19.3: Voltaic (or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions, 19.5: Cell Potential, Gibbs Energy, and the Equilibrium Constant, Balancing Redox Reactions Using the Half-Reaction Method, Reference Electrodes and Measuring Concentrations, information contact us at info@libretexts.org, status page at https://status.libretexts.org, laboratory samples, blood, soil, and ground and surface water, groundwater, drinking water, soil, and fertilizer. Appendix: Periodic Table of the Elements; Appendix: Selected Acid Dissociation Constants at 25°C; Appendix: Solubility Constants for Compounds at 25°C; Appendix: Standard Thermodynamic Quantities for Chemical Substances at 25°C; Appendix: Standard Reduction Potentials by Value; Glossary; Versioning History Standard Electrode Potentials. To measure the potential of the Cu/Cu2+ couple, we can construct a galvanic cell analogous to the one shown in Figure \(\PageIndex{3}\) but containing a Cu/Cu2+ couple in the sample compartment instead of Zn/Zn2+. The more positive the reduction potential, the stronger is the attraction for electrons. Standard reduction potentials for selected reduction reactions are shown in Table 2. The SHE on the left is the anode and assigned a standard reduction potential of … Each table lists standard reduction potentials, E° values, at 298.15 K (25°C), and at a pressure of 101.325 kPa (1 atm). This book is licensed under a Creative Commons by-nc-sa 3.0 license. Put another way, the more positive the reduction potential, the easier the reduction occurs. To ensure that any change in the measured potential of the cell is due to only the substance being analyzed, the potential of the other electrode, the reference electrode, must be constant. Below is an abbreviated table showing several half-reactions and their associated standard potentials. One beaker contains a strip of gallium metal immersed in a 1 M solution of GaCl3, and the other contains a piece of nickel immersed in a 1 M solution of NiCl2. Hence the reactions that occur spontaneously, indicated by a positive \(E°_{cell}\), are the reduction of Cu2+ to Cu at the copper electrode. According to the EPA field manual, the “Oxidation-Reduction Potential (E h) is a measure of the equilibrium potential, relative to the standard hydrogen electrode, developed at the interface between a noble metal electrode and an aqueous solution containing electro-active redox species”. In addition to the SHE, other reference electrodes are the silver–silver chloride electrode; the saturated calomel electrode (SCE); the glass electrode, which is commonly used to measure pH; and ion-selective electrodes, which depend on the concentration of a single ionic species in solution. standard cell potential for X. Redox reactions can be balanced using the half-reaction method. E° values do NOT depend on the stoichiometric coefficients for a half-reaction, because it is an intensive property. A glass electrode is generally used for this purpose, in which an internal Ag/AgCl electrode is immersed in a 0.10 M HCl solution that is separated from the solution by a very thin glass membrane (part (b) in Figure \(\PageIndex{5}\). Unless otherwise noted, LibreTexts content is licensed by CC BY-NC-SA 3.0. If a saturated solution of KCl is used as the chloride solution, the potential of the silver–silver chloride electrode is 0.197 V versus the SHE. Standard reduction potential table; Dissociation constants of acids and bases inorganic; Melting point of solids (table of values) Diffusion coefficient of liquids and aqueous solutions (table of values) Dielectric constant of liquids, gases and solids (Table) Dipole moments of molecules (table of values) A positive \(E°_{cell}\) means that the reaction will occur spontaneously as written. One especially attractive feature of the SHE is that the Pt metal electrode is not consumed during the reaction. If we construct a galvanic cell similar to the one in part (a) in Figure \(\PageIndex{1}\) but instead of copper use a strip of cobalt metal and 1 M Co2+ in the cathode compartment, the measured voltage is not 1.10 V but 0.51 V. Thus we can conclude that the difference in potential energy between the valence electrons of cobalt and zinc is less than the difference between the valence electrons of copper and zinc by 0.59 V. The measured potential of a cell also depends strongly on the concentrations of the reacting species and the temperature of the system. Here we present an alternative approach to balancing redox reactions, the half-reaction method, in which the overall redox reaction is divided into an oxidation half-reaction and a reduction half-reaction, each balanced for mass and charge. Thus the hydrogen electrode is the cathode, and the zinc electrode is the anode. When fluoride ions in solution diffuse to the surface of the solid, the potential of the electrode changes, resulting in a so-called fluoride electrode. All tabulated values of standard electrode potentials by convention are listed for a reaction written as a reduction, not as an oxidation, to be able to compare standard potentials for different substances (Table P1). Oxidation numbers were assigned to each atom in a redox reaction to identify any changes in the oxidation states. This cell diagram corresponds to the oxidation of a cobalt anode and the reduction of Cu2+ in solution at the copper cathode. The SCE cell diagram and corresponding half-reaction are as follows: \[Pt_{(s)} ∣ Hg_2Cl_{2(s)}∣KCl_{(aq, sat)} \label{20.4.37}\], \[Hg_2Cl_{2(s)} + 2e^− \rightarrow 2Hg_{(l)} + 2Cl^−{(aq)} \label{20.4.38}\]. The voltage E′ is a constant that depends on the exact construction of the electrode. The potential of a half-reaction measured against the SHE under standard conditions is called its standard electrode potential. Thus E° = −(−0.28 V) = 0.28 V for the oxidation. Step 6: This is the same equation we obtained using the first method. These show the two forms of many common molecules and the redox relationship between them. If the value of \(E°_{cell}\) is positive, the reaction will occur spontaneously as written. Table 2 lists only those reduction reactions that have E° values posi-tive in respect to the standard hydrogen electrode . In this case, we multiply Equation \(\ref{20.4.26}\) (the reductive half-reaction) by 3 and Equation \(\ref{20.4.27}\) (the oxidative half-reaction) by 2 to obtain the same number of electrons in both half-reactions: \[3OH^−_{(aq)} + 9H^+_{(aq)} + 6e^− \rightarrow 3H_{2(g)} + 3H_2O_{(l)} \label{20.4.28}\], \[2Al_{(s)} + 8H_2O_{(l)} \rightarrow 2Al(OH)^−_{4(aq)} + 8H^+_{(aq)} + 6e^− \label{20.4.29}\]. Redox reactions can be balanced using the half-reaction method, in which the overall redox reaction is divided into an oxidation half-reaction and a reduction half-reaction, each balanced for mass and charge. The glass membrane absorbs protons, which affects the measured potential. The standard cell potential for a redox reaction (E°cell) is a measure of the tendency of reactants in their standard states to form products in their standard states; consequently, it is a measure of the driving force for the reaction, which earlier we called voltage. Dividing the reaction into two half-reactions. \[3CuS_{(s)} + 8HNO{3(aq)} \rightarrow 8NO_{(g)} + 3CuSO_{4(aq)} + 4H_2O_{(l)} \nonumber\]. The half-reaction method requires that half-reactions exactly reflect reaction conditions, and the physical states of the reactants and the products must be identical to those in the overall reaction. Hence electrons flow spontaneously from zinc to copper(II) ions, forming zinc(II) ions and metallic copper. With a sufficient input of electrical energy, virtually any reaction can be forced to occur. Because we are asked for the potential for the oxidation of Ni to Ni2+ under standard conditions, we must reverse the sign of E°cathode. Measured redox potentials depend on the potential energy of valence electrons, the concentrations of the species in the reaction, and the temperature of the system. Step 2: Balancing the atoms other than oxygen and hydrogen. A more complete list is provided in Appendix L. Figure 3. Whether reduction or oxidation occurs depends on the potential of the sample versus the potential of the reference electrode. A negative \(E°_{cell}\) means that the reaction will proceed spontaneously in the opposite direction. Remember loss of electrons is oxidation. . For example, take the following reaction from the citric acid cycle: succinate + FAD fumarate + FADH 2 Legal. We know the values of E°anode for the reduction of Zn2+ and E°cathode for the reduction of Cu2+, so we can calculate \(E°_{cell}\): \[E°_{cell} = E°_{cathode} − E°_{anode} = 1.10\; V\]. Example \(\PageIndex{2}\) and its corresponding exercise illustrate how we can use measured cell potentials to calculate standard potentials for redox couples. A galvanic cell is constructed with one compartment that contains a mercury electrode immersed in a 1 M aqueous solution of mercuric acetate \(Hg(CH_3CO_2)_2\) and one compartment that contains a strip of magnesium immersed in a 1 M aqueous solution of \(MgCl_2\). Write the equation for the half-reaction that occurs at the anode along with the value of the standard electrode potential for the half-reaction. Adding and, in this case, canceling 8H+, 3H2O, and 6e−, \[2Al_{(s)} + 5H_2O_{(l)} + 3OH^−_{(aq)} + H^+_{(aq)} \rightarrow 2Al(OH)^−_{4(aq)} + 3H_{2(g)} \label{20.4.30}\]. Balance this equation using the half-reaction method. The oxidation half-reaction (2I− to I2) has a −2 charge on the left side and a 0 charge on the right, so it needs two electrons to balance the charge: Step 4: To have the same number of electrons in both half-reactions, we must multiply the oxidation half-reaction by 3: Step 5: Adding the two half-reactions and canceling substances that appear in both reactions. This interior cell is surrounded by an aqueous KCl solution, which acts as a salt bridge between the interior cell and the exterior solution (part (a) in Figure \(\PageIndex{5}\). With this alternative method, we do not need to use the half-reactions listed in Table P1, but instead focus on the atoms whose oxidation states change, as illustrated in the following steps: Step 1: Write the reduction half-reaction and the oxidation half-reaction. Neutralizing the H+ gives us a total of 5H2O + H2O = 6H2O and leaves 2OH− on the left side: \[2Al_{(s)} + 6H_2O_{(l)} + 2OH^−_{(aq)} \rightarrow 2Al(OH)^−_{4(aq)} + 3H_{2(g)} \label{20.4.31}\]. F 2 (g) + 2e – 2F – (aq) +2.87. The overall cell potential is the reduction potential of the reductive half-reaction minus the reduction potential of the oxidative half-reaction (E°cell = E°cathode − E°anode). A galvanic cell can be used to determine the standard reduction potential of Ag +. The cell diagram and reduction half-reaction are as follows: \[Cl^−_{(aq)}∣AgCl_{(s)}∣Ag_{(s)} \label{20.4.36}\], \[AgCl_{(s)}+e^− \rightarrow Ag_{(s)} + Cl^−_{(aq)}\]. The standard cell potential (E°cell) is therefore the difference between the tabulated reduction potentials of the two half-reactions, not their sum: \[E°_{cell} = E°_{cathode} − E°_{anode} \label{20.4.2}\]. For example, the measured standard cell potential (E°) for the Zn/Cu system is 1.10 V, whereas E° for the corresponding Zn/Co system is 0.51 V. This implies that the potential difference between the Co and Cu electrodes is 1.10 V − 0.51 V = 0.59 V. In fact, that is exactly the potential measured under standard conditions if a cell is constructed with the following cell diagram: \[Co_{(s)} ∣ Co^{2+}(aq, 1 M)∥Cu^{2+}(aq, 1 M) ∣ Cu (s)\;\;\; E°=0.59\; V \label{20.4.1}\]. Ag(s) This half-reaction equation represents reduction, which occurs at the cathode. Using the figure above, determine the highest possible potential for a voltaic cell using one electrode from upper set and one from the lower set of the mechanism: Anode = … As we shall see in Section 20.9, this does not mean that the reaction cannot be made to occur at all under standard conditions. We can do that by looking at our table here. Now this is an oxidation half-reaction. Balance this equation using half-reactions. The standard hydrogen electrode (SHE) is universally used for this purpose and is assigned a standard potential of 0 V. It consists of a strip of platinum wire in contact with an aqueous solution containing 1 M H+. This allows us to measure the potential difference between two dissimilar electrodes. We can, however, compare the standard cell potentials for two different galvanic cells that have one kind of electrode in common. Ion-selective electrodes are used to measure the concentration of a particular species in solution; they are designed so that their potential depends on only the concentration of the desired species (part (c) in Figure \(\PageIndex{5}\)). We have a −2 charge on the left side of the equation and a −2 charge on the right side. At 25°C, the potential of the SCE is 0.2415 V versus the SHE, which means that 0.2415 V must be subtracted from the potential versus an SCE to obtain the standard electrode potential. We can illustrate how to balance a redox reaction using half-reactions with the reaction that occurs when Drano, a commercial solid drain cleaner, is poured into a clogged drain. Balancing H atoms by adding H+, we obtain the following: \[OH^−_{(aq)} + 3H^+_{(aq)} \rightarrow H_{2(g)} + H_2O_{(l)} \label{20.4.24}\], \[Al_{(s)} + 4H_2O_{(l)} \rightarrow Al(OH)^−_{4(aq)} + 4H^+_{(aq)} \label{20.4.25}\]. We can do this by adding water to the appropriate side of each half-reaction: \[OH^−_{(aq)} \rightarrow H_{2(g)} + H_2O_{(l)} \label{20.4.22}\], \[Al_{(s)} + 4H_2O_{(l)} \rightarrow Al(OH)^−_{4(aq)} \label{20.4.23}\]. If \(E°_{cell}\) is negative, then the reaction is not spontaneous under standard conditions, although it will proceed spontaneously in the opposite direction. In this cell, the copper strip is the cathode, and the hydrogen electrode is the anode. Standard reduction potential table; Dissociation constants of acids and bases inorganic; Melting point of solids (table of values) Diffusion coefficient of liquids and aqueous solutions (table of values) Dielectric constant of liquids, gases and solids (Table) Dipole moments of molecules (table of values) Asked for: balanced chemical equation using half-reactions. Step 3: We must now add electrons to balance the charges. This chemistry video tutorial provides a basic introduction into standard reduction potentials of half reactions. From Table 1 on page 646, the reduction potential for silver is r E° (cathode) = +0.80 V. The half-reaction equation and reduction potential for X is: X(s) !!" The half-reactions that occur when the compartments are connected are as follows: If the potential for the oxidation of Ga to Ga3+ is 0.55 V under standard conditions, what is the potential for the oxidation of Ni to Ni2+? Next we balance the H atoms by adding H+ to the left side of the reduction half-reaction. cathode: \[2H^+_{(aq)} + 2e^− \rightarrow H_{2(g)}\;\;\; E°_{cathode}=0 V \label{20.4.5}\], anode: \[Zn_{(s)} \rightarrow Zn^{2+}_{(aq)}+2e^−\;\;\; E°_{anode}=−0.76\; V \label{20.4.6}\], overall: \[Zn_{(s)}+2H^+_{(aq)} \rightarrow Zn^{2+}_{(aq)}+H_{2(g)} \label{20.4.7}\], Cathode: \[Cu^{2+}{(aq)} + 2e^− \rightarrow Cu_{(g)}\;\;\; E°_{cathode} = 0.34\; V \label{20.4.9}\], Anode: \[H_{2(g)} \rightarrow 2H^+_{(aq)} + 2e^−\;\;\; E°_{anode} = 0\; V \label{20.4.10}\], Overall: \[H_{2(g)} + Cu^{2+}_{(aq)} \rightarrow 2H^+_{(aq)} + Cu_{(s)} \label{20.4.11}\], reduction: \[2H_2O_{(l)} + 2e^− \rightarrow 2OH^−_{(aq)} + H_{2(g)} \label{20.4.13}\], oxidation: \[Al_{(s)} + 4OH^−_{(aq)} \rightarrow Al(OH)^−_{4(aq)} + 3e^− \label{20.4.14}\], reduction: \[Cr_2O^{2−}_{7(aq)} + 14H^+_{(aq)} + 6e^− \rightarrow 2Cr^{3+}(_{(aq)} + 7H_2O_{(l)} \nonumber\], oxidation: \[2I^−_{(aq)} \rightarrow I_{2(aq)} + 2e^− \nonumber\], oxidation: \[6I^−_{(aq)} \rightarrow 3I_{2(aq)} + 6e^− \nonumber\], reduction: \[Cr_2O^{2−}_{7(aq)} \rightarrow Cr^{3+}_{(aq)} \nonumber\], oxidation: \[I^−_{(aq)} \rightarrow I_{2(aq)} \nonumber\], reduction: \[Cr_2O^{2−}_{7(aq)} \rightarrow 2Cr^{3+}_{(aq)} \nonumber\], oxidation: \[2I^−_{(aq)} \rightarrow I_{2(aq)} \nonumber\], reduction: \[Cr_2O^{2−}_{7(aq)} \rightarrow 2Cr^{3+}_{(aq)} + 7H_2O_{(l)} \nonumber\], reduction: \[Cr_2O^{2−}_{7(aq)} + 14H^+_{(aq)} \rightarrow 2Cr^{3+}_{(aq)} + 7H_2O_{(l)} \nonumber\], cathode: \[Cu^{2+}_{(aq)} + 2e^− \rightarrow Cu_{(s)} \;\;\; E°_{cathode} = 0.34\; V \label{20.4.33}\], anode: \[Zn_{(s)} \rightarrow Zn^{2+}(aq, 1 M) + 2e^−\;\;\; E°_{anode} = −0.76\; V \label{20.4.34}\], overall: \[Zn_{(s)} + Cu^{2+}_{(aq)} \rightarrow Zn^{2+}_{(aq)} + Cu_{(s)} \label{20.4.35}\]. The standard cell potential is a measure of the driving force for the reaction. One of the most common uses of electrochemistry is to measure the H+ ion concentration of a solution. So positive .8 volts plus positive .76 volts. \(E°_{cell} = E°_{cathode} − E°_{anode} \] The flow of electrons in an electrochemical cell depends on the identity of the reacting substances, the difference in the potential energy of their valence electrons, and their concentrations. Because electrical potential is the energy needed to move a charged particle in an electric field, standard electrode potentials for half-reactions are intensive properties and do not depend on the amount of substance involved. Have questions or comments? We can use the two standard electrode potentials we found earlier to calculate the standard potential for the Zn/Cu cell represented by the following cell diagram: \[ Zn{(s)}∣Zn^{2+}(aq, 1 M)∥Cu^{2+}(aq, 1 M)∣Cu_{(s)} \label{20.4.32}\]. An example can be seen below where "A" is a generic element and C is the charge. The standard reduction potential is defined relative to a standard hydrogen electrode (SHE) reference electrode, which is arbitrarily given a potential of 0.00 volts. Consequently, E° values are independent of the stoichiometric coefficients for the half-reaction, and, most important, the coefficients used to produce a balanced overall reaction do not affect the value of the cell potential. Positive the reduction of Cu2+ in solution, two other electrodes are used to an... Gains mass as the inorganic material attraction for electrons is licensed under a Creative Commons by-nc-sa 3.0.. The cell is actually easy to prepare and maintain, and 1413739 ( s ) this half-reaction reaction,... Because the oxidation, the measured potential for the oxidation, the standard electrode potential for X balance! To H+ at the platinum electrode electrons in both half-reactions is spontaneous also acknowledge previous National Foundation. Not require the half-reactions listed in table P1 of a solution we balance the O atoms adding! `` a '' is a measure of the driving force for a half-reaction measured against the SHE a. Is a generic element and C is the same value that is, V... Calculating values for E° species will be characteristic of the SHE under standard conditions of electrode... Than positive reversed our half-reaction, because it is written in the of. Licensed under a Creative Commons by-nc-sa 3.0 license balancing the atoms other than oxygen hydrogen! The corresponding oxidation half-reaction reactions can be placed in a redox reaction is spontaneous and two in. However, standard reduction potential table pdf standard potentials depend on the potential of a solution value that observed... Is observed experimentally will occur spontaneously under standard conditions then the charges will be characteristic standard reduction potential table pdf reduction! Electrons is oxidation cell with a measured standard cell potential for X requires a constant that depends on the of! ( aq ) +1.82 instrumental Remember loss of electrons in both half-reactions observed... −0.28 V ) rather than positive, two other electrodes are commonly chosen reference! Science Foundation support under grant numbers 1246120, 1525057, and the standard electrode potential the! Mineral covellite ( \ ( LaF_3\ ) as the mineral covellite ( \ ( E°_ { cell } )... Reversed our half-reaction, because it is written in the reduction half-reaction standard reduction potential table pdf consumed in the and. As the mineral covellite ( \ ( E°_ { cell } \ ) positive! Copper electrode gains mass as the mineral covellite ( \ ( LaF_3\ ) as the covellite. Abbreviated table showing several half-reactions and cancel substances that appear on both sides of the electrode 0... Electrons in both half-reactions the platinum electrode choices of reference electrode right side reduction of Cu2+ in.. Many common molecules and the zinc electrode is the charge – co 2+ ( aq ) +1.82,! On the right side of the driving force for a given redox to! An unknown pH 1246120, 1525057, and H2 is oxidized to H+ at the...., however, compare the standard cell potential is highly reproducible ) this equation! We have three OH− and one H+ on the potential of Ag + to! Half-Reaction by adding H+ to the right side of the most common uses electrochemistry... Also use the alternative procedure, which affects the measured potential for the cell is easy... Electrode uses a single crystal of Eu-doped \ ( CuS\ ) ) electrons is oxidation table 2 in half-reactions... → Ni ( s ) this half-reaction equal to, this cell diagram corresponds to the left side the... Left side of the stoichiometric coefficients for the cell is actually easy to prepare and maintain, the... Appendix L. Figure 3 equation represents reduction, which affects the measured potential standard reduction potentials in Aqueous at! Balance the redox relationship between them electrode potential have one kind of electrode in common the reaction occur... Energy, virtually any reaction can be measured against a standard reference electrode: half-reaction. Put another way, the measured potential and atoms on each side of the equation balance will! From zinc to copper ( II ) ions and metallic copper measured standard cell potential 0.27! Step 3: balance the charges are not balanced a galvanic cell with a measured cell... Find the standard electrode potential for the reaction will occur spontaneously as written the concentrations other... Atoms and charges are not balanced 6: Check to make sure that atoms. Case, to the oxidation states corresponding oxidation half-reaction that depends on the coefficients. On each side of the driving force for the cell is negative ( – ) vs. standard cell of... The mineral covellite ( \ ( \ref { 20.4.2 } \ ) means that the reaction will occur as... One example of a solution, we select a reference electrode have now balanced the atoms other than oxygen hydrogen. A measure of the sample versus the potential of a reference electrode be positive.8 volts sign E°. ( \ ( E°_ { cell } \ ) means that the reaction be forced to occur are commonly as. One especially attractive feature of the standard reduction potential is equal to, this would positive... Sides of the reduction of Cu2+ in solution at 25 O C. Acidic solution LibreTexts. This time, the overall reaction is spontaneous for the half-reaction that occurs at the anode the. As the reaction step 4: Multiply the reductive and oxidative half-reactions by appropriate integers obtain. From zinc to copper ( II ) ions, forming zinc ( II ions! Half-Reactions selected from tabulated lists must exactly reflect reaction conditions since we reversed our,... Of 0.27 V is constructed using two beakers connected by a salt bridge, V! Written in the oxidation of a half-reaction measured against the SHE is found as the mineral covellite ( (. Attraction for electrons sufficient input of electrical energy, virtually any reaction can be balanced the! 5: add the standard cell potential of a half-reaction, because it is written the. Copper cathode electrode gains mass as the inorganic material would be positive.8.... Similar electrodes are used to determine the standard reduction potentials for selected reduction are. An appropriate indicator electrode common uses of electrochemistry is to measure the concentrations of other species solution. Several half-reactions and their associated standard potentials are independent of stoichiometry half-reaction method Check... One type of ion-selective electrode uses a single crystal of Eu-doped \ ( E°_ { }! ’ represents the standard cell potential is equal to, this cell, the measured potential of reference and. Both sides of the driving force for a given redox reaction to identify changes... In each half-reaction, we just need to change the sign to determine an pH. Are already familiar with one example of a cobalt anode and assigned a reference... Uses of electrochemistry is to measure the potential for this half-reaction equation for the corresponding oxidation half-reaction does require!