Chem 321 Lecture 16 - Potentiometry

Chem 321 Lecture 16 - Potentiometry In lab you will use an ion-selective electrode to determine the amount of fluoride in an unknown solution. In this approach, as with pH measurements, the potential of an electrochemical cell is associated with the concentration (actually activity) of a particular ion in the solution being analyzed. In order to properly interpret the chemistry occurring in electrochemical cells, it is usually necessary to be able to assign oxidation numbers and to write balanced redox equations. These concepts are reviewed below. Assigning Oxidation Numbers Oxidation is an increase in oxidation number while reduction is a decrease in oxidation number. In order to recognize these processes, one must be able to assign proper oxidation numbers to the various elements associated with a particular chemical reaction. Remember that oxidation number, or oxidation state, is a bookkeeping device used to track electron transfer. The following guidelines will assist you in assigning oxidation numbers. chemical species oxidation number elemental form (e.g., Cu or Br2) zero monatomic ion (e.g., Na+ or S2-) charge on monatomic ion chemically combined hydrogen +1 (notable exception is in a metal hydride such as NaH where ox. no. = -1) chemically combined oxygen -2 (notable exception is in a peroxide such as H2O2 where ox. no. = -1) The sum of the oxidation numbers of all atoms in a given chemical species is equal to the net charge on the species. Guidelines for Balancing Redox Equations The method presented here is similar to the half-reaction method that is described in general chemistry textbooks. It is easy to apply to redox reactions which take place in aqueous solution, however, one must be careful to follow the sequence of steps exactly. 1. Assign oxidation numbers to all atoms in the equation and determine which substance is being oxidized and which is being reduced (from changes in oxidation numbers). Write separate half-reactions for each process. 2. Balance each half-reaction in the following way. a) By using appropriate coefficients, balance all atoms other than hydrogen and oxygen in the basic half-reaction. If hydrogen and/or oxygen are being oxidized or reduced, these atoms should be balanced as well. Keep in mind that atoms, except for hydrogen and oxygen as noted below, may appear in the half-reaction only in the form of substances indicated in the original equation. b) Add the necessary number of electrons (to account for the change in oxidation number) to the appropriate side of the equation; to the reactants' side if it is a reduction half-reaction, and to the products' side if it is an oxidation half-reaction. c) Balance charge by the addition of H+ ions if the reaction is in acidic solution, or OH- ions if the conditions are basic. d) Balance the hydrogen and oxygen atoms by the addition of water. 3. Combine the balanced half-reactions so that the electrons cancel exactly on each side of the final equation. This may necessitate multiplying one or both halfreactions by some integer to equalize the number of electrons involved in each half-reaction. 4. Cancel substances common to both sides of the equation. 5. Check to see if both atom and charge balance have been achieved. Electrochemical Cells A shorthand line notation for the description of an electrochemical cell is often used in place of detailed cell diagrams. The general convention is to indicate the cell anode on the left and the cell cathode on the right. They are usually separated by a pair of vertical parallel lines that signify a salt bridge. Anode 3 Cathode The anode is always the site of the oxidation half-reaction and the cathode is always the site of the reduction half-reaction. The salt bridge serves as an ion channel that permits charges to move from one electrode to another to maintain electrical neutrality in each half-cell while keeping the half-cell components separate. The line notation for each half-cell indicates the oxidized and reduced forms of the electroactive substance in the half-cell, along with any special reagent conditions. The substance that serves as the electrical contact between the half cell and the external circuitry is always placed at the extreme left and the extreme right positions in the line notation. Substances that are in contact and in different physical phases are separated by a single vertical line. To illustrate this, consider a cell with an anode that consists of a Ag wire coated with AgCl and immersed in a saturated KCl aqueous solution and a cathode consisting of a platinum wire in contact with an aqueous solution containing Fe2+ and Fe3+. The two electrodes are connected by a salt bridge. The line notation below describes this cell. Anode Ag(s) Cathode AgCl(s) KCl(sat’d) Fe2+(aq), Fe3+(aq) Pt(s) Note that because neither of the forms of the substance (Fe) that is being reduced in the cathode compartment is a suitable electrical contact for an external circuit a nonreactive electrical conductor like Pt is used for the half-cell contact. Also note that since both Fe2+ and Fe3+ are in the same phase (aqueous solution) they are separated by a comma and not a vertical line. This notation indicates that the anode involves the oxidation of Ag in the presence of chloride ion to form AgCl. The anode half reaction is: anode: Ag(s) + Cl-(aq) º AgCl(s) + e- In this cell, the reaction at the cathode involves the reduction of Fe3+ to Fe2+. cathode: Fe3+(aq) + e- º Fe2+(aq) In this particular electrochemical cell the anode half-cell is a reference electrode. A reference electrode is a half-cell with a constant, and usually known, half-cell potential. To see how this is achieved for a reference electrode such as the one above, recall the Nernst equation that describes the cell or half-cell potential. For reaction aA + bB º cC + dD, the potential (E) is given by where E0 is the standard cell (or half-cell) potential, R is the universal gas constant, T is the Kelvin temperature, n is the number of moles of electrons involved in the redox process, F is the Faraday and the log term involves the reaction quotient (Q) written using activities. At 25 EC the Nernst equation becomes: For the anode half-cell above, the Nernst equation is: Recall that the standard half-cell potential is the electrode potential when all reactants and products are at unit activity (a=1). To get a standard oxidation half-cell potential, reverse the reduction reaction and change the sign of the reduction potential. Substituting for E0 and n in the anode Nernst equation above gives: This indicates that the half-cell potential depends only on the chloride ion activity present in the electrode compartment. A constant half-cell potential is achieved by constructing the electrode with a large concentration of chloride ion. When a KClsaturated solution is used, the anode potential is -0.197 V. Recall that ∆G= -nFE. Thus, a negative potential is associated with a nonspontaneous process (∆G>0), and a positive potential is associated with a spontaneous process (∆G<0). A strong oxidizing agent (a substance that is easily reduced) will have a large, positive reduction half-cell potential (Ered). A strong reducing agent (a substance that is easily oxidized) will have a large, positive oxidation half-cell potential (Eox). Electrochemical cells, such as the one used in a pH meter, are often connected so that very little current flows in the cell. Consequently, very little reaction, which could alter the composition of the half-cell, occurs while the cell is being monitored. In a similar way, the Nernst equation for the cathode is: The overall cell potential (Ecell) is the sum of the half-cell potentials. Ecell = Eox + Ered For the cell above, Since the potential of the reference electrode (Eref) is fixed, the overall cell potential depends only on the relative Fe2+ and Fe3+ activities in the anode compartment. One analytical application of direct potentiometric measurements is in a potentiometric titration. In such an approach you construct an electrochemical cell that uses the titration solution as one of its electrodes. This is illustrated in the diagram below. In this case, the chloride in a dissolved sample is being titrated with a AgNO3 solution. As the titration proceeds, the concentration (and activity) of Cl- in solution decreases and this causes a change in the cell potential. By continuously monitoring the cell potential during the titration, it is possible to determine the equivalence point in the titration. In the cell shown in Figure 11.1, the titration solution containing chloride is serving as the cathode. Figure 11.1 Diagram of an electrochemical cell for the potentiometric titration of chloride The potentiometric chloride titration reaction is: Cl-(aq) + Ag+(aq) 6 AgCl(s) As the titration proceeds, the concentration and activity of chloride ion in the titration solution decrease, which in turn causes Ered and Ecell to increase (remember that the anode is a reference electrode with a constant half-cell potential). Initially the potential changes only slightly with addition of titrant. However, around the equivalence point the cell potential increases significantly for even small additions of titrant, resulting in a characteristic S-shaped titration curve similar to that found for acid-base titrations (see Fig. 11.2). Figure 11.2 Characteristic curve for a potentiometric chloride titration The equivalence point in a potentiometric chloride titration is at the inflection point in the steep portion of the curve and can be determined by various graphical methods. One can plot the rate of change in the potential per added titrant volume (∆Ecell/∆V) vs. VAgNO3 (the average titrant volume) and use the peak center of this pseudo-derivative plot as the equivalence point in this titration (see Fig. 11.3). Figure 11.3 Pseudo-derivative plot for a potentiometric chloride titration Potentiometric titrations such as this also provide a means to measure equilibrium constants. In this case, the Ksp for AgCl can be determined. At the equivalence point of the titration, equal amounts of silver ion and chloride ion must be present in the solution (i.e., [Ag+] = [Cl-]). Since the hydrated ion size is nearly the same for Ag+ and Cl-, the activity coefficients of these ions are essentially equal and aAg+ = a Cl- at the equivalence point. The solubility product expression at the equivalence point then becomes: Ksp = (aAg+)(aCl-) = (aCl-)2 The value of aCl- at the equivalence point can be evaluated from if the cell potential at the equivalence point and the reference half-cell potential are known. The value of Eref can be determined by measuring Ecell under the conditions where the cathode half-cell contains a solution of known potassium chloride concentration and known ionic strength. The chloride ion activity in the equation above is calculated from aCl- = [Cl-] γCl-, and the equation above is solved for Eref.