Bio-electrochemistry course

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1 Bio-electrochemistry course LabMET short course on microbial energy management and microbial fuel cells Peter Aelterman and Korneel Rabaey

2 Course overview Bacterial metabolism and the redox potential Basic electrochemistry Electron transfer losses Reactor based losses Analytical techniques Alleviating losses Bacterial electrodynamics

3 I. The bacterial metabolism and the redox potential

4 Bacterial metabolic pathways Metabolism = transfer of electrons from low to high potential, with charge balance. Aim: to generate ATP / NADH Oxidative versus fermentative patterns Oxidative: bacteria transfer electrons towards an electron acceptor Fermentative: bacteria transfer electrons towards substrate derived products, and generate hydrogen

5 Oxidative metabolism

6 Fermentative metabolism Electrons released from substrate are partially used to reduce substrate products Formation of acetate, butyrate, ethanol Hydrogen formation to regenerate NADH

7 Fermentation

8 Electron donor/acceptor: What is the potential? E (mv) : equilibrium potential of a compound (1/1 ox/red) at ph 7. Gives an indication of the intrinsic capacity of a compound to function as electron donor/acceptor Nernst: E = E R T n F [ C] [ A] Hence: potential dependent on many factors such as ph, temperature, concentrations, metabolic phase An electron acceptor for bacterium A can be an electron donor for bacterium B x c a [ D] [ B] d b

9 Redox reaction Redox potentials E ' 0 (mv) 2 H e - H Ferredoxin(Fe 3+ ) + e - Ferredoxin(Fe 2+ ) -420 NAD + + H + + 2e - NADH -320 S + 2 H + + 2e - H 2 S SO H e - H 2 S + 4 H 2 O -220 Pyruvate H e - Lactate FAD + 2 H e - FADH Fumarate H e - Succinate Cytochrome b(fe 3+ ) + e - Cytochrome b(fe 2+ ) +75 Ubiquinone + 2 H e - UbiquinoneH Cytochrome c(fe 3+ ) + e - Cytochrome c(fe 2+ ) +254 NO H e - NO H 2 O +421 NO H e - NH H 2 O +440 Fe 3+ + e - Fe O H e - 2 H 2 O +840

10 Electron donors and acceptors determine energy and metabolism Thermodynamics: G = - n. F. E (kj/mol) With n: the number of electrons exchanged F: Faraday s number (96485 C/mol) E: potential difference electron donor/acceptor

11 Examples Take glucose in aerobic metabolism (24 electrons); starting from NADH / O 2 : E = 1.16 V Hence G = kj/mol Glucose for sulphate reduction: only -185 kj/mol Hydrogen formation occurs at -420 mv energy investment to recover NADH C 6 H 12 O H 2 O 2 CH 3 COOH + 4 H CO 2

12 Examples Methanogenesis from glucose: for full conversion still only a G = -115 kj/mol BUT: CH 4 can be oxidized by other organisms WithO 2 : still 857 kj/mol methane, so 90% of energy can be recovered WithSO 4 2- : only 23 kj/mol methane, the rest of the energy is now comprised in the S 2- The anaerobic methane oxidation is omnipresent, despite this limited energy yield redox chains in sediments

13 Examples Evidently, the bacterial cell yield is directly related to the energetic yield of the electron donor/acceptor reaction Y ox ~ 0,4 ; Y H2 ferm ~ 0,03; Y meth ~ 0,05 Also substrate availability, temperature influence this Dehalogenation reactions, BTEX degradation are driven by small stepwise potential increases Often, intermediairy steps have a positive G

14 Rationale for redox chains Bacteria will always strive to use the highest potential electron acceptor if their enzymes allow this Bacteria have to find an electron acceptor

15 What if the electron acceptor is not soluble? In sediments, goethite is omnipresent FeIII = insoluble Dissimilatory versus assimilatory reduction Here we focus on dissimilatory reduction Bacteria develop mechanisms: Chemotaxis by means of a flagellum Membrane associated electron transfer (Geobacter type) Mobile shuttles transport electrons (Pseudomonas type Conductive pili??? (Schewanella? Geobacter?)

16 Direct electron transfer Electrode Cy? e- e-

17 Indirect electron transfer Electrode Bact. B 2 H + H 2 Med ox Med red Med ox Med red 2 H + H 2 2 H + H 2 Med ox Med red 2 H + H 2 PP OM Hy Hy Hy Cy IM e- e- e- e-

18 Biofilm Ecology = mediator producing species = non-producing species C n H 2N O n A C n H 2N O n B C n H 2N O n C C n H 2N O n D CO 2 CO 2 CO 2 e - e - e - CO 2 e - Reduced substrate e - ANODE

19 II. Basic electrochemistry

20 Electrochemistry = pretty simple in comparison to microbiology V(Volt) I (Ampère) Ω = V / I Q = I x t (Coulomb) P = V x I (Watt, J/s) E = P x t (Joule)

21 Electrochemical vs. Carnot Carnot: ε = (T 1 T 2 )/T 2 Efficiency Maximum 36% Electrochemical conversion: ε = n F V H Efficiency up to 110% e

22 Batteries and fuel cells Anode: oxidation occurs, electrons go to an electrical circuit Cathode: reduction occurs, electrons arrive from an electrical circuit

23 H 2 e - e - H 2 O O 2 2 H + H + Fuel cell Anode Cathode

24 Microbial Fuel Cell Glucose MED red e - e - MED ox H 2 O NAD + NADH CO 2 O 2 NAD + H + H + Anode Cathode

25 Biofuels: A Comparison Hydrogen gas: ~122 MJ/kg Biogas: ~56 MJ/kg Biodiesel: ~40 MJ/kg Ethanol: ~27 MJ/kg Glucose: ~16 MJ/kg Acetate: ~14 MJ/kg Wastewater COD: ~11 MJ/kg (domestic) Total WW: ~ MJ/kg Biodiesel: ~34 MJ/l Glucose: ~25 MJ/l Ethanol: ~22 MJ/l Acetate: ~21 MJ/l Wastewater COD: ~15 MJ/l (domestic) Hydrogen in metal hydr: ~5 MJ/l Biogas: ~0.021 MJ/l Hydrogen gas: ~0,010 MJ/l Total WW: ~0.010 MJ/l

26 We see: When we observe bacteria generating electricty Cell voltage: E c = E cathode E anode Current: I Ohms law: E c = R x I Power: P = E c x I When the electrical circuit is opened No current flows Maximum Cell voltage: Open Circuit Voltage (OCV)

27 When bacteria generate electricty We don t see: The potential of the anode and cathode The electrochemical kinetics The electrochemical losses Let s take a closer look at this!

28 Anode and cathode potential In a fuel cell: Anode = oxidation = negative terminal = black Cathode = reduction = positive terminal = red As described the Nernst equation describes the equilibrium potential of a chemical as: When no current flows, an exchange current occurs between an electrode and a solution: Ox + n e - Red (dynamic steady state) i = i a,0 + i c,0 = 0 => -i c,0 = i a,0 = i 0 = exchange current When E differs from equilibrium potential (E ) a current flows: E > E : [Red]/[Ox] : i a,0 > i 0 E < E : [Red]/[Ox] : -i c,0 > i 0 Thermodynamics: potential difference large enough to onset reaction

Anode and cathode potential By using a reference electrode the anode and cathode potential can be measured. The potential of the reference electrode is known and fixed (see table).

29 Anode and cathode potential By using a reference electrode the anode and cathode potential can be measured. The potential of the reference electrode is known and fixed (see table). Now, the electrode potential can be calculated versus a selected reference electrode Common Name SCE Calomel Mercurous sulphate Mercurous oxide Silver chloride Copper sulphate Zinc/Seaw ater Electrode Hg/Hg 2 Cl 2 / sat. KCl Hg/Hg 2 Cl 2 / 1M KCl V vs NHE Hg/Hg 2 SO 4 /sat. K 2 SO Hg/Hg 2 SO 4 /0.5M H 2 SO Hg/HgO/1 M NaOH Ag/AgCl/s at. KCl Cu/sat. CuSO Zn/seawat er -0.8

30 Electrochemical kinetics Exchange current density i o : indicates how fast/slow a reversible reaction occurs There exists a relation between the anodic/cathodic current and the electrode potential The Butler-Volmer equation describes how the electrochemical current density changes by a change of the electrode potential. i = i 0 e α n F ( E E R T 4 parameters: Exchange current Electrode potential : E Concentrations of the components at the surface (not shown) Temperature 0 ) e (1 α ) n F ( E E R T 0 )

31 Electrochemical kinetics E E = η = activation overpotential Activation overpotential = extra potential needed to counter the slow kinitics of a reduction or oxidation η = E E E= η + E Positive η: oxidation at anode observed E > E Negative η: reduction at cathode observed E < E

32 Electrochemical kinetics Current in function of E-E = η η = 0 : i = 0 : equilibrium i o determines the magnitude of η (or the voltage (E) applied) before a current starts to flow Curve a: reversible system = i o large (good electrode, good cat.) Curve c: irreversible system = i o small (bad electrode, bad cat.)

Ch. 13 Fundamentals of Electrochemistry

Ch. 13 Fundamentals of Electrochemistry 13.1 13-1. Basic Concepts of electrochemistry redox reaction : reactions with electron transfer oxidized : loses electrons reduced : gains electrons Fe 3+ + V 2+