Aquatic Redox Chemistry

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Aquatic Redox Chemistry

Paper No: 16 Environmental Chemistry Module: 23 Aquatic Redox Chemistry

Development Team

Principal Investigator

&

Co- Principal Investigator

Prof. R.K. Kohli

Prof. V.K. Garg & Prof. Ashok Dhawan Central University of Punjab, Bathinda

Paper Coordinator

Prof. K.S. Gupta

University of Rajasthan, Jaipur

Content Writer

Dr. Alka Sharma

University of Rajasthan, Jaipur Content Reviewer Prof. K.S. Gupta

University of Rajasthan, Jaipur

Anchor Institute Central University of Punjab

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Description of Module

Subject Name Environmental Sciences Paper Name Environmental Chemistry Module Name/Title Aquatic Redox Chemistry

Module Id EVS/EC-XVI/23

Pre-requisites A basic knowledge of redox chemistry

Objectives

1. To define redox (e- transfer) reactions, half reactions 2. To define and understand aquatic redox chemistry 3. To define redox potential, cell reactions and cell potential 4. To define standard electrode potential E⁰ and Nernst equation 5. To understand pE-scale, significance and measurement of pE values 6. To understand redox ladder

7. To understand redox potential in natural water, its interpretation and significance 8. To define and understand applications of redox potentials in aquatic systems Keywords Redox reactions, redox process, redox potential, pE-scale, redox ladder, aquatic redox

system

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Contents 1. Introduction 2. Half Reactions 3. Reduction Potential

4. Standard Reduction Potential 5. Interpretation and Significance 6. Cell Reaction and Cell Potential

7. Hydrogen Electrode and determination of Electrode Potential 8. pE-Scale

9. Significance of pE Values 10. Measurement of pE-Values 11. Solved Problem

12. pE and pH Relationship 13. Pourbaix diagram

14. Interpretation and Significance 15. Solved Problem

16. The Limits of pE and pH in Natural Waters 17. Redox Potentials in Natural Systems 18. Redox Ladder

19. Effect of redox on Metal Pollution

20. Suggesting Reading

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Aquatic Redox Chemistry Introduction

In aquatic environmental systems, including soils, sediments, aquifers, rivers, lakes, and water treatment systems, the most important and interesting chemical reactions occurring are the oxidation-reduction (redox) reactions. These reactions are central to major element cycling, to many sorption processes, to trace element mobility and toxicity, to most remediation schemes, and to life itself.

Why study Redox Reactions?

They fuel and constrain more or less all life processes.

They are a major determinant of chemical species present in natural environments.

Redox reactions are core to many emerging domains of the aquatic sciences research, which includes all aspects of the aquatic sciences: involving the hydrosphere, aquatic (i.e., aqueous) aspects of environmental processes in the atmosphere, lithosphere, biosphere, etc. The Aquatic Redox Chemistry has multidisciplinary roots (straddling mineralogy to microbiology) and interdisciplinary applications (e.g., in removal of contaminants from sediment, soil or water).

To characterize the oxidation-reduction status of surface environments, the geochemists, soil scientists and limnologists have used redox potential (Ered) measurements. The redox potential of aquatic, marine and soil systems is a measure of electrochemical potential or electron availability within these systems. The redox potential (Ered) is determined from the concentration of oxidants and reductants in the environment. Oxygen, nitrate, nitrite, manganese, iron, sulphate, and CO2 are some of the prominent inorganic oxidants; while organic substrates and reduced inorganic compounds are the well-known reductants.

The redox-potential is the evaluation of the equilibrium potential, i.e. reduction/oxidation potential, built at the interface between an electrode (a noble metal) and the solution consisting of electroactive redox species and is measured under standard state conditions (at 25 ⁰C, 1 atmospheric pressure and one unit activity for all species) with respect to the standard hydrogen electrode. The term ‘redox’ is the occurrence of both the chemical changes: oxidation and reduction, in a chemical reaction. The redox reactions i.e. oxidation-reduction reactions

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entail the changes of oxidation states of reactants in a reaction and it is the sum up of two half reactions. The two processes: oxidation (loss of electrons) and reduction (gain of electrons) takes place together during the same reaction, i.e. oxidation/reduction never takes place in isolation; hence, these reactions are called oxidation-reduction reactions or redox reactions. The potential of the overall reaction at the standard state can be depicted as E⁰ = E⁰ ox + E⁰red (where E⁰ ox and E⁰red are the potentials of oxidation half-reaction and reduction half-reaction respectively)

For example: In a reaction:

Copper (II) oxide + Hydrogen ^heat Copper + Water

Oxidation

CuO + H2 ^ Cu + H2O Reduction

Consider another reaction:

Zinc oxide + carbon Zinc + Carbon monoxide

Oxidation ZnO + C Zn + CO Reduction

In both these reactions, the hydrogen and carbon removes oxygen from copper(II) oxide and zinc oxide respectively, hence hydrogen and carbon are reducing agents or reductants. Similarly, those which are oxygen providers are called oxidizing agents or oxidants.

The modern electronic concept of oxidant and reductant is: one which accepts electrons is oxidant and one which donates electron is reductant.

For example: Cl2 + 2 I^ 2 Cl^ + I2

Here, Cl2 is an oxidant and I^ is a reductant; as I^ is oxidized to I2 by Cl2 and Cl2 is reduced to Cl^ by I^).

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Reducing agent (Compound A) Oxidizing agent (Compound B)

A oxidized (Electrons loss) B reduced (Electrons gain)

Compound A (Oxidized Form) Compound B (Reduced Form)

Thus, in a redox reaction which is brought about by loss and gain of electrons simultaneously, the oxidant is reduced and the reductant oxidized with an exchange of n electron (e^) which may be depicted as:

Oxidant + n e^ reductant

Some examples of Reductant in wetland soil are:

1. Organic matter & other organic compounds;

2. Reduced inorganic compounds, viz, Mn²⁺, S^2, CH4, H2, Fe²⁺, NH4 etc.

Oxidants are inorganic compounds, viz, O2, NO^3, FeOOH, SO42, HCO^3 etc.

It may be noted that in the overall redox reaction no free electrons are generated. The movement of electrons is from reductant to oxidant, thereby an electrical potential is developed between the two which is measured in volts and denoted by E°. Since the redox potential depends on the members of the pair in a reaction, hence it’s a relative rather than an absolute value. The standard redox potential values for elements, compounds and ions can be ascertained under standard conditions (unit molar concentration at 1 atm pressure and 25 ⁰C), relative to a standard hydrogen electrode (SHE) potential (which is arbitrarily given a potential, E° = zero volts). The greater the positive potential, the more anticipated it will be reduced.

The redox-potentials (also known as electrochemical version of Gibbs free energy G) are used to ascertain the direction and the free energy (the change in the system’s free energy G) of redox-reaction at standard states as:

G⁰ (eV) = nFE°

A

A

B

B

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where F is the Faraday constant number (96,485 C/mol or 96,485 J/mol/V or ≈100 kJ/mol/V) and n is the number of electrons involved in the redox-reaction). (The ⁰ symbol is for the substances involved in the reaction in their standard states).

For non-standard redox-reactions, the difference in redox-potential (precisely reduction potential E) is correlated to G as: ΔG (eV) = -nFE

Half-Reactions

A half reaction is either the reduction or the oxidation reaction component of a redox-reaction. The reactions occurring in an electrochemical cell is frequently described on the basis of the half-reaction concept.

A redox reaction is expressed as the difference of two reduction half-reactions.

For example:

Cu²⁺(aq) + 2e^ Cu(s) (reduction of Cu²⁺) (E⁰ = + 0.34 V) Zn²⁺(aq) + 2e^ Zn(s) (reduction of Zn²⁺) (E⁰ = - 0.76 V) Deducing the two half-reactions:

Cu²⁺(aq) + Zn(s) Cu(s) + Zn²⁺(aq)

Or, in other words, combining the two half-reactions, reduction of Cu²⁺ and oxidation of Zn(s) would result into a net redox-reaction:

Cu²⁺(aq) + 2e^ Cu(s) and Zn(s) Zn²⁺ + 2e^

Being half components of the redox-reaction, these reactions are termed as Half-reactions.

Similarly, the redox reaction: 2Na + Cl2 Na⁺ + 2Cl^ is made up of two half-reactions:

2Na 2Na⁺ + 2e^ (oxidation) Cl2 + 2e 2Cl (reduction)

Another example of half-reaction is the reduction of MnO4 in which atom transfer accompanies electron transfer; here the oxygen atoms are lost from MnO4 (aq) and H2O(l):

MnO4(aq) + 8H⁺(aq) + 5e^ Mn²⁺(aq) + 4H2O(l)

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Similarly, adding the simultaneously occurring two half reactions of metallic zinc atoms reaction with aqueous nickel ions give the net redox-reaction:

Zn(s) Zn²⁺(aq) + 2e^ 2 e^ + Ni²⁺(aq) Ni(s)

--- Ni²⁺(aq) + Zn(s) Ni(s) + Zn²⁺(aq)

The half-reactions depict the exact oxidation state changes happening in two half-reactions.

Reduction Potential Definition

By definition, redox-potential (reduction potential/electrode potential) is the tendency of a substance to accept electrons. The positive (high) values of electrode potential means that the elements or ions would readily accept electrons; on the other hand, negative (low) values indicate their easy capability of electron donation.

Few examples are cited below:

O2 + 4H ⁺ + 4e^ 2H2O E° = +1.23 (V) Fe³⁺ + e^ Fe²⁺ E° = +0.77 (V)

2H⁺ + 2e^ H2O E° = 0 (V) (reference) Na⁺ + e^ Na E° = 2.71 (V)

The standard reduction potential E^ored is measured relative to: 2H⁺(aq) + 2e^ H2(g) (which has an assigned value E^ored= 0.00 V and hence treated as reference)

Measurement of redox potential, E, allows quantitative appraisal of the force and tendency of the system. It can be measured by the difference between the potential of the hydrogen electrode (or more easily, the calomel electrode) and the potential of a platinum electrode immersed in the medium.

Interpretation and Significance

Redox (electron-transfer) reactions provide the energetic basis for the life process, and through this, play a decisive role in the geochemical cycle of the elements. Redox reactions are very significant in water-saturated

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environments, such as sediments, soils, and sludges. Actually, all aquatic organisms obtain their energy for metabolic processes from oxidation-reduction reactions. Photosynthetic organisms catalytically reduce CO2 to reduced organic matter by locking light energy, while non-photosynthetic organisms catalytically decompose the organic products of photosynthesis through energy-yielding redox reactions. One of the major elements in the terrestrial and aquatic environments – nitrogen – circulates by many microbially catalyzed redox reactions.

In truth, the only non-redox process in the entire nitrogen cycle is NH3 integration with and liberation from N- containing organic matter. The movements of many other elements also involve redox reactions, such as C, Fe, and S. The oxidizing power of anaerobic environments in the biosphere is mainly controlled by five molecules.

In decreasing order of energy produced, they are nitrate (NO3–), manganese dioxide (MnO2), ferric hydroxide (Fe(OH)3), sulfate (SO42–) and, under extreme conditions, carbohydrate (CH2O) itself.

Like Gibb’s free energy (G), the redox potentials (E) are not absolute. The redox potentials are measured with reference to the reduction of hydrogen ions to hydrogen gas at standard state conditions (SHE), i.e., 25 degrees Celsius (°C), 1 atmospheric pressure, and one unit activity for all species. The potential of this reaction, by convention, is taken as zero. Furthermore, as the redox potentials are comprised of two parts: the oxidation potential, and the reduction potential, the potential of the overall reaction of a cell at the standard state can be expressed as: E⁰cell= E⁰_ox + E⁰_red.

A positive E^ored indicate that a half-reaction will proceed in the direction indicated (reduction) when paired with the hydrogen half-reaction where as a negative E^ored means that a half-reaction will move in the opposite direction indicated (oxidation).

The numerical values of E°ox and E°red of a substance in a half-reaction are the same with opposite in sign (until the conditions are changed). Thus, numerically, the potential for oxidation half-reaction is the negative of the potential of the reduction half-reaction, i.e. E°ox =  E°red (with opposite sign). The voltage produces by an electrochemical cell is determined by summing up all the potentials in circuit, as:

E^ocell = E^ohalf-reactions.

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Table 1: Electrode (reduction) potentials of selected half-reactions at 25 ⁰C.

Half-reaction E⁰ (V)

---

F2 + 2e^  2F^ +2.87

MnO4— + 8H⁺ + 5e^  Mn²⁺ + 4H2O +1.49

Cl2 + 2e^  2Cl^ +1.36

Cr2O72 + 14H⁺ + 6e^  2Cr³⁺ + 7H2O +1.33 NO3— + 6H⁺ + 5e^  ½ N2+ 3H2O +1.26

O2 + H⁺ + 4e^  2H2O +1.23

Br2 + 2e^  2Br^ +1.09

NO3— + 4H⁺ + 3e^  NO+ 2H2O +0.96 NO3— + 2H⁺ + 4e^  NO2— + H2O +0.85

Ag⁺ + e^  Ag +0.80

Hg22+ + 2e^  2Hg +0.79

Fe³⁺ + e^  Fe²⁺ +0.77

I2 + 2e^  2I^— +0.54

Cu²⁺ + 2e^  Cu +0.34

SO4 + 10H⁺ + 8e^  H2S + 4H2O +0.31 CO2 + 4H⁺ + 4e^  C + 2H2O +0.21 SO42 + 4H⁺ + 2e^  H2SO3 + H2O +0.20

N2 + 6H⁺ + 6e^  2NH3 +0.09

2H⁺ + 2e^  H2 0.00

Fe²⁺ + 2e^  Fe 0.44

Zn²⁺ + 2e^  Zn 0.76

Al³⁺ + 3e^  Al 1.66

Mg²⁺ + 2e^  Mg 2.37

Na⁺ + e^  Na 2.71

Ca²⁺ + 2e^  Ca 2.87

K⁺ + e^  K 2.92

Li⁺ + e^  Li 3.05

S t r o n g

R e d u c i n g

a g e n t S

t r o n g

o x i d i z i n g

a

g

e

n

t

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The value of redox potential (E) (for non-standard states) under concentration conditions can be related to standard potential (E⁰) by the Nernst equation:

E = E⁰+ RT/nF * ln {(Ox)/(Red)} = E⁰ RT/nF * ln {(Red)/(Ox)} = E⁰ RT/nF * ln Q Or at 298 K, this is expressed as: E = E⁰ 0.0592/n * log {(Red)/(Ox)}

where, E⁰ = standard electrode potential, R = ideal gas constant (8.314 J/mol-K); T = absolute temperature (Kelvin); n = number of electrons involved in the reaction; F = the Faraday constant number and (Red)/(Ox) is the concentration ration of reduced and oxidized forms of a given reaction pair, i.e. the reaction quotient.

The important factors upon which the redox potential of a substance depends are: nature of the substance, its affinity for electron, concentration of reductants and oxidants (referred as redox pair) and temperature.

--- Solved Problem (based on Nernst equation):

For the half reaction 2H⁺ + 2e^ = H2. What is E1/2 at pH 5 and P_H2 = 1 atm?

Solution: pH =  log [H⁺] = 5, therefore [H⁺] = 10^-5 M.

In the above half-reaction, n = 2.

Putting these values in the Nernst equation: E = E⁰  0.0592/2 * log P_H2/(H⁺)² E = E⁰  0.0592/2 * log (10⁵)²

Since, for the given half-reaction, the E⁰1/2 = 0 V, hence, E = 0.000 0.0592/2 * (10) Therefore, E =  0.296 V.

---

Many electrode combinations are possible in electrochemical cells, and it is convenient to specify a standard potential, E^o, for each electrode by referencing it to the SHE (whose standard potential is defined as zero). There are many half-reactions whose electrode potential cannot be measured, because the electron transfer reaction at an electrode is too slow.

These potentials
can
nevertheless
be
calculated
from
the
free
energy
of
appropriate
redox
 reactions.

For
example, the
formation
of
NO
from
N2
and
O2 is
a
redox
reaction:

O2
+
N2
 2NO which
can
be
divided
into
the
half-reactions:

O2 + 4e– + 4H+ 2H2O and 2NO + 4e– + 4H+ N2 + 2H2O

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From
the
free
energy
of
the
overall
reaction 173.4 kJ, we obtain a cell potential of -0.45 V (using equation

G = –nFE). Then knowing that the standard potential of the oxygen electrode is 1.24 V, we can readily evaluate the standard potential for half reaction as 1.69 V, (1.24
V
–[‐0.45
V]),
even
though it is impossible to measure this potential directly because the electron transfer between the electrode and NO and N2 molecules is too slow to establish a reversible potential.

Significance of Redox Electrodes, Redox Potentials & Redox Reactions in the determination of:

i. mobility and toxicity of chemical species in the environment (natural aquatic systems vary widely in redox conditions; therefore, fate may constantly change);

ii. equilibrium constant: single electrode potentials may be used to determine equilibrium constants of ionic reactions;

iii. solubility product of a sparingly soluble salt;

iv. pH, using hydrogen electrode

Cell Reaction and Cell Potential

Let us see how the redox-reactions are a source of electric current in the electrochemical cells (a device for producing current from a chemical (redox) reaction).

For example, in the redox-reaction: Cu²⁺(aq) + Zn(s) Cu(s) + Zn²⁺(aq); the electrons released from a half-reaction, i.e. oxidation of Zn(s) are consumed by the other half-reaction, i.e. reduction of Cu²⁺(aq).

Since both the reactions occur on the zinc electrode itself (dipped in CuSO4 solution), there is no net charge.

Now, if the two half-reactions occur in two separate compartments (one with Zn rod dipped in ZnSO4 solution and other with Cu rod dipped in CuSO4 solution) and both well connected by a wire, there will be a net flow of electrons from the reductant in one compartment to the oxidant in another through wire. However, the flow of current will be instant and later it will stop due to charge built up in the two compartments. The current flow can be restarted just by connecting the two compartments by a Salt bridge (a U-tube filled with an electrolyte such as NaCl, KCl, K2SO4 etc.) thus providing a passage to ions from one compartment to other without extensive mixing of the two solutions. This will complete the circuit and the electrons pass freely through the wire to maintain the net charge zero in the two compartments. This kind of circuitry cell is a simple Voltaic (Galvanic) cell where electrical current is generated by a spontaneous redox reaction.

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In an electrochemical cell, the flow of electrons from one electrode to the other is due to the half-reactions occurring in the anodic and cathodic compartments and the addition of these two half-reactions gives the net chemical change called as cell reaction. Thus, the E° (electrode potential) for a given substance can be determines by constructing an electrochemical cell consisting of two half-cells. If the flow of electrons is from oxidant (hydrogen half-cell) to the other half-cell, the substance has a positive redox potential; whereas the E°

negative indicate the substance is reductant. Knowing the redox potentials values of two substances will help to predict whether a redox reaction between them is theoretically possible.

In electrochemical cells, or in redox reactions that happen in solution, the thermodynamic driving force can be measured as the cell potential. Chemical reactions are spontaneous in the direction of -ΔG, which is also the direction in which the cell potential (defined as Eanode - Ecathode) is positive. A cell operating in the spontaneous direction (for example, a battery that is discharging) is called a galvanic cell. A cell that is being driven in the non-spontaneous direction is called an electrolytic cell.

Hydrogen Electrode and Determination of Electrode Potential

2 H+ (aq) + 2e^  H2 (g)

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The Standard hydrogen electrode (SHE) is a redox electrode which forms the basis of the thermodynamic scale of oxidation-reduction potentials. The absolute electrode potential of SHE is estimated to be 4.44 ± 0.02 V at 25 °C, but for constructing a base for comparison with all other electrode reactions, hydrogen’s standard electrode potential, E⁰, is declared to be zero at all temperature. The absolute electrode potential, according to IUPAC, is the electrode potential of a metal measured with respect to a universal reference system (without any additional metal-solution interface).

The standard electrode potential, E⁰, is measured under standard conditions: 25 ⁰C, 1 M concentration for each ion participating in the reaction, a partial pressure of 1 atm for each part of the reaction, and metals in their pure state. The standard reduction potential is defined relative to a SHE reference electrode (arbitrary with a potential 0.00 V). Since, E⁰ (SHE) = 0, E⁰ for other half reaction can be >0 [oxidizes H2 (g) or <0 [is oxidized by H2 (g)].

The conventions for Eº:

a. Eº values (units of volts) are compared on the basis of half reactions, which by convention are written as reductions;

b. all substances are assumed to be at unit activity;

c. all Eº values are determined relative to the reduction potential of the standard hydrogen electrode (SHE).

1. If the Eº for a given half-reaction is >0, that couple has the potential (under standard conditions) to oxidize the SHE;

2. A negative Eº indicates a couple that can reduce the SHE (at standard conditions).

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The most important factor affecting redox reactions of some species is pH, and the effects are very well illustrated by Eh-pH (or pe-pH, or pE-pH, or Pourbaix) diagrams. Since it is frequently difficult to determine which half-reactions are actually coupled in nature, the concept of pE is preferred.

pE may be defined as the activity of the free electron in water, thus only half reactions can be focused upon.

Forever, the reactions are expressed as the reduction half reaction. From the equation: G⁰ = –nFE, a negative

G⁰ corresponds to a positive E⁰, and thus to a potentially spontaneous reduction half-reaction (at standard conditions) versus the SHE.

For graphical expressions, pE is negative log of electron activity, i.e., pE = -log(e^) = F E⁰ / 2.3 RT = 1/n log K

= -1/n G⁰ / 2.3 RT.

Or, the pE equation can be expressed as: pE = pE⁰ - log (Red) / (Ox); where pE⁰ can be calculated from log K. If the equilibrium distribution of (Red)/(Ox) is known, this equation can be used to solve for the pE of the

environment.

For example, for the half-reaction Cu²⁺ + 2e^ Cu (s);

Since n = two electrons involved and K= 1/({Cu²⁺} {e^}²), therefore, pE will be expressed as:

pE = 1/2 log K + 1/2 Log {Cu²⁺}

[as pE⁰ = 1/n log K and pE = pE⁰ + (1/n) log {ox}/{red}]

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pE⁰ is readily obtainable from thermodynamic data, now multiplying both the sides of above equation by 2.303 RT/F, equation becomes:

pE = F/(2.303RT) EH

pE⁰ = 16.90 E⁰H

pE⁰ = -1/n (1.753x10^-4)(G⁰)

G⁰ = G⁰f (products) - G⁰f (reactants)

Significance of pE Values

At equilibrium, solution has one pE value; if this is assumed, the ratio of oxidized to reduced species activities can be calculates; moreover, determination of prominent species is possible. Usually, electrons move from high to low activity (Ae); small pE corresponds to reducing environments, reduced species predominates; and large pE correspond to oxidizing environments, oxidized species predominates. The pE diagrams can be an asset to know how species predominate as a function of pE.

Measurement of pE Values

The pE⁰ of a half reaction expresses the electron activity required to maintain reactants and products at unit activities. There are two main types of redox calculations. The first is the calculation of what controls the pE of the environment. This is analogous to calculating the pH of the environment.

The second type of calculation is to determine how trace species respond or distribute themselves with respect to that pE. Again by analogy, when we know the pH, we can calculate the pH dependent speciation of trace species.

--- Solved Problem:

For the reduction: Fe³⁺ + 2e^ Fe²⁺ K = 10¹³

The electron activity of the solution must be held to the very low value of 10^-13 (pE = 13) to maintain equal activities of the two ions; this corresponds to usually referred as an oxidizing environment.

The value of K = 10¹³ actually refers, of course, to the reaction; Fe³⁺ + H2 Fe²⁺ + 2H⁺ K = {Fe²⁺} {H⁺} /{Fe³⁺} {PH5

2}

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So that the equilibrium condition {Fe³⁺} = {Fe²⁺} would require a pH of 13 at this unit pressure of H2 or a hydrogen partial pressure of 10²⁶ atm at zero pH.

---

pE and pH Relationship:

Pourbaix Diagram

A Pourbaix diagram (also known as Potential/pH diagram, EH-pH diagram or a pE/pH diagram) is essentially an electrochemical phase diagram, the best representation of the possible thermodynamically stable phases of an aqueous electrochemical system (i.e. redox-active substances). The diagrams are invented by the Russian born, Belgium Chemist Marcel Pourbaix (1904-1998) and are named after him.

Figure 2: Pourbaix diagram for iron at ionic concentrations of 1.0 mM

The diagrams have two axes, the vertical axis labeled as pE or EH represent the voltage potential with respect to SHE (as EH (SHE)/V; where H stands for hydrogen); and the horizontal axis measures pH. The lines depict the conditions under which two phases coexist in equilibrium and also the redox and acid-base reactions. The

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shaded area represents the conditions of potential and pH for the stable phase of the substance, whereas outside the shaded region represents the thermodynamically unstable phase which either gets reduced or oxidized.

Eh-pH diagram showing the predominance fields for oxidized (upper right) and reduced (lower left) forms of selected redox-active species. Dashed diagonal lines are for the H2/H2O (lower) and H2O/O2 (upper) couples and together they enclose the conditions over which water is stable.

Interpretation and Significance

A Pourbaix diagram designates mainly three regions: ‘immunity’, ‘corrosion’ and ‘passivity’ instead of stable species. These regions describe the stability of a particular substance in a specific environment. ‘Immunity”

indicate the safe and non-attacked region which is opposite to the ‘corrosion’ region; ‘passivity’ other hand depict the relative stability (i.e. when a metal forms a stable oxide or other salt coating on its surface).

For example, in the Pourbaix diagram for Fe (above):

Areas in the Pourbaix diagram mark regions where a single species (Fe²⁺(aq), Fe3O4(s), etc.) is stable. More stable species tend to occupy larger areas.

Lines mark places where two species exist in equilibrium.

 Pure redox reactions are horizontal lines - these reactions are not pH-dependent

 Pure acid-base reactions are vertical lines - these do not depend on potential

 Reactions that are both acid-base and redox have a slope of -0.0592 V/pH x4H⁺⁄4e^-)

Illustration of equilibria in the iron Pourbaix diagram (numbered on the plot):

1. Fe²⁺ + 2e^- → Fe (s) (pure redox reaction - no pH dependence) 2. Fe³⁺ + e^- → Fe²⁺ (pure redox reaction - no pH dependence) 3. 2 Fe³⁺ + 3 H2O → Fe2O3(s)+ 6H⁺ (pure acid-base, no redox)

4. 2 Fe²⁺ + 3 H2O → Fe2O3(s)+ 6H⁺ + 2e^- (slope = -59.2 x 6/2 = -178 mV/pH) 5. 2 Fe3O4(s) + H2O → 3 Fe2O3(s) + 2H⁺ + 2e^- (slope = -59.2 x 2/2 = -59.2 mV/pH)

For an element, such as iron, the water redox lines have special significance on a Pourbaix diagram. In other case, such as liquid water, it is stable only in the region between the dotted lines. Below the H2 line, water is unstable relative to hydrogen gas, and above the O2 line, water is unstable with respect to oxygen; while, for

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active metals such as Fe, the region where the pure element is stable is typically below the H2 line. This means that iron metal is unstable in contact with water, undergoing reactions:

Fe(s) + 2H⁺ → Fe²⁺(aq) + H2 (in acid) Fe(s) + 2 H2O → Fe(OH)2(s) + H2 (in base)

Iron (and most other metals) are also thermodynamically unstable in air-saturated water, where the potential of the solution is close to the O2 line in the Pourbaix diagram. Here the spontaneous reactions are:

4 Fe(s) + 3 O2 + 12H⁺ → 4 Fe³⁺ + 6 H2O (in acid) 4 Fe(s) + 3 O2 → 2 Fe2O3(s) (in base)

--- Solved Problem: Consider the spontaneous reaction of Zn metal in a solution of Cu²⁺

Zn(s) + Cu²⁺(aq) Zn²⁺(aq) + Cu(s)

Without doing the experiment, could we have predicted that this reaction would be spontaneous as written?

Solution: We can calculate G^o, or we can calculate a related quantity, E^o. First write the half-reactions as reductions:

Cu²⁺(aq) + 2e^- Cu(s) E^ored = +0.340 V Zn²⁺(aq) + 2e^- Zn(s) E^ored = -0.763 V

E^ored are the standard reduction potentials for these two half-reactions. The half-reaction with the more positive E^ored will occur as written (as a reduction).

So for the Cu²⁺ and Zn²⁺ half-reactions, the Cu²⁺ E^o (+0.340 V) is more positive than the Zn²⁺ E^o (-0.763 V), so the Cu²⁺ will be reduced, and Zn metal will be oxidized.

---

The Limits of pE and pH in Natural Waters:

The limits of stability of water itself can be an index of pE and pH of natural waters. The stability of water depends on the availability of e^s so that it is neither reduced nor oxidized by these half reactions:

Reducing limit: 4H2O + 4e^ → 2H2 + 4OH^

Oxidizing limit: 2H2O → O2 + 4H⁺ + 4e^ E⁰ = 1.23V, pE⁰ = 20.75 Only the oxidation state of H changes in this reaction, and hence is equivalent to:

4H⁺ + 4e^ → 2H2 E⁰ = 0V, pE⁰ = 0

The oxidizing limit for the stability of water is described as a function of E⁰ (or pE) and pH:

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E=E⁰ + 0.059/n log(PO2 [H⁺]⁴/[H2O]). Putting n=4 and PO2=1 atm, E= E⁰  0.059pH = -0.059pH + 1.23V or in terms of pE, the above equation is expressed as: pE= - pH + 20.75

Only the oxidation state of H changes in this reaction, and hence is equivalent to:

4H⁺ + 4e^ → 2H2 E⁰ = 0V, pE⁰ = 0

The oxidizing limit for the water-stability is described as a function of E⁰ (or pE) and pH: E=E⁰ + 0.059/n log(PO2 [H⁺]⁴/[H2O]). By putting n=4 and PO2=1 atm, E= E⁰  0.059pH = -0.059pH + 1.23V or in terms of pE, the above equation is expressed as: pE= - pH + 20.75

For present atmosphere, using PO2=0.2 atm, the expression becomes:

E=  0.059pH + 1.22V (indicating minor shift)

Similarly, reducing limit for water-stability is predicted (H2=1atm): E=E⁰+0.059/n log([H⁺]⁴/PH22)

By putting n=4 & PH2=1atm, E=E⁰+0.059/4*4log([H⁺]⁴/PH22) =E⁰0.059pH =0.059pH+0V or in terms of pE, the above equation is expressed as: pE= - pH

pE/ pH and E

/ pH diagrams (Plot between E

vs pH or pE vs pH):

The parallelogram depicts the Water-Stability on Earth’s surface in terms of EH and pH.

Line-a – depicting Reducing limit and Line-b – depicting Oxidizing limit. Below line-a (zone-I) is too reducing for water; whereas above line-b (zone-III) is too oxidizing for water. Zone falling in between Reducing &

Oxidizing limits, i.e. between the two parallel lines a & b (zone-II) is just right for water.

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Similarly, Pourbaix diagram for other ions in the water stability field can be drawn.

The creation of phase diagram for any solute ion, or combination of solute ions will illustrate the speciation (form) of that/those ion(s) in EH vs pH or pE vs pH space for a given solute concentration in solution. Though creating phase diagram for mixed ions are quite complex in comparison to single ion diagrams.

Applications:

Monitoring of industrial waste water Swimming pool water monitoring Soil investigation

Freshwater habitat quality

Redox potentials in natural systems:

At almost neutral pH (7-8), the redox potentials in natural waters range from about –400 mV to +800 mV. They are bounded in the negative range by the reduction of H₂O to hydrogen gas (H_2(g)) and in the positive range by the oxidation of H₂O to O_2(g).

In water and sediment, four representative ranges of redox potentials exist:

i. Range I: For oxygen-bearing waters: water saturated with oxygen may have a redox potential within the first range (710 to 800 mV at pH 7 to 8).

ii. Range II: in systems where some oxygen has been consumed, the redox potentials may range between – 100 to 710 mV (at pH 7 to 8). It is representative of many ground and soil waters where O₂ has been consumed (by degradation of organic matter), but SO₄

2

is not yet reduced. In this range soluble Fe (II) and Mn (II) are present; their concentration is redox-buffered because of the presence of solid Fe (III) and Mn (III, IV) oxides.

iii. Range III: though the potential is not sensitive to oxygen concentrations and has been observed to remain nearly constant down to values of 0.1% O₂ saturation. It is characterized by SO₄

2

/ HS or SO₄

2

/ FeS₂ redox equilibria.

iv. Range IV: for anaerobic sediments and sludges: Within the second range, solid Fe (III), and Mn (III, IV) are reduced to soluble Fe (II) and Mn (II) when organic matter is mineralized. Phosphorus, which is

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initially bound to Fe (III) hydroxides, is released to the overlying water as a result of the reduction of Fe (III) solid phases. Under even more reduced conditions (Range 3, -240 to –100 mV at pH 7 to 8), the concentration of Fe (II) and Mn (II) rises further; SO₄

2

begins to be reduced, accompanied by the precipitation of Manganous Sulfide (MnS) and Ferrous Sulfide (FeS), and the formation of pyrite.

Many researchers have studied the relationships between oxidation-reduction potential and the physical, chemical, and biological processes in soil-water systems. One of the pioneers in this field was Mortimer (1941, 1942). Mortimer studied the factors which control the rate of nutrient supply to phytoplankton in systems of lake water and sediment deposits.

Figure 3: Redox intensity representative ranges in soil and water. (pE = 16.9 E

o

).

Redox Ladder:

The change in the pE of a fresh natural water in contact with sediment as a function of amount of organic matter decomposed or in other words, the Redox Ladder may be defined as: the step by step sequencing of common redox. The redox reactions occurring in an aquatic environment, a step wise pE sketch is formed in which at a particular place or time, pE is fixed until a particular oxidant is consumed.

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As illustrated in fig., at almost 7 pH, the change in the pE of fresh natural water in contact with sediment as a function of amount of organic matter decomposed. The amount of the organic matter reacted is depicted horizontally (thus its length depends on the availability of specific solid phases for reaction).

Effect of redox on metal pollution:

Changes in the redox potential can have important consequences for environmental pollution, especially with respect to metal ions such as cadmium, lead, and nickel. In general, the solubility of heavy metals is highest in oxidizing and acidic environments (Figure: Eh/pH as a function of different aquatic environments). At neutral to alkaline pHs in oxidizing environments, these metals often adsorb onto the surface of insoluble Fe(OH)3 and MnO2 particles, especially when phosphate is present to act as a bridging ion. When the redox potential shifts to only slightly oxidizing or slightly reducing conditions as a result of microbial action, and the pH shifts toward the acidic range, Fe(OH)3 and MnO2 in soils and sediments are reduced and solubilized. The adsorbed metal ions likewise become solubilized and move into groundwater (or into the water column of lakes when there is Fe(OH)3 or MnO2 in the sediment). Conversely, if sulfate is reduced microbially to HS^– metal ions are immobilized as insoluble sulfides. But if sulfide rich sediments are exposed to air through drainage or dredging operations, then HS^– is oxidized back to sulfate, and the heavy metal ions are released.

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Figure 4: EH/pH as a function of different aquatic environments. Oval enclosed by dashed line indicates region of highest solubility of heavy metals. [Source: Adapted from W. Salomons (1995). Long‐term strategies for handling contaminated sites and large‐scale areas. In Biogeodynamics of Pollutants in Soils and Sediments, W Salomons and W.M. Stigliani, eds. (Berlin: Springer‐Verlag)].

A particularly important instance of biological redox mediation of heavy-metal pollution occurs in the case of mercury. Inorganic mercury, in any of its common valence states, Hg⁰, Hg22+, and Hg²⁺, is not toxic when ingested; it tends to pass through the digestive system, although Hg⁰ is highly toxic when inhaled. But the methylmercury ion (CH3)Hg⁺ is very toxic, regardless of the route of exposure. The environmental route to toxicity involves sulfate reducing bacteria that live in anaerobic sediments. As part of their metabolism these bacteria use methyl groups to produce acetate. When exposed to Hg²⁺ the bacteria transfer the methyl groups to the mercury, producing (CH3)Hg⁺; because methylmercury is soluble, it enters the aquatic food chain, where it is bioaccumulated in the protein-laden tissue of fish.

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1. Stanley E. Manahan, Environmental Chemistry, 9^th Edition, CRC Press, New York, 2009.

2. Colin Baird and Michael Cann, Environmental Chemistry, 4^th Edition, WH Freeman, New York, 2008.

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