Tro, Chemistry: A Molecular Approach 1 - starab.kau.edu.sastarab.kau.edu.sa/GetFile.aspx?id=164354&Lng=AR&fn=Chapter 19_LE… · Tro, Chemistry: A Molecular Approach 21 Electrochemical

Tro, Chemistry: A Molecular Approach 1

Electrochemistry

19.1 Redox Reactions.

19.2 Galvanic Cells.

19.3 Standard Reaction Potential.

19.4 Spontaneity of Redox Reactions.

19.5 The Effect of Concentration on Cell Emf.

19.6 Batteries.

19.7 Corrosion.

19.8 Electrolysis.

3

Terms used with electricity:

•Galvanometer - an instrument for detecting electric current.

•Voltmeter - an instrument to measure the potential difference between two

half-cells in a voltaic cell.

•Amperes - the SI units for current flow = # e- / s

•Volts - a measurement of electric potential difference between two

electrodes.

•Electromotive force, emf - another term for volts.

•Cations - positive ions attracted to the cathode.

•Anions - negative ions attracted to the anode.

•Electrolysis - the process by which an electric current produces a chemical

change.

•Voltaic cell - a device used to produce electric energy from an oxidation-

reduction reaction.

•Battery - two or more electrochemical cells operating as a unit.

•External circuit - a "wires" connected to a battery providing a path for

electricity to flow.

•Internal circuit - the electrolyte inside a battery through which ions can

move.


19.1 Redox Reactions

Electrochemistry is the branch of chemistry that deals

with the interconversion of electrical energy and

chemical energy.

Electrochemical processes are redox (oxidation -

reduction) reactions.

In redox reactions, electrons are transferred from one

substance to another.



Redox Reaction• one or more elements change oxidation number

all single displacement, and combustion,

some synthesis and decomposition

• always have both oxidation and reduction

split reaction into oxidation half-reaction and a

reduction half-reaction

• aka electron transfer reactions

half-reactions include electrons

• oxidizing agent is reactant molecule that causes oxidation

contains element reduced

• reducing agent is reactant molecule that causes reduction

contains the element oxidized


Oxidation & Reduction

• oxidation is the process that occurs whenelement loses electrons

compound adds oxygen

oxidation number of an element increases

compound loses hydrogen

half-reaction has electrons as products

• reduction is the process that occurs whenoxidation number of an element decreases

element gains electrons

compound loses oxygen

compound gains hydrogen

half-reactions have electrons as reactants




Rules for Assigning Oxidation States

5. in their compounds, nonmetals have oxidation

states according to the table below

nonmetals higher on the table take priority

Nonmetal Oxidation State Example

F -1 CF4

H +1 CH4

O -2 CO2

Group 7A -1 CCl4

Group 6A -2 CS2

Group 5A -3 NH3

Balancing Redox Equations

Suppose we are asked to balance the equation showing

the oxidation of Fe(III) ions by dichromate ions () in

acidic medium.

The following steps will help us balance the equation





Practice - Balance the Equation

ClO3-1 + Cl-1 Cl2 (in acid)


Practice - Balance the Equation

ClO3-1 + Cl-1 Cl2 (in acid)

+5 -2 -1 0

oxidationreduction

ox: 2 Cl-1 Cl2 + 2 e-1 } x5

red: 2 ClO3-1 + 10 e-1 + 12 H+ Cl2 + 6 H2O} x1

tot 10 Cl-1 + 2 ClO3-1 + 12 H+ 6 Cl2 + 6 H2O

1 ClO3-1 + 5 Cl-1 + 6 H+1 3 Cl2 + 3 H2O

Worked Example 19.1


Electrochemical Cells• 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 voltaic cell

aka galvanic cells

• nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy


Electrochemical Cells• oxidation and reduction reactions kept separate

half-cells

• electron flow through a wire along with ion flow

through a solution constitutes an electric circuit

• requires a conductive solid (metal or graphite)

electrode to allow the transfer of electrons

through external circuit

• ion exchange between the two halves of the system

electrolyte


Electrodes• Anodeelectrode where oxidation occursanions attracted to itconnected to positive end of battery in electrolytic

cellloses weight in electrolytic cell

• Cathodeelectrode where reduction occurscations attracted to itconnected to negative end of battery in electrolytic

cellgains weight in electrolytic cellelectrode where plating takes place in electroplating

19.2 Galvanic Cell

The experimental apparatus for generating

electricity through the use of a spontaneous

reaction is called a galvanic cell or voltaic cell,

Figure 19.1 shows the essential components of a

galvanic cell.



The difference in electrical potential between the anode

and the cathode is measured by a voltameter (Figure 19.2)

The voltage across the electrodes of a galvanic cell is

called the cell voltage, or cell potential .

Another common term for cell voltage is the

electromotive force or emf (E), which is a measure of

voltage, not force.



Current and Voltage• the number of electrons that flow through the system per

second is the current unit = Ampere

1 A of current = 1 Coulomb of charge flowing by each second

1 A = 6.242 x 1018 electrons/second

Electrode surface area dictates the number of electrons that can flow

• the difference in potential energy between the reactants and products is the potential difference unit = Volt

1 V of force = 1 J of energy/Coulomb of charge

the voltage needed to drive electrons through the external circuit

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 case 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


Standard Reduction Potential• a half-reaction with a strong tendency to

occur has a large + half-cell potential

• when two half-cells are connected, the electrons will flow so that the half-reaction with the stronger tendency will occur

• we cannot measure the absolute tendency of a half-reaction, we can only measure it relative to another half-reaction

• we select as a standard half-reaction the reduction of H+ to H2 under standard conditions, which we assign a potential difference = 0 v standard hydrogen electrode, SHE


19.3 Standard Reduction Potentials

19.3

Standard reduction potential (E0) is the voltage associated with

a reduction reaction at an electrode when all solutes are 1 M and

all gases are at 1 atm.

E0 = 0 V

Standard hydrogen electrode (SHE)

Reduction Reaction

2e- + 2H+ (1 M) H2 (1 atm)


Half-Cell Potentials• SHE reduction potential is defined to be exactly 0 v

• half-reactions with a stronger tendency toward

reduction than the SHE have a + value for E°red

• half-reactions with a stronger tendency toward

oxidation than the SHE have a value for E°red

• E°cell = E°oxidation + E°reduction

E°oxidation = E°reduction

when adding E° values for the half-cells, do not multiply the

half-cell E° values, even if you need to multiply the half-

reactions to balance the equation





Worked Example 19.2

Worked Example 19.3

19.4 Spontaneity of Redox Reaction

How E⁰ for cell is related to thermodynamic

quantities such as ΔG⁰ and K


E°cell, DG° and K

• for a spontaneous reaction

one the proceeds in the forward direction with the

chemicals in their standard states

DG° < 1 (negative)

E° > 1 (positive)

K > 1

• DG° = −RTlnK = −nFE°cell

n is the number of electrons

F = Faraday’s Constant = 96,485 C/mol e−




Worked Example 19.4

Worked Example 19.5

19.5 The Effect of Concentration on Cell Emf

There is a mathematical relationship between

the emf of a galvanic cell and the

concentration of reactants and products in a

redox reaction under nonstandard-state

conditions. This equation called

Nernest Equation

E = E⁰ - RT /nF ln [oxd]/[red]

From equation (18.13) page 821

ΔG = ΔG⁰ + RT ln Q

because : ΔG = -nFE and ΔG⁰ = -nFE⁰



During the operation of a galvanic cell, electrons flow

from the anode to the cathode, resulting in product

formation and a decrease in reactant concentration.

Thus Q increases, which means that E decreases, the cell

reaches equilibrium.

At equilibrium, there is no net transfer of electrons,

so E= 0 and Q = K, where K is the equilibrium constant



Worked Example 19.6

Worked Example 19.7


Concentration Cells• it is possible to get a spontaneous reaction when the oxidation

and reduction reactions are the same, as long as the electrolyte concentrations are different

• the difference in energy is due to the entropic difference in the solutions

the more concentrated solution has lower entropy than the less concentrated

• electrons will flow from the electrode in the less concentrated solution to the electrode in the more concentrated solution

oxidation of the electrode in the less concentrated solution will increase the ion concentration in the solution – the less concentrated solution has the anode

reduction of the solution ions at the electrode in the more concentrated solution reduces the ion concentration – the more concentrated solution has the cathode

Concentration Cell

Because electrode potential depends on ion concentration, it is possible to construct a

galvanic cell from two half-cells composed of

the same material but differing in ion

concentrations. Such a cell is called

a concentration cell.


Concentration Cell

when the cell concentrations

are different, electrons flow

from the side with the less

concentrated solution

(anode) to the side with the

more concentrated solution

(cathode)

Cu(s) Cu2+(aq) (0.010 M) Cu2+

(aq) (2.0 M) Cu(s)

19.6 Batteries

A battery is a galvanic cell, or a series of combined

galvanic cells, that can be used as a source of direct

electric current at a constant voltage.

The operation of battery is similar in principle to that

of the galvanic cells.






19.7 Corrosion

Corrosion is the term usually applied to the deterioration

of metals by an electrochemical process. There are many

examples of corrosion: Rust on iron, tarnish of silver, and the

green patina formed on copper and brass.

Corrosion causes damage to buildings, bridges, ships, and cars.

The most familiar example of corrosion is the formation of rust

on iron.

Oxygen gas and water must be present for iron to rust.

Corrosion

The main steps are believed to be as follows:

A region of the metal’s surface serves as the anode,

The oxedation reaction is:

Fe(s) Fe²+(aq) + 2e

The electron given up by iron reduce atmospheric oxygen

to water at the cathode, which is another region of the

same metal surface:

O2 + 4H + 4e 2H2O(l)

The overall reaction is:

2Fe(s) + O2(g) + 4H(aq) 2Fe²+(aq) + 2H2O(l)

We find the standard emf for this process:

Eºcell = Eºcathode – Eºanode

=1.23 V – (-0.44 V) = 1.67 V

Note that this reaction occurs in an acidic

medium, see figure 19.14 which shows the

mechanism of the rust formation.



Cathodic protection is a process in which the metal that is

to be protected from corrosion is made the cathodic in

what amounts to a galvanic cell.

Figure 19.15 shows how an iron nail can be protected

from rusting by connecting the nail to a piece of zinc.

With out such protection, an iron nail quickly rusts in

water.

Rusting of underground iron pipes and iron storage tanks

can be prevented or greatly reduced by connecting them

to metals such as zinc and magnesium, which oxidize

more readily than iron as in figure 19.16


Figure 19.15


19.8 Electrolysis

In contrast to spontaneou redox reactions, which result

in the conversion of chemical energy into electrical

energy, electrolysis is the process in which electrical

energy is used to cause a nonspontaneous chemical

reaction to occure.

An electrolytic cell is an apparatus for carrying out

electrolysis.

Electrolysis of molten Sodium Chloride




Faraday’s Law

• the amount of metal deposited during

electrolysis is directly proportional to the charge

on the cation, the current, and the length of time

the cell runs

charge that flows through the cell = current x time

Quantitative Aspects of Elctrolysis

The quantitative treatment of electrolysis was developed primarily

by Faraday.

He observed that:

The mass of product formed (or reactant consumed) at an

electrode is proportional to both the amount of electricity

transferred at the electrode and the molar mass of the substance.

The relationship between charge (in coulombs, C) and current is:

1C = 1A x 1s

That is a coulomb is the quantity of electrical charge passing any

point in the circuit in 1 second when the current is 1 ampere.


Figure 19.20 shows the steps involved in calculating the quantities

of substances produced in electrolysis.

Finish





Terms used with electricity:

•Galvanometer - an instrument for detecting electric current.

•Voltmeter - an instrument to measure the potential difference between two half-

cells in a voltaic cell.

•Amperes - the SI units for current flow = # e- / s

•Volts - a measurement of electric potential difference between two electrodes.

•Electromotive force, emf - another term for volts.

•Cations - positive ions attracted to the cathode.

•Anions - negative ions attracted to the anode.

•Electrolysis - the process by which an electric current produces a chemical change.

•Voltaic cell - a device used to produce electric energy from an oxidation-reduction

reaction.

•Battery - two or more electrochemical cells operating as a unit.

•External circuit - a "wires" connected to a battery providing a path for electricity to

flow.

•Internal circuit - the electrolyte inside a battery through which ions can move.

Documents

Tro, Chemistry: A Molecular Approach 1 - starab.kau.edu.sastarab.kau.edu.sa/GetFile.aspx?id=164354&Lng=AR&fn=Chapter 19_LE… · Tro, Chemistry: A Molecular Approach 21 Electrochemical