Bachelorarbeit. In situ state of charge measurement in vanadium redox flow batteries. Niklas Janshen

Transcription

1 Error! No text of specified style in document. 1 Bachelorarbeit Niklas Janshen In situ state of charge measurement in vanadium redox flow batteries Fakultät Technik und Informatik Department Maschinenbau und Produktion Faculty of Engineering and Computer Science Department of Mechanical Engineering and Production Management

2 Error! No text of specified style in document. 2 Niklas Janshen In situ state of charge measurement in vanadium redox flow batteries Bachelorarbeit eingereicht im Rahmen der Bachelorprüfung im Studiengang Maschinenbau/Entwicklung und Konstruktion am Department Maschinenbau und Produktion der Fakultät Technik und Informatik der Hochschule für Angewandte Wissenschaften Hamburg Erstprüfer/in: Prof. Dr. Thorsten Struckmann Zweitprüfer/in : M. Eng. Simon Ressel Abgabedatum:

3 Error! No text of specified style in document. 3 Zusammenfassung Niklas Janshen Thema der Bachelorthesis In situ Ladezustandsbestimmung in Vanadium Redox Flow Batterien Stichworte Vanadium Redox Flow Batterie, Ladezustandsbestimmung, Elektrolytdichte Kurzzusammenfassung Der Ladezustand von All-Vanadium Redox Flow Batterien (VRBs) kann über die Messung von Elektrolyteigenschaften bestimmt werden. Im Rahmen dieser Bachelor Thesis wurde eine Übersicht über die existierenden Methoden der Ladezustandsbestimmung in VRBs erstellt und die Eignung einer Dichtemessung zur Ermittlung des Ladezustands einer VRB evaluiert. Es wurden Lade- und Entladezyklen mit einer 1.6 M Vanadium und 4 M Sulfat Elektrolytlösung analysiert und die Temperaturabhängigkeit der Dichte experimentell bestimmt. Es wurde gezeigt, dass durch die Messung der Elektrolytdichte und Temperatur der Ladezustand der Halbzellen von VRB in situ mit einer Genauigkeit von 2.4 \% für den Anolyten und 7.4 \% für den Katholyten bestimmt werden kann. In der Zukunft kann mit der evaluierten Methode die Kreuzkontamination durch die Membran von VRBs untersucht werden. Zusätzlich kann mit der Vermessung von Elektrolyten unterschiedlicher Zusammensetzung der Anwedungsbereich erweitert werden. Niklas Janshen Title of the paper In situ state of charge measurement in vanadium redox flow batteries Keywords Vanadium redox flow batteries, state of charge monitoring, electrolyte density Abstract The State Of Charge (SOC) of All-Vanadium Redox Flow Batteries (VRB s) can be monitored by measuring the electrolyte properties. Within this bachelor thesis, an overview on the existing SOC monitoring methods for VRB's was given. The applicability of a method to monitor the SOC of a VRB by measuring the electrolyte density was evaluated. Charge and discharge cycles with an electrolyte solution of 1.6 M vanadium and 4 M total sulfate were analysed and the temperature dependency of the density was determined experimentally. It was shown, that the SOC of both half-cells of a VRB can be monitored in situ with an accuracy of 2.4 \% for the anolyte and 7.4 \% for the catholyte by measuring the density and temperature. In future research, the crossover through the membrane of VRB s can be studied by utilizing the evaluated method. In addition, the scope of the SOC monitor can be extended by investigating electrolytes with different compositions.

4 List of Abbreviations and Symbols a) Symbols Symbol Unit Description c i [mole l 1 ] molar concentration or molarity of i c [ml % g 1 ] slope of regression d [% K 1 ] temperature correction factor e [%] correction factor E 0,cell [V ] standard cell potential E 0,cell [V ] formal potential E cell [V ] cell potential f [%] y-intercept of regression F [C mole 1 ] Faraday constant I [A] current I cell [A] cell current m g ml 1 C 1 slope of regression n [mole] amount of substance Q [C] charge Q theo [C] theoretical capacity R [J K 1 mole 1 ] gas constant R 2 [-] coefficient of determination t [s] time T [K] temperature I

5 V [m 3 ] volume z [-] number of electrons per reaction γ [-] activity coefficient i [-] error of i κ [ms cm 1 ] electrolyte conductivity ρ [g ml 1 ] electrolyte density ρ 0 [g ml 1 ] standard electrolyte density φ 0 [V ] standard reduction potential b) Abbreviations blablablablablabla Abbreviation ch dch EES GfE IR OCV RE SOC VRB WE Full form charge discharge Electrical Energy Storage Gesellschaft für Elektrometallurgie Infrared Open Circuit Potential Reference Electrode State Of Charge All-Vanadium Redox Flow Battery Working Electrode II

6 Contents List of Abbreviations and Symbols I 1 Introduction The All-Vanadium Redox Flow Battery (VRB) Electrochemistry of VRB Components of the VRB cell Thesis Objectives State of Charge Monitoring Methods & Electrolyte Properties Formation,Charge and Discharge of the Electrolyte SOC Monitoring Methods in VRB Electrolyte Properties Experimental Temperature Dependence of the Electrolyte Density Setup Measurement Routine Charge and Discharge Cycling Setup Measurement Routine Electrolyte Density Dependent SOC Monitoring Temperature Influence Influence of the SOC on the Electrolyte Density Calibration Validation Measurement Accuracy Conclusions and Outlook Conclusions Method comparison Outlook Required future research efforts Crossover Bibliography 56 1

7 A Appendix III A.1 Error Computation III A.2 Validation of Experimental Set Up VI A.3 Measuring Data VII A.4 Data Sheet - GfE Electrolyte XI 2

8 1 Introduction The worlds energy demand is rapidly increasing with the growing population. Simultaneously fossil fuels are running low and getting harder and more expensive to extract. In addition, by consuming the oil, coal and natural gases of this planet, CO 2 is emitted. This greenhouse gas is under suspicion to be the biggest contributor to global warming. If in the future, the energy demand wants to be met and the CO 2 emission reduced, the energy production sector will have to change extensively. The International Energy Agency has predicted, that the worlds energy demand will grow by 33 %, while the need for electricity will grow by over 60 % until 2035 [1]. Renewable energies will contribute 30 % to the increase of electricity production from 2011 to While hydro power will hold around 50 % of the shares of the renewables, electricity production from wind energy will increase from 10 % to 24 %, while solar energy will increase from 1.4 % to 10 % [1]. Wind and solar are both abundant sources of energy and accessible in all parts of the world, but they are not constantly available. To provide grid stability it is important to generate as much electricity as the demand requires. Due to the difficulty in forecasting solar and especially wind energy production, more attention has been raised to Electrical Energy Storage (EES). The EES can balance demand and generation by accumulating electric energy during periods of high production and low demand and introduce electricity into the grid during periods of the opposite behavior. Additionally, this balancing service is economically valuable, because energy can be stored at low prices during off-peak times and released with high profit during peak times. Furthermore EES can be an assurance in the case of blackouts and can be used for reliable energy supply in off-grid regions [2, 3]. A promising technology for stationary EES is the redox-flow-battery (RFB) [4], which stores electrical energy in two redox couples dissolved in electrolytes. A major advantage of the RFB is the separation of power from energy capacity, which makes it possible to match the specific requirements for a variety of applications. The allvanadium redox flow battery (VRB) is one of the most developed RFB s [5] and a promising stationary EES technology [2]. 1.1 The All-Vanadium Redox Flow Battery (VRB) In the following section the basic principles and underlying electrochemistry of the VRB will be explained and the important components of the VRB will be introduced. In the end, difficulties in operating a VRB will be explained and the thesis objectives will be outlined. 3

9 Figure 1.1: Principle of a VRB [6] In figure 1.1 a schematic of a VRB setup, including the processes occurring during charging and discharging, is depicted. The VRB consists of two electrolyte cycles connected to an electrochemical cell. The latter is divided by an ion exchange membrane into half-cells, the negative half-cell on the right side and the positive half-cell on the left side of the figure. The electrolytes are stored in two tanks and are pumped through the half-cells during operation of the battery. The charge transfer takes place at the electrode surfaces inside both half-cells. The electrolyte is an aqueous solution of vanadium salts in sulfuric acid and is denoted as anolyte, the electrolyte in the negative- and as catholyte, the electrolyte in the positive half-cell. To charge or discharge the battery, the cell has to be connected to a power source or an electrical load Electrochemistry of VRB The following section is focussing on the underlying electrochemistry of a VRB. First the important basic terms and definitions of reduction-oxidation-reactions will be explained, followed by the reactions taking place in a VRB. In the end the equations used in this thesis will be explained and possible occurring side reactions will be introduced. 4

10 Fundamentals and Definitions of Reduction-Oxidation-Reactions To simplify the understanding of reduction-oxidation-reactions (redox-reactions) the following fundamentals and definitions are helpful. A redox-reaction is a chemical reaction in which one or more electrons are transferred between the reacting species. Oxidant + e Product (1.1) Reductant Product + e (1.2) In a reduction reaction (1.1) the oxidant gains an electron, while in an oxidation reaction (1.2) the reductant loses an electron to another species.therefore, the oxidant is reduced and the reductant is oxidized. A reduction decreases the oxidation state, while an oxidation increases the oxidation state of the involved species. e Reductant + Oxidant Oxidant + Reductant (1.3) e When reduction and oxidation reactions appear at both sides of a chemical equation the reaction is called a redox-reaction(1.3). The molarity The molarity or molar concentration is an important unit in chemistry when referring to the concentration of electrolytes. The molar concentration c i is defined as the ratio of the number of moles of the solute n i to the volume of the solution V : c i = n i V Redox Reactions of VRB s in [ ] mole l or [ M] (1.4) When looking at figure 1.1 it can be seen, that the employed redox couples in a VRB are V 2+ /V 3+ in the negative half-cell and V 4+ /V 5+ in the positive half-cell. The vanadium ions in the oxidation states 4 and 5 are known to form a V-O bond [7] and therefore V 4+ occurs as V O 2+ and V 5+ is present as V O 2 +. The reactions taking place in the negative and positive half-cell and the corresponding standard reduction potentials are the following: Negative half-cell charge V 3+ + e discharge V2+ ; φ 0 = 0.26V (1.5) Positive half-cell VO 2+ + H 2 O charge H + + e discharge ; φ + 0 = 1.00V (1.6) 5

11 Overall redox-reaction charge V 3+ + VO 2+ + H 2 O discharge V2+ + V O H + ; E 0,cell = 1.26V (1.7) From the chemical reaction, (1.5) it can be seen, that during charge V 3+ is reduced to V 2+. The electron needed for the reduction is transferred through the outer electrical circuit from the oxidation of V O 2+ to V O + 2 in the positive half-cell. The transferred electron leads to an unbalanced charge of the half-cells. To compensate this, one proton H + from the catholyte crosses the ion exchange membrane towards the anolyte. The opposite reactions occur during discharge of the VRB. The V 2+ is oxidized to V 3+ + in the negative half-cell, losing the electron needed for the reduction of V O 2 to V O 2+ in the positive half-cell. The electron is transferred through the outer electrical circuit to the positive half-cell and one proton crosses the membrane from the anolyte to the catholyte. When looking at equation (1.6) it can be noticed that in the positive half-cell during charge two hydrogen ions are produced and one water molecule is consumed for each oxidized V O 2+ ion. In conclusion, a fully charged VRB contains only V 2+ in the negative half-cell and V O + 2 in the positive half-cell, while a fully discharged VRB contains solely V 3+ in the negative half-cell and V O 2+ in the positive half-cell. The standard reduction potentials of the negative φ 0 and positive φ 0 + half-cell are given in equations (1.5) and (1.6), respectively [8]. The potential is the driving force for the species to be oxidized or reduced and can experimentally be determined when either an electrolyte with V 2+ /V 3+ or V O 2+ /V O + 2 is placed in one half-cell and is measured against a standard hydrogen electrode (SHE) in the other half-cell, under standard conditions. These are defined as K, 1 M of both redox couples and 1 bar pressure. The standard cell potential E 0,cell of the VRB cell can be obtained by: E 0,cell = φ + 0 φ 0 (1.8) The standard cell potential can also be measured between the electrodes of the VRB at standard conditions without current flow. Faraday s Law To determine the charge needed to fully charge or discharge the VRB, Faraday s first law of electrolysis can be applied. Q = I t = z F n (1.9) According to Faraday the amount of charge Q, which flows through the outer electrical circuit can be expressed as the number of electrons per reaction z multiplied with the Faraday constant F and the moles of the active species n dissolved in the electrolyte. This amount of charge will be proportional to the applied current I during the time t. 6

12 Nernst Equation As mentioned before, the standard cell potential E 0,cell can experimentally be determined as the voltage measured between the negative and positive electrode in a VRB during standard conditions. By applying the Nernst equation, the cell potential E cell can be determined theoretically for varying conditions [9]: [ E cell = E 0,cell R T cv z F ln O + c 2 V 2+ (c H +) 2 γv O + γ 2 V 2+ (γ H +) 2 ] (1.10) c V O 2+ c V 3+ γ V O 2+ γ V 3+ Where c i is the molar concentration and γ i the activity coefficients of the vanadium species, c + H the proton concentration in the positive half-cell, R the gas constant and T the temperature. The activity coefficients however cannot directly be measured [9] and to circumvent this lack of information the measurable formal potential E 0,cell is introduced: [ E 0,cell = E 0,cell + R T γv z F ln O + γ 2 V 2+ (γ H +) 2 ] (1.11) γ V O 2+ γ V 3+ By combining equations (1.10) and (1.11), the activity constants can be included into the formal potential and the standard cell potential can be determined using the following equation: E cell = E 0,cell R T z F ln [ cv O + 2 c V 2+ (c H +) 2 ] c V O 2+ c V 3+ (1.12) Another common way to avoid the use of the activity coefficients is to assume that they cancel each other out and are equal to one [9], which leads to: [ E cell = E 0,cell R T cv z F ln O + c 2 V 2+ (c H +) 2 ] (1.13) c V O 2+ c V 3+ It should be noted however, that the Nernst equation is only valid for zero current flow in the VRB. Therefore E cell is denoted as Open Circuit Voltage (OCV), also often referred to as open circuit potential. State Of Charge (SOC) The State Of Charge (SOC) is a characteristic value to indicate how much energy is stored in a battery. The SOC can generally be expressed as the percentage of the amount of stored charge at a certain point Q(x) to the maximum possible stored amount of charge Q theo, also be referred to as the theoretical capacity: SOC(x) = Q(x) Q theo 100% (1.14) 7

13 In a VRB the energy is stored within two electrolytes and therefore the SOC of the whole battery is dependent on both half-cells. When using Faraday s Law (1.9) the SOC of a VRB can be calculated for each half-cell by: Negative half-cell Positive half-cell n V 2+ SOC = 100% (1.15) n V 2+ + n V 3+ n V O2 + SOC = 100% (1.16) n V O2 + + n V O 2+ By using the definition of the molarity (1.4) and assuming that the electrolyte mass and volume are constant the SOC of the entire VRB can be calculated by: SOC = c V 2+ c V 2+ + c V % = c V O2 + c V O2 + + c V O % (1.17) Side Reactions The before mentioned chemical equations (1.5), (1.6) were explained under the assumption of an ideal behavior. However, in a real operating battery side reactions can occur. 2 H e H 2 (1.18) 2 H 2 O 4 H + + O e (1.19) Hydrogen evolution (1.18) can occur at low potentials in the negative half-cell, while oxygen evolution and proton production (1.19) can occur at high potentials in the positive half-cell [10], both at a high SOC of the VRB. In addition, the rate of hydrogen evolution is observed to be higher than the rate of oxygen evolution [11]. Due to this, more current will flow into the hydrogen evolution reaction in the positive half-cell, than into the actual charge of the catholyte. Consequently, the electrolytes in the half-cells will have an unbalanced SOC. In addition to the side reactions mentioned above, the V 2+ -ion is reported to be highly unstable. Therefore it can either be oxidized by water [7, 12] or by the oxygen in the air [8, 10] Components of the VRB cell In the following section, the components of a VRB will be introduced and the basic function will be explained. 8

14 Electrode As already mentioned, the electron transfer takes place on the surfaces of the electrodes of the VRB. The cell potential can be measured between the electrode of the positive and the negative half-cell in the absence of current flow. In order to increase the mass transfer and reduce the pressure drop, often porous graphite felts are used as electrodes in VRB. During charge or discharge however, a degradation of the electrode material can occur [13] which can influence the electrode behavior and therefore the measured potentials. Electrolyte In section 1.1.1, the electrolyte was introduced as an aqueous solution of vanadium salts in sulfuric acid. Due to the sulfuric acid H 2 SO 4 more protons H +, as indicated by the reaction of the positive half-cell and sulfate ions SO 4 are present in the solution. The composition of the electrolyte is denoted as the concentration of vanadium c V and total sulfate c SO4 of the electrolyte. Membrane As already discussed in section 1.1.1, the ion exchange membrane is responsible for the transport of protons from one half-cell to the other. Furthermore it plays a crucial role in the set up of the cell, by separating the negative from the positive half-cell. In general there are two types of membranes, only permitting certain ions the passage: 1. Anion exchange membranes 2. Cation exchange membranes The first membranes obviously only permit anions to cross itself, while the second membranes only allow cations to pass from one half-cell to the other. However, in operating VRB s other ions or molecules also happen to migrate through the membrane. This process is called crossover and can also be divided into two types: 1. Selective crossover 2. Bulk crossover Selective crossover describes certain ions or molecules crossing the membrane. Therefore it includes vanadium ions, sulfate ions and water molecules. Bulk crossover occurs when the electrolyte in its entirety crosses the membrane. Both types will lead to an imbalance of electrolyte volume of the half-cells and cause an SOC decrease due to the reaction of the vanadium species with each other. This SOC decrease is often referred to as self-discharge of the VRB. Up to this point two types of imbalances were introduced: a volumetric imbalance between the anolyte and catholyte and an imbalance in the SOC s of the half-cell 9

15 electrolytes. Both types of imbalance, limit the SOC of the entire cell, because one half-cell will limit the charge of the other one. This will lead to a capacity loss and requires rebalancing methods [11]. 1.2 Thesis Objectives When examining the previous section, the following difficulties in the operation of a VRB can be observed: ˆ Overcharging of the battery will lead to undesirable side reactions. ˆ Side reactions, bulk and selective crossover can lead to an imbalance and therefore to a capacity loss. ˆ The SOC of both half-cells need to be monitored in order to detect imbalances and whether one half-cell is in danger of overcharging. Thesis Objectives In order to solve the above listed problems the applicability of a new SOC monitoring method should be evaluated. The SOC of a VRB should be monitored by measuring the electrolyte density and temperature in situ. In detail the following points should be investigated: ˆ Literature research on existing SOC monitoring methods for VRB s should be done and an overview of these should be given. The focus should be on the electrolyte composition, advantages, limitations and on the feasibility of those methods in order to compare them with the new monitoring method. ˆ The temperature influence on the electrolyte density should be experimentally investigated for the anolyte and catholyte of VRB s. Therefore an experimental set up and measurement routine should be developed. ˆ The applicability of an in situ density and temperature measurement of the positive and negative electrolyte to monitor the SOC of a VRB should be evaluated. ˆ The method should be validated in comparison with another validated, existing SOC monitoring method. For that purpose either an OCV or an in situ half-cell potential measurement should be utilized. ˆ An outlook should be given, focussing on the possibility of determining the crossover through the membrane of a VRB, by the application of the evaluated method. 10

16 2 State of Charge Monitoring Methods & Electrolyte Properties The following chapter is divided into three sections. In the first one the detailed processes which occur during charge and discharge of the VRB will be explained. In the second section an overview on the existing SOC monitoring methods will be given, while in the third section studies on the electrolyte density will be presented. 2.1 Formation,Charge and Discharge of the Electrolyte From the previously introduced redox reaction equations (1.5), (1.6) and (1.7) the exact process taking place in the VRB cell can not be observed yet. They neither consider the complete composition of the electrolyte nor the transfer of hydrogen ions through the membrane. In the following section, the processes which take place during charge and discharge of the VRB will be explained. In section 1.1.2, the electrolyte was introduced as an aqueous solution of vanadium species in sulfuric acid. In order to achieve higher energy densities, the vanadium concentration can be increased. Thereby, the dissolved V 5+ -ions become unstable and can precipitate more easily [14]. To avoid precipitation of vanadium species, in some cases a stabilizing agent is added [15]. To produce electrolytes employed in VRB s, different vanadium salts and techniques to dissolve them, are used. The electrolyte used for experiments in this thesis will be used as an example to explain the detailed reactions, which occur during charge and discharge of the VRB. The electrolyte from Gesellschaft für Elektrometallurgie mbh (GfE), used for measurements in this thesis, is produced by the dissolution of 0.8 M vanadyl sulfate (V OSO 4 ) and 0.4 M vanadium(iii) sulfate (V 2 [SO 4 ] 3 ) in a 2 M sulfuric acid solution. Furthermore 0.05 M phosphoric acid (H 3 P O 4 ) is added to increase the stability of the electrolyte. However, in the following section the effect of the phosphoric acid will be neglected. The ratio of V OSO 4 to V 2 (SO 4 ) 3 needs to be 2:1, to obtain a solution with an equal concentration of V 3+ and V O 2+ referred to as V It should be noted, that in the following processes concerning the electrolyte, the sulfuric acid is assumed to be fully dissociated. However, the acid dissociation generally takes two steps [7]: H 2 SO 4 H HSO 4 (2.1) HSO 4 2 H + + SO 4 2 (2.2) 11

17 In the first step (2.1) one bisulfate ion (HSO 2 4 ) and one proton is produced for each dissociated sulfuric acid molecule. In the second step (2.2) another proton and one sulfate ion (SO 4 ) arise for each dissociated bisulfate ion. Whether the dissociation occurs completely or only to the first step is uncertain [6] and amongst others dependent on the total sulfate concentration [7]. Before operating the cell, the formation of the electrolyte has to take place. For this process, the V 3.5+ electrolyte solution is filled into both tanks and charged until the anolyte solely contains V 3+ ions while in the catholyte only V O 2+ ions are present Formation of the Electrolyte In figure 2.1 the formation of an electrolyte solution of 1M V 3.5+ in variable sulfuric acid concentration is depicted. After a first observation, it can be seen that the dissolution of one V OSO 4 molecule leads to one V O 2+ ion and one sulfate ion, while the dissolution of one V 2 (SO 4 ) 3 molecule results in two V 3+ ions and three sulfate ions. In addition for each dissociated sulfuric acid molecule, two protons and one sulfate ion are produced. For further comprehension of the formation and the transfer of hydrogen through the membrane it is helpful to focus on the processes taking place in the negative and positive half-cell separately. By studying the formation in the negative half-cell, shown in figure 2.1, it can be seen that the V 3+ ions arising from the dissolution of the vanadium(iii) sulfate stay in their present form, while the V O 2+ ions are split into one oxide ion (O 2 ) and one V 4+ ion. The oxide ion forms one water molecule with two protons, available from the sulfuric acid dissociation, while the V 4+ ion needs one electron to be reduced to V 3+. The electron needed for the reduction comes from the oxidation of the V 3+ ion in the positive half-cell. When looking at the positive half-cell reactions it can be seen that the V O 2+ ions arising from the dissolution of the vanadyl sulfate stay in their present form, while the V 3+ ions are oxidized to V 4+ donating the previously mentioned electron. To form the V O 2+ ion, one water molecule is electrolysed to produce one oxide ion and two protons. The V 4+ ion and the oxide ion then form the V O 2+ ion, while the protons stay dissolved and increase the proton concentration of the positive half-cell. The electron transferred through the outer electrical circuit leads to an unbalanced charge of the half-cells and one proton is transferred through the ion exchange membrane. The outer electrical circuit is depicted as a dotted line, the membrane as a double line and the proton crossing the membrane as a dashed line. After the formation, the electrolyte has an SOC of 0% according to equation (1.17) and can be used for the actual charge and discharge process. 12

18 Negative half-cell Positive half-cell 0.5 VOSO V 2 SO 4 3 x H 2 O a H 2 SO V 2 SO VOSO 4 x H 2 O a H 2 SO 4 0.5VO SO V 3+ a 2H a SO 4 0.5V SO 0.5VO O 2 a 2H + 2 1H + a SO 4 0.5V e 0.5O 2 1H + 0.5e 0.5H + 0.5V 4+ 1 V 3+ (x H 2 O) 2a H + a SO 4 2 1VO 2+ (x 0. 5 H 2 O) 2a H + a SO 4 2 Figure 2.1: Schematic of the formation process in a VRB for an electrolyte solution of 1M vanadium in variable sulfuric acid concentration Charge/Discharge of the Electrolyte In figure 2.2 the charge and discharge process for the previously mentioned electrolyte solution is depicted. The arrows on the right side of the figure follow the processes occurring during charge. However, by reversing them the discharge would be shown. On the left side of the figure the SOC and the direction of the charge and discharge process is shown. So from top to bottom the electrolyte starts with 0 % SOC and ends with 100% SOC. The starting conditions are the same as those from the end of the formation depicted in figure 2.1. Furthermore the outer electrical circuit is depicted as a dotted line, the membrane as a double line and the transferred proton as a dashed line. After a first observation of the schematic, it can be seen that the sulfate ions in both half-cells stay in their present form and do not vary in concentration throughout the entire charging process. The same observance can be made for the water molecules in the negative half-cell. It can also be seen, that the V 3+ ion in the negative half-cell is reduced to V 2+, by receiving the electron of the V O 2+ ion from the positive half-cell. The latter is thereby oxidized to V O 3+, which forms a V O 2 + ion with an oxide molecule. The oxide molecule is produced by the electrolysis of one water molecule, which produces two more protons in the positive half-cell. Due to the fact that an electron is transferred through the outer electrical circuit, one proton crosses the ion exchange membrane to balance the charge of the half-cells. At the time the electrolyte has an SOC of 100% the proton concentration in both half-cells increased, while the water concentration in the positive half-cell decreased. After looking at the formation, charge and discharge process as well as the acid dissociation it can be seen, that the proton concentration in both half-cells is dependent on the degree of acid dissociation, the total vanadium and the sulfuric acid 13

19 charge SOC=0% 1 V 3+ Negative half-cell 2 (x H 2 O) 2a H + a SO 4 1VO 2+ (x 0. 5 H 2 O) Positive half-cell 2a H + a SO 4 2 discharge 1e 1e 1VO 3+ 1O 2 2H + 1H + SOC=100% 1 V 2+ (x H 2 O) 2a H + a SO 4 2 1VO 2 + (x 1. 5 H 2 O) 2a H + a SO 4 2 Figure 2.2: Schematic of the charge and discharge process for an electrolyte solution of 1M vanadium in variable sulfuric acid concentration. concentration. The increase of the water concentration in the negative half-cell during formation and the decrease in the positive half-cell during formation and charge only depends on the total vanadium concentration. The total sulfate concentration is dependent on the sulfuric acid and the vanadium sulfate concentration. It should be noted that the effect of the dissociation of water was neglected in this section. After looking at the charge, discharge and formation, it can be seen that the electrolyte undergoes significant changes. Due to this, electrolyte properties like viscosity, density, conductivity and ph-value vary during these processes [7]. 2.2 SOC Monitoring Methods in VRB In the previous chapter the basic electrochemistry of VRB was introduced, followed by a brief description of the components of the VRB cell. This insight leads to important criteria for the assessment of SOC monitoring methods. As mentioned in section 1.2, the detection of any imbalances arising from side reactions or crossover is a necessary ability of a SOC monitoring system. Whether or not a method can be executed in or ex situ is of importance for the future implementation in commercial systems and therefore crucial for the feasibility. Furthermore the advantages or limitations compared to other methods will be discussed in order to compare them with the evaluated method. In addition, the knowledge of the electrolyte composition is critical for comparison with own present or future results. As mentioned in section 2.1 different techniques are used to produce vanadium electrolytes. However, it is assumed, that those will not influence the electrolyte properties, if the same composition is obtained. Only by adding stabilizing agents the electrolyte properties might change. Consequently, 14

20 SOC monitoring methods coulomb counting via potential measurements spectroscopic methods via electrolyte properties potentiometric titration Open Circuit Potential electrolyte halfcell potentials UV/Vis infrared other conductivity absorption transmission Figure 2.3: Overview of SOC monitoring methods used in VRB. a comparison with results which are obtained with electrolyte solutions of the same composition, but without stabilizing agents can be difficult. In the analysed literature, either self produced electrolytes were used without any stabilizing agents or the 1.6 M vanadium and 4 M total sulfate from GfE was used. Because of this, the techniques to produce the electrolytes are not mentioned in the following section. The existing SOC monitoring used for VRB s, which were found during the literature research, are depicted in figure 2.3. Some of them can be further divided into SOC monitoring by measuring electrolyte properties or potentials and spectroscopic methods Coulomb Counting The cell current integration over time is called coulomb counting. The SOC can hence be determined using a theoretical calculated possible charge amount Q theo by applying Faraday s law (1.9): SOC(t) = t 0 Idt Q theo 100% (2.3) This is a common method for a rough SOC estimation and was for instance used by Aaron et al. [16] for the characterization of their cell design at an SOC of approximately 60 % with an electrolyte solution of 1 M vanadium and 6 M total sulfate. Coulomb counting is a very simple method and easy to apply during operation, but it does not consider any side reactions and therefore any current flowing in those will lead to errors. If these errors are not corrected they will accumulate over each charge and discharge cycle. 15

21 2.2.2 Potentiometric Titration For the potentiometric titration the electrolyte needs to be extracted from the VRB and diluted. Afterwards the potential between two electrodes is recorded while the active species are oxidized with an appropriate oxidant. The recorded voltage is plotted over the titrated volume and any voltage jumps indicate the oxidation states of the vanadium species. Becker et al. [17] used potentiometric titration for the characterization of the current density distribution at different SOC s in VRB s with an electrolyte solution of 1.6 M vanadium and 4 M total sulfate purchased from GfE. Petchsingh et al. [18] used potentiometric titration for the calibration of their spectroscopic SOC monitoring method for the positive half-cell in VRB for varying electrolyte compositions, M vanadium and 4 M total sulfate. Potentiometric titration is a precise method to determine the SOC, but it is only applicable ex situ. It includes also complicated steps and is therefore time-consuming. Furthermore the anolyte and catholyte need to be treated with different oxidants and diluted in different acids [17, 19] SOC Monitoring via Potential Measurements Potentiometric titration not only includes potential measurement, but chemical treatment as well, the following methods are only dependent on the measurement of potentials. These can generally be divided into measuring the cell potential of the VRB, referred to as Open Circuit Voltage (OCV) and the potential of either the anolyte or catholyte, referred to as electrolyte half-cell potentials. Open Circuit Voltage (OCV) As mentioned in 1.1.1, the SOC in terms of equation (1.17) is dependent on the ratio of V 5+ to V 4+ ions and V 2+ to V 3+ ions. Therefore the SOC can be calculated, dependent on the OCV, using one of the forms of the Nernst equations. As mentioned before, there are several forms of the Nernst equation. In some studies the proton concentration in the positive half-cell is assumed to be constantly equal to one. In figure 2.4, the OCV over SOC for an electrolyte solution of 1 M vanadium in 1 M sulfuric acid is shown. The dashed line shows the OCV for a constant proton concentration of 1 M in the positive half-cell throughout the charge and discharge process while the solid line corresponds to the OCV with a varying proton concentration according to the schematic depicted in figure 2.2. Generally, it can be seen that the OCV increases strongly during the first 5 % of the SOC range. Afterwards it increases approximately linearly and ends with another strong increase from 95 %- 100% SOC. In addition an average deviation of approximately 0.05 V can be calculated between the two depicted graphs. Tang et al. [20] determined the SOC via OCV in order to compare the results with a spectroscopic measured SOC for an electrolyte solution of 1 M vanadium and 5 M total sulfate. The OCV was calculated with an initial proton concentration of 4 M 16

22 OCV with varying c H + OCV OCV [V] SOC [%] Figure 2.4: Calculated OCV for an electrolyte solution of 1 M vanadium in 1M sulfuric acid in dependence of the SOC with included c H + variation, solid line, and without c H +, dashed line. and a standard potential of 1.26 V using equation (1.13) and a mean deviation of approximately 7 % between the calculated and the measured OCV values was found. P. Pyka [21] monitored the SOC via OCV measurements for an electrolyte solution of 1.6 M and 4 M total sulfate from GfE. Two models were compared: one including the varying proton concentration with a standard cell potential of 1.26 V using equation (1.13) and one neglecting the proton concentration with a formal potential of 1.4 V using equation (1.12). Both were found to be equally precise, while deviations appeared at an SOC of above 90 %. Furthermore, the formal potential of the second model was adjusted to V. The above presented SOC monitoring methods using OCV measurements can be used in situ, but they cannot be applied during cell operation. Ngamsai et al. [22] determined the SOC via OCV for electrolyte solutions with compositions of M vanadium and 3-5 M total sulfate for the calibration of an SOC monitoring method utilizing the electrolyte conductivity. The OCV was reported to increase with increasing total sulfate concentration and was measured in an OCV-cell. The cell was implemented before the actual cell, which enables SOC monitoring during cell operation. Nevertheless, in the additional cell more crossover can occur. 17

23 In conclusion, SOC monitoring via OCV measurement can be applied without any additional equipment, except the OCV-cell, but it cannot detect any imbalance. As shown in figure 2.4, the effect of the varying proton concentration increases the OCV with a constant value. At the moment, the varying proton concentration is neglected, but the formal potential is adjusted to a higher value, the SOC can be measured with the same precision. Thereby a known influence is neglected in order to use a less complex model. Furthermore the measured potentials might be influenced by electrode degradation processes [13], which can occur during cell operation. Electrolyte Half-Cell Potentials By applying the Nernst equation separately to both half-cells the theoretical reduction potentials of the positive φ + and the negative half-cell φ can be obtained: Negative half-cell Positive half-cell φ = φ 0 R T z F ln [ cv 2+ c V 3+ φ + = φ + 0 R T [ ] z F ln c V O 2+ c V O2 + (c H +) 2 ] (2.4) (2.5) For the measurement of the reduction potentials, now referred to as the electrolyte half-cell potentials, a reference electrode (RE) with a known potential is needed [23]. To measure the electrolyte half-cell potential the reference electrode is placed in short distance to a working electrode (WE), both immersed inside the electrolyte. The type of the reference electrode in the following section is given as RE against WE. Corcuera et al. [11] monitored the SOC via electrolyte half-cell potentials measured by a Hg/Hg 2 SO 4 in 2 M H 2 SO 4 against a carbon rod immersed in both tanks of the VRB. The proton concentration was assumed to be one and with this a formal half-cell potential of V for the negative and V for the positive half-cell was measured. Rudolph et al. [24, 25] measured the electrolyte half-cell potentials by the use of a Ag/AgCl against a unknown working electrode in the tanks for an electrolyte solution of 1.6 M vanadium and 4 M total sulfate from GfE. The determined SOC was used to calibrate an SOC monitoring method, which utilizes an infrared (IR) sensor. Depending on the tank size and electrolyte volume used, the two abovementioned methods can lead to errors, due to stratification of the electrolyte in the tanks. P. Pyka [21] monitored the SOC via electrolyte half-cell potentials measured by a Hg/HgSO 4 in 2.5 M H 2 SO 4 against a glassy carbon rod for an electrolyte solution of 1.6 M vanadium and 4 M total sulfate from GfE. Flow through RE s employed after the cells were used to avoid the stratification in the tanks. For the anolyte, φ 0 was adjusted to V and for the catholyte, φ 0 + was adjusted to V. For the positive half-cells good agreements between the calculated and measured SOC 18

24 values were found until an SOC of 80 %, while for the negative half-cell deviations under 10 % and above 90 % were reported. By monitoring the half-cell potentials an imbalance can be detected. However, reference electrodes are influenced by the proton concentration [10], which is further dependent on the sulfuric acid and vanadium concentration as well as the sulfuric acid dissociation. Additionally, the electrolyte can diffuse into the reference and alter the potential. Furthermore, to ensure the potential is stable, the RE has to be referenced itself after an experiment. Another error can occur due to impurities inside the reference electrode, which can corrosively react with the electrode material, change the activity of the reacting species or the properties of the electrolyte [23]. Limitations of SOC monitoring via potential measurement in general are, that the potential of the VRB around an SOC of 50 % changes very little with increasing or decreasing SOC. Therefore, in this SOC range small offsets can lead to significant errors. In addition, the theoretical potentials need to be adjusted to the measured values Spectroscopic Methods As already mentioned in section 2.1, the electrolyte changes significantly during charging and discharging. From table 2.1 it can be seen, that each of the four oxidation states of the vanadium species has their specific colour [26]. Spectroscopic measurement methods utilize this dependency to measure the SOC via absorption or transmission spectra. UV/Vis Spectroscopy Spectroscopic measurements in the visible spectrum of light are called UV/Vis spectroscopy, which can further be divided into absorption- and transmission spectra measurements. Skyllas et al. [26] monitored the SOC via absorption spectra measurements using a wavelength of 750 nm for an electrolyte solution of 2 M vanadium and 5 M total sulfate. The measurement was conducted ex situ on reference solutions and reported as not feasible for the catholyte due to the high absorbance of the V 4+ /V 5+ solution. However for the anolyte a feasibility over the range of % was reported. P. Pyka [21] monitored the SOC of an electrolyte solution of 1.6 M vanadium and 4 M total sulfate from GfE in situ in a bypass utilizing a wavelength of 405 nm. For the positive half-cell the method was reported as not feasible due to non-linear behavior of the absorbance spectra. Petchsingh et al. [18] monitored the SOC of the positive half-cell utilizing two wavelengths and a mathematical model to simulate the spectra of the V 4+ /V 5+ solution. Different electrolyte solutions with M vanadium and 4 M total sulfate were 19

25 negative half-cell positive half-cell oxidation state specific colour violet green yellow blue Table 2.1: Different colours of the electrolyte corresponding to the oxidation state of the vanadium species used. Liu et al. [27, 28] monitored the SOC of the positive- and negative half-cell in situ by investigating the entire transmission spectra. Different electrolyte solutions with M vanadium and M total sulfate, confirmed by potentiometric titration with an average error of 1.1 %, were used and a dependence of the V 4+ /V 5+ spectra between an SOC of % on the total sulfate concentration was found. Infrared Rudolph et al. [24, 25, 29] utilized an IR sensor in the range of 950 nm to monitor the SOC of an electrolyte solution of 1.6 M vanadium and 4 M total sulfate from GfE in situ. The sensor was calibrated using the previously mentioned Ag/AgCl RE. With this an error between the calibration curve and the 8th degree polynomial equation of 1 % was reported. For the catholyte the method can only detect 0 % and 100 % SOC. In conclusion, spectroscopic methods can monitor the SOC in situ, but additional equipment is needed. A bypass and an additional pump in case of [21] or a home made transmission spectrum analytical system in case of [27, 28]. Furthermore, the monitoring of the positive half-cell is reported to be impossible [21, 26], limited possible [24, 25, 29] or only by establishing a huge database utilizing the entire transmission spectrum of the catholyte [27]. Also in [18] a series of different electrolyte mixtures needed to be analysed to establish the monitoring method. Additionally, in [27, 28] at an SOC between 50 % and 100 %, a dependence of the catholyte spectra on the total sulfate concentration was found SOC Monitoring via Measurement of Electrolyte Properties As mentioned in section 2.1, the electrolyte undergoes changes during the charge and discharge process. Thereby, after studying the change of electrolyte properties with varying SOC, the latter can be monitored by measuring the former. Conductivity Skyllas et al. [26] measured the electrolyte conductivity ex situ of different reference solutions of M vanadium and 4-5 M total sulfate at 22 C. 20

26 Ngamsai et al. [22] monitored the SOC in situ via electrolyte conductivity for 1-2 M vanadium and 3-5 M total sulfate solutions. In both studies, the conductivity was reported to increase with increasing total sulfate concentration and to decrease with increasing vanadium concentration. In addition, Skyllas et al. [26] monitored the SOC for a 2 M vanadium and 5 M total sulfate electrolyte solution at 10, 20 and 30 C and a linear increase of the conductivity with increasing SOC for both half-cells was reported. Corcuera et al. [11] monitored the SOC via electrolyte conductivity for an electrolyte solution of 1.6 M and 4.2 M total sulfate at 10, 22 and 45 C. A Comparison of the calculated conductivity values with experimental data showed good correlation and a higher precision for the positive half-cell. All in all, the feasibility of the conductivity as an indicator for the SOC was reported in all studies. Nevertheless, the high dependency of the conductivity on the temperature increases the measurement effort. In addition, varying total sulfate, vanadium and proton concentration, due to crossover or side reactions, will lead to errors if they are not considered. ph-value As mentioned in section 1.1.2, the hydrogen concentration in both half-cells will ideally change during charge and discharge according to figure 2.2. The ph-value represents the amount of hydrogen ions in an aqueous solution. S. Ressel [6] studied the ph-value in dependence of the SOC in situ for an electrolyte solution of 1.6 M and 4 M total sulfate. The ph-value as an SOC monitor was reported as feasible for the positive half-cell. The ph-value in the negative half-cell remained constant throughout the experiment. Other Skyllas et al. [7] presented in two graphs the dependency of the electrolyte viscosity and density on the SOC for a 2 M vanadium and 5 M total sulfate electrolyte at 10, 20 and 30 C. However, they referred to a patent from 1990 [30], in which neither the two graphs nor the SOC as a function of density and viscosity were mentioned. Therefore it remains uncertain how these values were measured and they will not be considered in the following sections. Furthermore it should be noted, that the exact determination of the boundaries of the SOC range are not easy to obtain. In the end of a charge process, the exact value of 100 % SOC is difficult to detect because the possibility of overcharging the VRB. Furthermore at the end of the discharge, the VRB can go below 0 % into the formation region. In conclusion, an indicator is needed to detect whether the battery is fully charged or discharged. In table 2.2, the indicators for 0 % and 100 % SOC, used in the analysed literature, are listed. The found indicators are explained in the following paragraph. 21

27 indicator for 100 % SOC specification and reference cut off current 4 ma/cm 2 [16, 20], 6 ma/cm 2 [10] change of electrolyte colour [11, 18, 22, 26] constant measurement signal OCV [22], electrolyte half-cell potential [18, 31], absorption [21] indicator for 0 % SOC potential jump electrolyte half-cell potential [18] maximum in absorption [21, 24] reference solution only applicable for the catholyte - dissolution of V OSO 4 in sulfuric acid [10, 26, 31] Table 2.2: Indicators for a fully charged and discharged VRB, an SOC of 100 % and 0 %, respectively. When charging the VRB at a constant voltage, the current drops over the charging process, but it will not reach zero, because of the side reactions. One way to define a fully charged battery is to use a cut off current. However this indicator is highly dependent on the ohmic resistance of the used VRB cell. A cell with a high ohmic resistance will reach the cut off current earlier than one with a lower ohmic resistance. Another common way is to simply take the moment when the colour of the electrolyte changes to violet in case of the anolyte and yellow in case of the catholyte. The third found indicator for a fully charged VRB is the time when a measurement signal, which represents the SOC, stays constant. Measurement signals, which can be utilized for this purpose are the OCV, the electrolyte half-cell potential and the absorption. As an indicator for 0 % SOC the half-cell potential can be monitored, because after the formation process a sharp increase in the potential of both half-cells was reported and theoretically supported by the Nernst equation. Spectroscopic methods use the maximum in the absorption signal. Another way often used to ensure an SOC of 0 % is to dissolve V OSO 4 in sulfuric acid to obtain a solution containing only V Electrolyte Properties In the following section, a short introduction into the aqueous chemistry will be given, followed by a brief overview of the temperature stability of the vanadium electrolyte in order to establish the temperature range for the measurement routine in the following chapter. Afterwards, studies on the density of vanadium electrolyte solutions. They will be used for comparison with own measured densities in order to estimate the effect of crossover on the density. In general, aqueous electrolyte solutions are produced by using water as the solvent 22

28 and adding one or more solutes to produce dissolved cations and anions. In the case of the VRB, the solutes are vanadium sulfates and sulfuric acid. The charge of those ions will have an impact on the dipole water molecules and hence disturb the prior configuration of them. This will cause a change of the physical properties of the solution. The ions in the aqueous solution have stronger charges than the hydrogen and oxygen ions of the water and therefore the water molecules are forced to form hydration spheres around the ions. Within these spheres, the oxygen ions are oriented towards the positively charged ion, while the hydrogen ion is attracted towards the negatively charged ion [12]. Hence, in case of the vanadium electrolyte, the oxygen ions are attracted towards the vanadium ions and the hydrogen ions are oriented towards the sulfate ions. In [32], Xiao et al. reported that both, V 2+ and V 3+ electrolyte solutions are stable until minus 25 C and increase in stability with increasing temperatures. For the V 4+ and V 5+ electrolyte solutions, the opposite behavior occurred and V 5+ was reported to precipitate at 35 C. Studies on Electrolyte Density F. Rahman [14] studied the density of different electrolyte solutions, containing 2-5 M V 5+ and 5-7 M total sulfate at 20 C. The V 5+ solutions were prepared by electrolytic oxidation of V 4+ solutions, afterwards analysed by inductively coupled plasma mass spectrometry and adjusted if needed. The measured densities are presented in table 2.3. V 5+ concentration [M] total sulfate concentration [M] Table 2.3: Densities of different V 5+ electrolyte solutions at 20 C in g/ml [14] A. Mousa [12] studied the densities of V 2+ and V 3+ solutions over a range of C for various vanadium and total sulfate concentrations in the case of V 3+. The V 2+ solutions were reported to be highly unstable and therefore only the densities at 25 C for 1-2 M vanadium and 3-4 M total sulfates were reported. The V 2+ and V 3+ solutions were prepared by electrolytic reduction of V 4+ solutions and the concentrations were confirmed by potentiometric titration, with a reported relative error of 10 %. The measured densities are presented in table 2.5 and 2.4 for the V 3+ - and V 2+ solutions, respectively. 23

29 c V SO4 [M] ρ [g/ml] Table 2.4: Densities of V 2+ electrolyte solutions at 25 C in g/ml [12] Table 2.5: Densities of different V 3+ electrolyte solutions in the range of C in g/ml [12] The V 4+ solutions used for both presented studies were prepared by the reaction of an equimolar mixture of vanadium trioxide (V 2 O 3 ) and vanadium pentoxide (V 2 O 5 ) in diluted sulfuric acid according to equation 2.6 [12, 14]: V 2 O 3 + V 2 O 5 + 8H + 4 VO H 2 O (2.6) The densities were measured, by recording the mass of a glass density bottle with known volume and a relative error of 2 % in the case of [12]. From tables (2.5) and (2.3), it can be seen that the density of vanadium solutions increases with increasing total sulfate and vanadium concentration. In both studies the increase due to higher vanadium concentration was reported to be larger. 24