Redox Geochemistry Oxidation Reduction Reactions R E 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 G s y a s Free electrons do not exist in solution Any electron lost from one species in solution must be immediately gained by another Ox1 Red2 Red1 Ox2 L O E Fundamental electromagnetic relations Electric charge q is measured in coulombs C The magnitude of the charge of a single electron is 1 602 x 10 19 C 1 mole of electrons has a charge of 9 649 x 104 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 Fe2 Fe3 e Reduction half rxn e leaves left goes right O2 4 e 2 H 2O SUM of the half reactions yields the total redox reaction 4 Fe2 4 Fe3 4 eO2 4 e 2 H2O 4 Fe2 O2 4 Fe3 2 H2O Examples Balance these and write the half reactions Mn IV H2S Mn2 S0 H CH2O O2 CO2 H2O H2S O2 S8 H2O Redox Couples For any half reaction the oxidized reduced pair is the redox couple Fe2 Fe3 e Couple Fe2 Fe3 H2S 4 H2O SO42 10 H 8 e Couple H2S SO42 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 pe log ae 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 MnO2 s 4H 2e Mn2 2H2O l The equilibrium constant is K a Mn 2 a H4 ae2 Solving for the electron activity aMn 2 ae 4 Ka H 1 2 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 H2 g H eBy convention G o f H G so K 1 K o f H2 G a H ae 1 pH22 o f e 1 0 THE STANDARD HYDROGEN ELECTRODE If a cell were set up in the laboratory based on the half reaction H2 g H eand the conditions a H 1 pH 0 and p H2 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 H 2 1 a tm P la tin u m e le c tr o d e H2 g H e a H 1 ELECTROCHEMICAL CELL H 2 1 a tm V P la tin u m e le c t r o d e P la t in u m e le c tr o d e S a lt B r id g e a H Fe 1 H2 g H e 3 Fe 2 Fe3 e Fe2 ELECTROCHEMICAL CELL We can calculate the pe of the cell on the right with respect to SHE using a Fe2 pe log a 3 Fe 12 8 If the activities of both iron species are equal pe 12 8 If a Fe2 a Fe3 0 05 then pe log 0 05 12 8 14 1 The electrochemical cell shown gives us a method of measuring the redox potential of an unknown solution vs SHE 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 303RT 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 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 Gr Gr 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 acidbase pair Nernst Equation Consider the half reaction NO3 10H 8e NH4 3H2O l We can calculate the Eh if the activities of H NO3 and NH4 are known The general Nernst equation 2 303RT is 0 Eh E log Q n The Nernst equation for this reaction at 25 C is a NH 0 0592 4 Eh E 0 log 10 8 a NO a H 3 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 Fe2 Fe3 1 Eh 0 77V When two redox species like Fe2 and O2 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 Ehsys Eobs EhZobell EhZobell observed Ehsys the Eh of the water sample Eobs the measured potential of the water sample relative to the reference electrode EhZobell the theoretical Eh of the Zobell solution EhZobell 0 428 0 0022 t 25 EhZobell 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 …