 is anode b/c it has lower reduction potential

Ag in AgNO3 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 E⁰cell 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

Electrolytic cells

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.

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 CellsRechargeable 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 H2SO4

Anode half-reaction is as follows

e- move through circuit to cathode where PbO2 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, PbO2 and H2SO4 are regenerated

Battery can be recharged as long as all reactants are present and all reactions are reversible