Chemistry
Cell Potential
Cell potential, also known as electromotive force, is a measure of the potential difference between the two electrodes of an electrochemical cell. It represents the driving force for the movement of electrons through the cell. Cell potential is a key concept in understanding the behavior of electrochemical cells and is crucial in determining the feasibility and direction of redox reactions.
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eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
cell ).Example 10.6: Determining the Cell Potential of a Concentration CellWhat is the Cell Potential of a galvanic cell composed of two Ag0 /Ag+ half cells, one with a silver ion concentration of 0.01 M and the other with a silver ion concentration of 1 M.1. Determine the oxidation-reduction reaction.The cell with the lower concentration (0.01 M) becomes the anode:Ag 0s →Ag +aq0.01 M+1e −anodeThe cell with the higher concentration becomes the cathode:Ag+ (aq) [1.0 M] + 1e− → Ag0 (s) (cathode)The chemical equation for the oxidation-reduction reaction is:Ag0 (s) + Ag+ [1.0 M] → Ag+ [0.01 M] + Ag0 (s)2. Determine the Cell Potential.E = −0.059log Qne −where: E 0 = 0, n e − = 1, Q = [0.01 M]/[1.0 M]E =E 0– 0.059 × log/Ag +0.01 MAg +1.0 M= 0 V –0.059− 2= 0.118 VCase Study: The pH MeterSolution pH can be determined by the color of a chemical indicator, as described in Chapter 7 . But, these measurements are only approximate determinations of solution pH and a more quantitative method is required when more accurate measurements of pH are needed. Solution pH can be determined accurately by using the principles of a concentration cell based on the H2 (g)/H+ (aq) half reactions. The potential generated between two half cells with different H+ (aq) concentrations is measured to determine the difference in the H+ (aq) concentrations electrochemically. This was originally accomplished by using two hydrogen electrodes as shown in Fig. 10.7 . Both electrodes were constructed like the SHE, but one was submerged in a solution of unknown pH (the sensing electrode) and the second was submerged in a solution of 1 M H+
- Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
The AP Chemistry exam requires students to see the relationship among ideas presented in different parts of the AP syllabus. An example of the connection between different chapters is seen in the relationships between free energy, Cell Potential, and equilibrium constants.Relationship of Cell Potential to Gibbs Free Energy
For a voltaic cell, the overall Cell Potential will always be positive and signifies a spontaneous redox reaction. The relationship between the overall Cell Potential and the free energy is shown in the following equation:
The free energy of the reaction can also be used to calculate the equilibrium constant, K, for a reaction.n =number of electrons lost or gained F =Faraday’s constant; 96,500 coulombs/mole E ° =standard Cell Potential
The combination of the two equations can be used to relate the Cell Potential of a redox reaction to its equilibrium constant.R =gas constant, 8.31 J/mol K T =absolute temperature, K K =equilibrium constant TEST TIP You should make sure that the units of ΔG are in Joules, not kilojoules, when calculating an equilibrium constant.EXAMPLE: For the following equation, calculate the standard Cell Potential, the free energy, and the equilibrium constant. SOLUTION:ΔG and Nonstandard Conditions
Although the data for standard conditions is prevalent, most reactions do not occur under these specific circumstances. It is still important to predict spontaneity, though, so it will be necessary to calculate ΔG rather than ΔG°. Following is the relationship between ΔG and ΔG°:R = 8.31 J/mol KT = absolute temperature, KQ = reaction quotientΔG ° = Gibbs free energy under standard conditionsΔG can be related to the Cell Potential at nonstandard conditions, as well. Here, E is calculated using the Nernst equation.Electrolytic Cell
In addition to Voltaic cells, there is another type of electrochemical cell called an electrolytic cell.
eBook - ePubElectrical Engineering
Fundamentals
- Viktor Hacker, Christof Sumereder(Authors)
- 2020(Publication Date)
- De Gruyter Oldenbourg(Publisher)
6 Electrochemistry6.1 Basic electrochemical concepts
With special regard to electrical engineering, this chapter covers the branch of electrochemistry that deals with the generation and storage of electric current. The electrochemical oxidation and reduction reactions take place at the phase boundaries of the electrode and the electrolyte.Galvanic cellChemical energy is transformed into electrical energy, current is produced, and electrochemical reactions take place spontaneously (negative free enthalpy). Galvanic cells are categorised into three subgroups:- Primary cells
- Secondary cells
- Fuel cells
Electrolytic cellElectric energy is transformed into chemical energy. Two electrodes made of electron-conducting material, and the electrolytes with ion conductivity are conductively connected62 to each other. At the two spatially separated electrodes electrochemical reactions take place.Half-cellA half-cell consists of one single electrode and an electrolyte into which the electrode is submerged (e.g. copper in a copper sulphate solution). If a (metal) electrode is submerged into a metal salt solution (same metal), the surface of the electrode becomes charged. With base metals (e.g. zinc) some metal atoms enter the solution and the released electrons stay on the surface of the electrode, which is now negatively charged. The positively charged metal ions remain bound to the negatively charged metal surface. Thereby an electrical double layer is formed where the negative and the positive charges balance each other out. When two half-cells are combined, a galvanic cell (connected through ionic conductor and electron conductor) is formed.Anode
eBook - ePubPEM Fuel Cells
Theory and Practice
- Frano Barbir(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
However, in practice this potential, called the open circuit potential, is significantly lower than the theoretical potential, usually less than 1 V. This suggests that there are some losses in the fuel cell, even when no external current is generated. When the electrical circuit is closed with a load (such as a resistor) in it, as shown in Figure 3-1b, the potential is expected to drop even further as a function of current being generated, due to unavoidable losses. There are different kinds of voltage losses in a fuel cell caused by the following factors: • Kinetics of the electrochemical reactions • Internal electrical and ionic resistance • Difficulties in getting the reactants to reaction sites • Internal (stray) currents • Crossover of reactants Figure 3-1 Fuel cell with a load: (a) in open circuit; (b) load connected. Although mechanical and electrical engineers prefer to use voltage losses, (electro)chemical engineers use terms such as polarization or overpotential. They all have the same physical meaning: the difference between the electrode potential and the equilibrium potential. From the electrochemical engineer’s point of view, this difference is the driver for the reaction, and from a mechanical or electrical engineer’s point of view, this represents the loss of voltage and power. 3.2.1 Activation Polarization Some voltage difference from equilibrium is needed to get the electrochemical reaction going, as shown previously (Equation 3-17). This is called activation polarization, and it is associated with sluggish electrode kinetics. The higher the exchange current density, the lower the activation polarization losses
eBook - ePub- Odne Stokke Burheim(Author)
- 2017(Publication Date)
- Academic Press(Publisher)
The contribution from mixing change both into free energy Δ g and reversible potential E r e v. Sometimes there are other effects that make the available energy deviate even further, like the lack of supreme selectivity in the liquid junction potential. Thus the open circuit potential E O C P is whatever potential we can measure from the cell. b) The Cell Potential is the potential available for work during a reaction process, that is, when current is flowing. It is the work delivered to the outside of the cell, and thus we must subtract for all the internal losses: E c e l l = E r e v − r j − η c − (a + b ⋅ log j). The power density thus becomes P c e l l = E r e v j − r j 2 − η c j − (a + b ⋅ log j) j. Solution to Problem 6.3 Tafel Kinetics from Experiments. a) This is a straightforward calculation. However, because we are going to be interested in numerical determination of Tafel kinetics in c), we add one extra line to the. table: Nr. 1 2 3 4 5 6 7 8 9 10 E / V A g C l 2 −0.20 −0.20 −0.21 −0.22 −0.25 −0.30 −0.39 −0.44 −0.56 −0.68 j /A cm −2 0.000 0.002 0.005 0.010 0.021 0.062 0.398 1.00 10.0 100 log j – 2.81 −2.33 −2.02 −1.67 −1.21 −0.40 0.00 1.00 2.00 i – j 1–2 2–3 3–4 4–5 5–6 6–7 7–8 8–9 9–10 b i − j −0.07 −0.10 −0.12 −0.12[--=PLGO-SEPARAT. OR=--]−0.12 −0.12 b) Graphically, we proceed as follows, following the figures from left to right: First, we plot the potential versus the logarithmic current density. Next, we determine the open Cell Potential. Third, we draw the Tafel slope from the region where we have a linear E- log j behavior. Here we can read the slope, the coefficient b. Finally, where the Tafel slope crosses the open Cell Potential line, we read (by drawing a vertical line toward the x -line) the logarithmic value of the exchange current density and in turn calculate the current density
eBook - ePub- Allen J. Bard, Roger Parsons, Joseph Jordan(Authors)
- 2017(Publication Date)
- CRC Press(Publisher)
2The Single Electrode Potential: Its Significance and Calculation *ROGER PARSONS† Laboratoire d'Elect rochimie Interfaciale du CNRS, Meudon, FranceIn Chapter 1 the relation of standard electrode potentials to standard thermodynamic quantities was considered. It was emphasized there that by considering whole cell reactions there is no difficulty in principle in relating these quantities, and also that these relationships provide all that is required for the solution of practical problems. Nevertheless, the question of single electrode potentials and their relation to other single ionic properties has existed since the early days of electrochemistry and interest in this question still remains active. This chapter attempts to clarify some of these problems.I. Electrical Potentials in Real Systems
Many of the problems in real systems arise because of the attempt to adapt the classical concept of an electrical potential to a system of real condensed phases. In classical electrostatics the potential is defined in terms of the energy of a test charge and the nature of this test charge is unimportant provided that the potential difference between two points in a uniform medium is considered. This occurs in a number of real situations: for example, twopoints in vacuum or in a dilute gas or, in the most important example, two pieces of metal of the same composition which make up the terminals of a potentiometer or digital voltmeter. It is this last example which is used every time a Cell Potential difference is measured as described in Chapter 1 .However, the classical definition of potential causes difficulty as soon as an attempt is made to consider the potential difference between two points in different media: for example, between a point in vacuum and a point in a condensed phase. This may be made clear by noting that the energy change in taking a test charge from one point to the other depends on the nature of the test charge because of the short-range “chemical” interaction between the test charge and the condensed phase. There is no complete solution to this problem because there is no method for the separation of “chemical” and “electrical” forces since they are both essentially electrical in nature.
eBook - ePub- Adrian Waygood(Author)
- 2018(Publication Date)
- Routledge(Publisher)
Within the cell itself, we see a negative charge moving vertically upwards, through the electrolyte*, along one of the flux lines, and against the force of repulsion offered by the negative electrode. The work done in moving this negative charge increases the potential of that charge, in much the same way as raising a mass against the force of gravity acquires potential energy. The rise in the potential of this negative charge is acquired at the expense of the energy available through the cell’s chemical reaction. At the same time, another negative charge is being repelled by the negative electrode and is moving through an external conductor that links the two electrodes (for the sake of clarity, no load is shown) – to the right of the cell, in Figure 5.9. As this negative charge moves ‘down’ through the conductor, along the flux line towards the positive electrode, the potential that this negative charge acquired moving through the electrolyte, is given up, in much the same way as a mass gives up its potential energy as it falls under the influence of gravity. Again, this is a simplification* of what is actually taking place, chemically, within the cell itself (as we have learnt, it’s ions that move through the electrolyte, not electrons), but the principle applies: the potential acquired by any charged particle (electrons, ions, whatever) as it moves between electrodes within the cell is then lost as it moves through the external conductor. *As we have learnt, in the case of a cell, electrons don’t actually move through the electrolyte but, rather, ions do
eBook - ePubCorrosion Engineering
Principles and Solved Problems
- Branko N. Popov(Author)
- 2015(Publication Date)
- Elsevier(Publisher)
2.5, the electrical energy results from chemical energy that has been released when water is formed: H 2 + 1 2 O 2 → H 2 O Fig. 2.5 Schematic of the hydrogen/oxygen fuel cell. The Cell Potential is estimated using equation: E H 2, O 2 | H 2 O = E H 2, O 2 | H 2 O o + 2.303 RT 4 F log P H 2 P O 2 1 2 This equation was derived considering the partial electrode reaction of hydrogen oxidation and oxygen reduction, H 2 → 2 H + + 2 e − 1 2 O 2 + H 2 O + 2 e − → 2 OH − The overall reaction. is: H 2 + 1 2 O 2 + H 2 O → 2 H + + 2 OH − (2.58) According to the Nernst. equation: E cell = E cell o + 2.303 RT 4 F log P H 2 P O 2 1 / 2 a H 2 O a H + 2 a OH − 2 a H + a OH − = K w E cell = E cell o − 2.303 RT F log K w + 2.303 RT 2 F log P H 2 P O 2 1 / 2 E cell = e H + | H 2 o[--=PLGO-SEPA. RATOR=--]+ e OH − | O 2 o − 2.303 RT F log K w = 0.00 + 0.401 − 0.059 log 10 − 14 E cell = 0.00 + 0.41 + 0.83 = 1.23 V E cell = 1.23 + 0.059 2 log P H 2 P O 2 1 / 2 (2.59) According to Eq. (2.59), the potential of the fuel cell increases by the log of the partial pressure of both oxygen and hydrogen, and is independent of the pH of the solution. 2.8.4 Determination of electrode potential of a standard Weston cell The overall reaction in the Weston cell as presented in Fig. 2.6a and b is Cd, Hg|CdSO 4 |CdSO 4 |Hg 2 SO 4, Hg. Fig. 2.6 (a) Schematic of the standard Weston cell and (b) hypothetical electric potential profile in the Weston cell. The potential of this cell is very stable with a small temperature coefficient. The left half-cell (anode) reaction in the Weston cell is reversible to cadmium ions, Cd Hg → Cd 2 + + 2 e − while the right-hand side (cathode) is reversible to sulfate ions: Hg 2 SO 4 + 2 e − → 2 Hg + SO 4 2 − The overall reaction in the cell is: Cd Hg + Hg 2 SO 4 → Cd 2 + + SO 4 2 − + 2 Hg The emf of the Weston cell
