REDOX CLASSIFICATION OF NATURAL WATERS

Oxic waters

- waters that contain measurable dissolved oxygen.

Suboxic waters

- waters that lack measurable oxygen or sulfide, but do contain significant dissolved iron (> ~0.1 mg L -1 ).

Anoxic waters

- waters that contain both dissolved iron and sulfide.

DEFINITION OF Eh

Eh - the potential of a solution relative to the SHE.

Both pe and Eh measure essentially the same thing. They may be converted via the relationship: 

pe



Eh

2 .

303

RT

Where  = 96.42 kJ volt -1 eq -1 (Faraday’s constant).

At 25°C, this becomes

pe

 16 .

9

Eh

or

Eh

 0 .

059

pe

Eh – Measurement and meaning

• Eh is the driving force for a redox reaction • No exposed live wires in natural systems (usually…)  where does Eh come from?

• From Nernst  redox couples exist at some Eh (Fe 2+ /Fe 3+ =1, Eh = +0.77V) • When two redox species (like Fe 2+ and O 2 ) come together, they should react towards equilibrium • Total Eh of a solution is measure of that equilibrium

FIELD APPARATUS FOR Eh MEASUREMENTS

PROBLEMS WITH Eh MEASUREMENTS

• Natural waters contain many redox couples NOT at equilibrium; it is not always clear to which couple (if any) the Eh electrode is responding.

• Eh values calculated from redox couples often do not correlate with each other or directly measured Eh values.

• Eh can change during sampling and measurement if caution is not exercised.

• Electrode material (Pt usually used, others also used) – Many species are not

electroactive

(do NOT react at electrode) • Many species of O, N, C, As, Se, and S are not electroactive at Pt – electrode can become poisoned by sulfide, etc.

Figure 5-6 from Kehew (2001). Plot of Eh values computed from the Nernst equation vs. field-measured Eh values.

Other methods of determining the redox state of natural systems

• For some, we can directly measure the redox couple (such as Fe 2+ and Fe 3+ ) • Techniques to directly measure redox SPECIES: – Amperometry (ion specific electrodes) – Voltammetry – Chromatography – Spectrophotometry/ colorimetry – EPR, NMR – Synchrotron based XANES, EXAFS, etc.

Free Energy and Electropotential

• Talked about electropotential (aka emf, Eh)  driving force for e transfer • How does this relate to driving force for any reaction defined by D G r ??

– D

G r = n

D

E or

D

G 0 r = n

D

E 0 Where n is the # of e ’s in the rxn,



is Faraday’s constant (23.06 cal V -1 ), and E is electropotential (V)

• pe for an electron transfer between a redox couple analagous to pK between conjugate acid base pair

Electromotive Series

• When we put two redox species together, they will react towards equilibrium, i.e., e- will move  which ones move electrons from others better is the electromotive series • Measurement of this is through the electropotential for half-reactions of any redox couple (like Fe 2+ and Fe 3+ ) – Because D G r = n D E, combining two half reactions in a certain way will yield either a + or – electropotential (additive, remember to switch sign when reversing a rxn) -E  D G r , therefore  spontaneous • In order of decreasing strength as a reducing agent  strong reducing agents are better e- donors

Biology’s view  upside down?

Reaction directions for 2 different redox couples brought together??

More negative potential



reductant // More positive potential



Example – O 2 /H 2 O vs. Fe 3+ /Fe 2+



O 2 oxidizes Fe 2+ is spontaneous!

oxidant

Nernst Equation

Consider the half reaction: NO 3 + 10H + + 8e  NH 4 + + 3H 2 O(l) We can calculate the Eh if the activities of H + , NO 3 , and NH 4 + are known. The general Nernst equation is

Eh



E

0  2 .

303

RT n

 log

Q

The Nernst equation for this reaction at 25°C is

Eh



E

0  0 .

0592 8 log  

a NH

4 

a NO

3 

a

10

H

  

Let’s assume that the concentrations of NO 3 and NH 4 + 3  10 -7 have been measured to be 10 -5 M and M, respectively, and pH = 5. What are the Eh and pe of this water?

First, we must make use of the relationship

E

0   D

G r n



o

For the reaction of interest D r G° = 3(-237.1) + (-79.4) - (-110.8)

E

0  = -679.9 kJ mol -1 679 .

9  0 .

88 volts ( 8 )( 96 .

42 )

The Nernst equation now becomes

Eh

 0 .

88  0 .

0592 8 log  

a NH

 4

a NO

3 

a

10

H

   substituting the known concentrations (neglecting activity coefficients)

Eh

 0 .

88  0 .

0592 log 8  3 5  10  7

  

10  0 .

521 volts and

pe

 16 .

9

Eh

 16 .

9 ( 0 .

521 )  8 .

81

Stability Limits of Water

• H 2 O  2 H + + ½ O 2(g) + 2e Using the Nernst Equation: -

Eh



E

0  0 .

0592 log

n

1

p

1 2

O

2

a H

2  • Must assign 1 value to plot in x-y space (P O2 ) • Then define a line in pH – Eh space

UPPER STABILITY LIMIT OF WATER (Eh-pH)

To determine the upper limit on an Eh-pH diagram, we start with the same reaction 1/2O 2 (g) + 2e + 2H +  H 2 O but now we employ the Nernst eq.

Eh



E

0  0 .

0592 log

n

1

p

1 2

O

2

a H

2 

Eh



E

0  0 .

0592 log 2 1

p

1 2

O

2

a H

2 

E

0   D

G r

0

n

   (  237 .

1 ) ( 2 )( 96 .

42 )  1 .

23 volts

Eh



Eh

 1 .

23  0 .

0296 log 1

p O

2 2

a H

2  1 .

23  0 .

0148 log

p O

2  0 .

0592

pH

As for the pe-pH diagram, we assume that p O2 = 1 atm. This results in

Eh

 1 .

23  0 .

0592

pH

This yields a line with slope of -0.0592.

LOWER STABILITY LIMIT OF WATER (Eh-pH)

Starting with H + + e  1/2H 2 (g) we write the Nernst equation

Eh



E

0  0 .

0592 log

p H

1 2 2 We set p H2 0. Thus, we have 1

a



H

= 1 atm. Also, D G r ° = 0, so E 0 =

Eh

  0 .

0592

pH

C 2 HO

Redox titrations

• Imagine an oxic water being reduced to become an anoxic water • We can change the Eh of a solution by adding reductant or oxidant just like we can change pH by adding an acid or base • Just as pK determined which conjugate acid-base pair would buffer pH, pe determines what redox pair will buffer Eh (and thus be reduced/oxidized themselves)

Making stability diagrams

• For any reaction we wish to consider, we can write a mass action equation for that reaction • We make 2-axis diagrams to represent how several reactions change with respect to 2 variables (the axes) • Common examples: Eh-pH, P O2 -pH, T-[x], [x]-[y], [x]/[y]-[z], etc

Construction of these diagrams

• For selected reactions: Fe 2+ + 2 H 2 O  FeOOH + e + 3 H +

Eh



E

0  0 .

0592 1 log    3

a H



a Fe

2     How would we describe this reaction on a 2-D diagram? What would we need to define or assume?

• How about: • Fe 3+ + 2 H 2 O  FeOOH (ferrihydrite) + 3 H + K sp =[H + ] 3 /[Fe 3+ ] log K=3 pH – log[Fe 3+ ] How would one put this on an Eh-pH diagram, could it go into any other type of diagram (what other factors affect this equilibrium description???)

Redox titrations

• Imagine an oxic water being reduced to become an anoxic water • We can change the Eh of a solution by adding reductant or oxidant just like we can change pH by adding an acid or base • Just as pK determined which conjugate acid-base pair would buffer pH, pe determines what redox pair will buffer Eh (and thus be reduced/oxidized themselves)

Redox titration II

• Let’s modify a bjerrum plot to reflect pe changes 100 90 80 70 60 50 -4 H 2 S(aq) -2 0 2 4 pe 6 8 - SO 4 10 12 Greg Mon Oct 25 2004

Redox titrations in complex solutions

• For redox couples not directly related, there is a ladder of changing activity • Couple with highest + potential reduced first, oxidized last • Thermodynamics drives this!!

Oxic Post - oxic Sulfidic Methanic O 2

The Redox ladder

Aerobes H 2 O NO 3 Dinitrofiers N 2 MnO 2 Mn 2+ Maganese reducers Fe(OH) 3 Fe 2+ Iron reducers SO 4 2 H 2 S Sulfate reducers CO 2 Methanogens CH 4 H 2 O H 2 The redox-couples are shown on each stair-step, where the most energy is gained at the top step and the least at the bottom step. (Gibb’s free energy becomes more positive going down the steps)