Seminar 6Electrochemistry. Modes of red-ox reaction balancing. Voltage series of metals. Theory of potential creation. Factors affecting the size of metal electrodes. Nernst's formula and its modifications. Normal electrode potential. Impact of environmental pH on the potential value. Types of electrodes and their examples (hydrogen and glass

electrode). The use of electrodes to create measuring cells. Construction and operation of a concentration cell. Electromotive force measurement. Examples of graphs, tasks and their solutions.

Tomasz Tuzimski

Medical University of Lublin, Faculty of Pharmacy with Medical Analytics Division, Chair of Chemistry,

Department of Physical Chemistry, 4A Chodźki Street, 20-093 Lublin, Poland

E-mail: tomasz.tuzimski@umlub.pl E-mail: tomasz.tuzimski@spoko.pl

Voltage series of metals (the galvanic series or electropotential series)

1. The most active metals can displace hydrogen from water.

Cs0 Rb0 K0 Ba0 Sr0 Ca0 Na0 Mg0 (T)

Me0 + H2O → MeOH + ½ H2↑

Me0 + 2H2O → Me(OH)2 + H2↑

K0 + H2O → KOH + ½ H2↑

Na0 + H2O → NaOH + ½ H2↑

Ba0 + 2H2O → Ba(OH)2 + H2↑

Ca0 + 2H2O → Ca(OH)2 + H2↑

Mg0 + 2H2O T→ Mg(OH)2 + H2↑

General reaction scheme

General reaction scheme

2. The metals lying in front of hydrogen are more active than hydrogen, thereforedisplace hydrogen from nonoxidizing acids (they do 'dissolve' in nonoxidizing acids).

General reaction scheme Me0 + HX → MeX + H2↑

Zn0 + H2SO4 (dilute, cold) → ZnSO4 + H2↑

General reaction scheme Me0 + 2HX → MeX2 + H2↑

Cd0 + 2HCl → CdCl2 + H2↑

Nonoxidizing Acids

HCl (aq), HBr (aq), HF (aq), H2S (aq)

H2SO4 (aq) – when diluteand cold!

H3PO4 (aq)

Most organic acids, e.g.,

CH3COOH

(acetic acid)

HCOOH

(formic acid)

In general, the reactionof an acid with a metal isa redox reaction in whichthe metal is oxidized and

the acid is reduced.

3. A more active metal will displace a less active one from its compounds (salts).

Zn0(s) + CuSO4(aq) → Cu0(s) + ZnSO4(aq)More „active” metal + Salt of less „active” metal

Cu0(s) + ZnSO4(aq)→ no reaction!Less „active” metal + Salt of more „active” metal

In other words,

the less „active” copper is unable

to displace the more „active” zinc from solution of its salt.

A more active metal will displace a less active one from its compounds (salts).

Cu0(s) + ZnSO4(aq)→ no reaction!Less „active” metal + Salt of more „active” metal

The less „active” copper is unableto displace the more „active” zinc

from solution of its salt.

The reaction of copper ions with zinc, viewed at the atomic level

Amphoteric metals (e.g., Zn0(s))Amphoteric metals react with strong acids and strongbasic:

Zn0 + 2HCl → ZnCl2 + H2↑

Zn0 + 2KOH → K2ZnO2 + H2↑

but DO NOT REACT WITH WATER:

Zn0 + H2O → no reaction!

Oxidizing Acids: HNO3 (dilute, conc.), H2SO4 (only conc.!)

Remember:

H2SO4 (aq) is a nonoxidizing acid when dilute and cold!

5. The metals behind hydrogen dissolve only in oxidizing acids(and the oxidizing mixture (co-called "royal water")).

LEAST ACTIVECu0 Ag0 Hg0 Au0

Oxidizing acids have „two potentials”, but the reaction with the biggest potential difference occurs only!

You can predict the reaction!

The winner is: greater potential difference!The stronger boxer wins!

Reducing agent / Oxidizing agentredox reaction

The reducing agent is the substance that isoxidized.The oxidizing agent is the substance that isreduced.A redox reaction is a chemical reaction inwhich changes in oxidation numbers occur.Oxidation is an increase in oxidation number.Reduction is a decrease in oxidation number.

Rules for Assigning Oxidation Numbers

1. The oxidation number of any free element (an elementnot combined chemically with a different element) is zero,regardless of how complex its molecules are.

2. The oxidation number for any simple, monatomic ion(e.g., Na+ or Cl-) is equal to the charge on the ion.

3. The sum of all the oxidation numbers of the atoms in amolecule or polyatomic ion must equal the charge on theparticle.

4. In its compounds, fluorine has an oxidation numer of -1.5. In its compounds, hydrogen has an oxidation number of

+1.6. In its compounds, oxygen has an oxidation numer of -2

(except H2O2, in which oxygen has an oxidation numberof -1).

Rules for Assigning Oxidation Numbers

Rules for Assigning Oxidation Numbers

Using Oxidation Numbers to Follow Redox Reactions

Ion-Electron Method of redox reaction balancing

The ion-electron method should be applied to balancing redox reactions.

In the ion-electron method are added: H+ and H2O, which help balance redox equation for Acidic Solution

OH- and H2O, which help balance redox equation for Basic SolutionH2O, which help balance redox equation for Neutral Solutionand electrons (to compensate the oxidation changes for all types ofsolutions).

These components of reactions (listed on the top) are added to cell's half-reactions for oxidation and

reduction (to balance the overall reaction).

Ion-Electron Method of redox reaction balancing(H+ and H2O help balance redox equation for Acidic Solution)

1. Divide the equation into two half-reactions.2. Balance atoms other than H and O.3. Balance H by adding H2O.4. Balance H by adding H+.5. Balance net charge by adding electron e-.6. Make e- gain e- loss, then add half-reactions.7. Cancel anything that’s the same on both

sides.

Balance redox equation for Acidic Solution (Step by Step)

Balance redoxequation for

Basic Solution (Additional Steps)

Nernst's formula and its modifications

𝐸 = 𝐸0 +𝑅𝑇

𝑛𝐹ln[𝑂𝑋𝐼𝐷𝐼𝑍𝐸𝐷 𝐹𝑂𝑅𝑀]

[𝑅𝐸𝐷𝑈𝐶𝐸𝐷 𝐹𝑂𝑅𝑀]

𝐸 = 𝐸0 +𝑅𝑇

𝑛𝐹ln

𝑎𝑀𝑒

𝑎𝑀𝑒0

a = const = 1

𝐸 = 𝐸0 +𝑅𝑇

𝑛𝐹ln 𝑎𝑀𝑒

𝐸 = 𝐸0 +0.05916

𝑛log 𝑎𝑀𝑒

𝐸 = 𝐸0 +0.05916

𝑛log[𝑀𝑒 ]

n+

n+

n+

n+

Impact of environmental pH on the potential value

MnO4-(aq) + 5e- + 8H+(aq) ↔ Mn2+(aq) + 4H2O

𝑶𝑿𝑰𝑫𝑰𝒁𝑬𝑫 𝑭𝑶𝑹𝑴

MnO4-(aq) + 3e- + 2H2O↔ MnO2(s) + 4OH-(aq)

𝑶𝑿𝑰𝑫𝑰𝒁𝑬𝑫 𝑭𝑶𝑹𝑴 𝑹𝑬𝑫𝑼𝑪𝑬𝑫 𝑭𝑶𝑹𝑴

(brown precipitate)

MnO4-(aq) + e- ↔ MnO4

2-(aq)𝑶𝑿𝑰𝑫𝑰𝒁𝑬𝑫 𝑭𝑶𝑹𝑴 𝑹𝑬𝑫𝑼𝑪𝑬𝑫 𝑭𝑶𝑹𝑴

(green solution)

Acidic 𝑹𝑬𝑫𝑼𝑪𝑬𝑫 𝑭𝑶𝑹𝑴(discoloration)

𝐸 = 𝐸0 +0.05916

𝑛lo𝑔

[𝑂𝑋𝐼𝐷𝐼𝑍𝐸𝐷 𝐹𝑂𝑅𝑀]

[𝑅𝐸𝐷𝑈𝐶𝐸𝐷 𝐹𝑂𝑅𝑀]

𝐸 = 𝐸0 +0.05916

5lo𝑔

MnO4−(aq) H+(aq) 8

Mn2+(aq) [H2O]4

Neutral

𝐸 = 𝐸0 +0.05916

3lo𝑔

MnO4−(aq) [H2O]2

MnO2(s) [OH−(aq)]4

Basic OH-

𝐸 = 𝐸0 +0.05916

1lo𝑔

MnO4−(aq)

MnO42−(aq)

Setting up a galvanic cellIf a shiny piece of metallic copper is placed into a solution of silver nitrate, aspontaneous reaction occurs. Gradually, a grayish white deposit forms on the copperand the solution itself becomes pale blue as hydrates Cu2+ ions enter the solution.

The equation is: Cu0(s) +2Ag+(aq) → Cu2+(aq) + 2Ag0(s)

Although the reaction is exothermic, no usable energy can be harnessed from itbecause all energy is dispersed as heat.

(a) A coil of copper wire stands next to a beaker containing a silver nitrate solution.(b) When the copper wire is placed in the solution, copper dissolves, giving the solution its blue color, and the

metallic silver deposit as glittering crystals on the wire.(c) After a while, much of the copper has dissolved and nearly all of the silver has deposited as the free metal.

Setting up a galvanic cellTo produce electrical energy, the two half-reactions involved in the net reaction must be madeto occur in separate containers or compartments called half-cells.An apparatus to accomplish this- a galvanic cell – is made up of two half-sells, as illustrated onFigure below:

On the left, a silver electrode dips into a solution of AgNO3. On the right, a copper electrode dips into a solution of Cu(NO3)2.

The two electrodes are connected by an external electricalcircuit and the two solutions are connected by a salt bridge.

In the silver half-cell, the followinghalf-reaction occurs:Ag+(aq) + e- → Ag0(s) (reduction)

In the copper half-cell, the followinghalf-reaction occurs:

Cu0(s) → Cu2+(aq) + 2e- (oxidation)

Cu0(s) + 2Ag+(aq) + 2e- → Cu2+(aq) + 2Ag0(s) + 2e- (cell reaction)

Standard cell notation

The anode half-cel is specified on the left, with the electrode material of the anode given first(in this case, the anode is copper metal).

The single vertical bar represents a phase boundary(here, between the copper electrode and the solution that surrounds it).

The double vertical bars represent the salt bridge, which connects the solutions in the two half-cells.A salt bridge is a tube filled with the solution of a salt composed of ions not involved in the cell

reaction (often KNO3 or KCl are used).

Changes that take place at the anode and cathode in the copper – silver galvanic cell

At the anode, Cu2+ ions enter the solution when copper atoms are oxidized, leaving electronsbehind on the electrode. Unless Cu2+ ions move away from the electrode or NO3

- ions movetoward it, the solution around the electrode will become positively charged.

At the cathode, Ag+ ions leave the solution and become silver atoms by acquiring electrons from the electrode surface. Unless more silver ions move toward the cathode or negative ions move

away, the solution around the electrode will becaome negatively charged.

Reduction of silverions at the cathodeextracts electrons

from the electrode, so the electrode

becomes positivelycharged.

Oxidation of copperatoms at the anode

leaves electronsbehind on the

electrode, whichbecomes negatively

charged.

The zinc-copper cell (Daniell cell)

Half-reactions for a Daniell cell are these:

Zinc electrode (anode):

Zn0(s) → Zn2+ (aq) + 2 e–

Copper electrode (cathode): Cu2+ (aq) + 2 e– → Cu0(s)

Cell potentials

A galvanic cel has an ability topush electrons the externalcircuit. The magnitude of thisability is expressed as a potential.Potential is expressed in anelectrical unit called the volt (V),which is a measure of the amountof energy, in joules, that can bedelivered per SI unit of chargé(the coulomb) as the currentmoves through the circuit. Thus, acurrent flowing under a potentialof 1 volt can deliver 1 joule ofenergy per coulomb.

1V = 1J/C

The potential generated by galvanic cell has also been calledan electromotive force (emf), suggesting that it can be thoughtof as the force with which anelectric current is pushed througha wire. Electrical current and emf areoften likened to the flow of waterin a house („current”) and the pressure of the water („emf”).

Cell potential (Ecell) and Standard cell potential (Eocell)

The maximum potential that a given cell can generate is calledits cell potential (Ecell). The cell potential (Ecell) depends on:- the composition of the electrodes,- the concentrations of ions in the half-cells,- the temperature.

The standard cell potential (Eocell) – this is potential of the cell

when all of the ion concentrations are 1.00 M, the temperatureis 25°C, and any gases that are involved in the cel reactions areat a pressure of 1 atm.

Standard electrode (reduction or cell) potential (Eocell)

Standard electrode potentials are usually tabulated as reduction potentials. However, the reactions are reversible and the role of a particularelectrode in a cell depends on the relative oxidation/reductionpotential of both electrodes. The oxidation potential for a particularelectrode is just the negative of the reduction potential. A standardcell potential can be determined by looking up the standardelectrode potentials for both electrodes (sometimes called half-cellpotentials).The one that is smaller will be the anode and will undergo oxidation.The cell potential is then calculated as the sum of the reductionpotential for the cathode and the oxidation potential for the anode.

E°cell = E°red(cathode) – E°red(anode) = E°red(cathode) + E°oxi(anode)

Examples of standard electrode (reduction

or cell) potentials(Eo

cell) at 25°C

Types of electrodes and their examples hydrogen electrode (reference electrode)

0V a ≈ 1+]

[H+]

Types of electrodes and their examples (glass electrode)

(a) Photograph of a typical glass electrode.(b) Cutaway view of the construction of a glass

electrode and a reference electrode.

=

A glass electrode is a type of ionselective electrode that is used todetermine the concentration ofhydrogen ions (pH measurement).

The glass electrode works in the pHrange from 2 to 10, in which itspotential is a linear function.

Construction and operation of a concentration cell

A concentration cell is an electrochemical cell where the twoelectrodes are the same material, the electrolytes on the two half-cells involve the same ions, but the electrolyte concentration differsbetween the two half-cells:

Me0(s) I Men+ (concentration 1) II Men+ (concentration 2) I Me0(s)

concentration 1 < concentration 2

Where the half-cell reactions for oxidation and reduction are:

Oxidation: Me0(s) → Men+ (concentration 1) + n e–

Reduction: Men+ (concentration 2) + n e– → Me0(s)Overall reaction: Men+ (concentration 2) → Men+ (concentration 1)

The cell's emf is calculated through Nernst equation as follows:

𝐸 = 𝐸0 −0.05916

𝑛log

[𝑀𝑒 𝑛 + (𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 1)]

[𝑀𝑒 𝑛 + (𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 2)]

Electromotive force measurement

In general, the electromotive force can be calculated:

ΔE = E↗higher potential – E↘lower potential

ΔE = Ecathode – Eanode

ΔE = Eoxidizing agent – Ereducing agent

Homework

In the near future...

Please check in the next days...

Sources(books to be used during all seminars)

• James E. Brady, Fred Sense, Chemistry Matter and Its Changes, John Wiley & Sons, Fourth Edition, 2004, ISBN 0-471-44891-5 (WileyInternational Edition).

• J. Brady, N. Jespersen, A. Hyslop – Chemistry, 7 ed., International student version. Wiley, 2015

• D. A. Skoog, F. J. Holler, S. R. Crouch - Principles of instrumentalanalysis. - 6th ed. , Thomson, Brooks/Cole, 2007

• J. Crowe. T. Bradshaw, P. Monk, Chemistry for the Biosciences. The essential concepts., Oxford University Press, 2006.

• D. A. Skoog, D. M. West, F. J. Holler, S. R. Crouch; Analytical Chemistry. An Introduction, 7th ed., 2000

• J. A. Beran; Laboratory Manual for Principles of General Chemistry 7. P. Monk, Physical Chemistry , Wiley 2004.

• Wikipedia – free encyclopedia.