Redox Geochemistry

Oxidation – Reduction Reactions Oxidation - a process involving loss of electrons. Reduction - a process involving gain of electrons. Reductant - a species that loses electrons. Oxidant - a species that gains electrons. Free electrons do not exist in solution. Any electron lost from one species in solution must be immediately gained by another. Ox 1 + Red 2  Red 1 + Ox 2 LEO says GER

Fundamental electromagnetic relations: Electric charge (q) is measured in coulombs (C). –The magnitude of the charge of a single electron is x C. 1 mole of electrons has a charge of x 10 4 C which is called the Faraday constant (F) –q=n*F The quantity of charge flowing each second through a circuit is called the current (i). The unit of current is the ampere (A) 1 A = 1 C/sec The difference in electric potential (E) between two points is a measure of the work that is needed when an electric charge moves from one point to another. Potential difference is measured in volts (V) 1 V = 1 J/C –The greater the potential difference between two points, the stronger will be the "push" on a charged particle traveling between those points. A 12 V battery will push electrons through a circuit 8 times harder than a 1.5 V battery. Ohm’s Law: V = I R  potential is equal to current * resistance

Half Reactions Often split redox reactions in two: –oxidation half rxn  e- leaves left, goes right Fe 2+  Fe 3+ + e- –Reduction half rxn  e- leaves left, goes right O e -  2 H 2 O SUM of the half reactions yields the total redox reaction 4 Fe 2+  4 Fe e- O e -  2 H 2 O 4 Fe 2+ + O 2  4 Fe H 2 O

Examples Balance these and write the half reactions: Mn (IV) + H 2 S  Mn 2+ + S 0 + H+ CH 2 O + O 2  CO 2 + H 2 O H 2 S + O 2  S 8 + H 2 O

Redox Couples For any half reaction, the oxidized/reduced pair is the redox couple: –Fe 2+  Fe 3+ + e- –Couple: Fe 2+ /Fe 3+ –H 2 S + 4 H 2 O  SO H e- –Couple: H 2 S/SO 4 2-

Half-reaction vocabulary part II Anodic Reaction – an oxidation reaction Cathodic Reaction – a reduction reaction Relates the direction of the half reaction: A  A + + e - == anodic B + e -  B - == cathodic

ELECTRON ACTIVITY Although no free electrons exist in solution, it is useful to define a quantity called the electron activity: The pe indicates the tendency of a solution to donate or accept a proton. If pe is low, there is a strong tendency for the solution to donate protons - the solution is reducing. If pe is high, there is a strong tendency for the solution to accept protons - the solution is oxidizing.

THE pe OF A HALF REACTION - I Consider the half reaction MnO 2 (s) + 4H + + 2e -  Mn H 2 O(l) The equilibrium constant is Solving for the electron activity

WE NEED A REFERENCE POINT! Values of pe are meaningless without a point of reference with which to compare. Such a point is provided by the following reaction: ½H 2 (g)  H + + e - By convention so K = 1.

THE STANDARD HYDROGEN ELECTRODE If a cell were set up in the laboratory based on the half reaction ½H 2 (g)  H + + e - and the conditions a H + = 1 (pH = 0) and p H 2 = 1, it would be called the standard hydrogen electrode (SHE). If conditions are constant in the SHE, no reaction occurs, but if we connect it to another cell containing a different solution, electrons may flow and a reaction may occur.

STANDARD HYDROGEN ELECTRODE ½H 2 (g)  H + + e -

ELECTROCHEMICAL CELL ½H 2 (g)  H + + e - Fe 3+ + e -  Fe 2+

We can calculate the pe of the cell on the right with respect to SHE using: If the activities of both iron species are equal, pe = If a Fe 2+ /a Fe 3+ = 0.05, then The electrochemical cell shown gives us a method of measuring the redox potential of an unknown solution vs. SHE. ELECTROCHEMICAL CELL

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: Where  = kJ volt -1 eq -1 (Faraday’s constant). At 25°C, this becomes or

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  G r ??  G r = - n  E –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

Nernst Equation Consider the half reaction: NO H + + 8e -  NH H 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 The Nernst equation for this reaction at 25°C is

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

CALIBRATION OF ELECTRODES The indicator electrode is usually platinum. In practice, the SHE is not a convenient field reference electrode. More convenient reference electrodes include saturated calomel (SCE - mercury in mercurous chloride solution) or silver-silver chloride electrodes. A standard solution is employed to calibrate the electrode. Zobell’s solution - solution of potassium ferric-ferro cyanide of known Eh.

CONVERTING ELECTRODE READING TO Eh Once a stable potential has been obtained, the reading can be converted to Eh using the equation Eh sys = E obs + Eh Zobell - Eh Zobell-observed Eh sys = the Eh of the water sample. E obs = the measured potential of the water sample relative to the reference electrode. Eh Zobell = the theoretical Eh of the Zobell solution Eh Zobell = (t - 25) Eh Zobell-observed = the measured potential of the Zobell solution relative to the reference electrode.

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 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.

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 ). Reducing waters (anoxic) - waters that contain both dissolved iron and sulfide.

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

The Redox ladder H2OH2O H2H2 O2O2 H2OH2O NO 3 - N2N2 MnO 2 Mn 2+ Fe(OH) 3 Fe 2+ SO 4 2- H2SH2S CO 2 CH 4 Oxic Post - oxic Sulfidic Methanic 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)