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JOURNAL OF CHEMICAL EDUCATION
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A N EXPERIMENT FOR THE pH METER ROBERT L. PECSOK University of California, Los Angeles, California
Tm
present trend for course material in quantitative analysis is to reduce the number of time-consuming, classical procedures to a minimum, retaining only a sufficient number to introduce the student to careful techniques. The increasing scope of analytical chemistry and the increasing use of analytical instruments in industrial laboratories require us to reconsider carefully the content of our courses. The introduction of instrumental methods to sophomore courses has not found universal acceptance; however, electrolytic determinations have been standard for many years and are discussed in considerable detail in the popular textbooks. Colorimetric determinations are now a part of most elementary courses; the instrument may be as simple as Nessler tubes, but more probably a Duboscq or photoelectric colorimeter. The basic principles of a pH meter are covered in the theoretical treatment of oxidation-reduction chemistry of most courses. In view of the importance and frequency of the determination of hydrogen ion concentration we feel that the use of the pH meter has a definite place in elementary courses. The cost of such a program is not unreasonable. Our investment in pH meters is less than that in platinum electrodes for electrogravimetric work. Although acidimetry is part of our first semester's work we postpone.the pH meter until the second semester, partly because of the fewer students involved, and partly because there are several experiments which must be staggered due to the apparatus involved. In this manner we are able to keep our equipment in use nearly constantly, and six meters are ample for a class of fifty. Few of the standard texts discuss the pH meter, and almost none give detailed experimental procedures.
Since inevitably some students must do the laboratory work before it can be discussed in lecture, they are provided with a summary of the principles involved in the constmction and operation of the meter. This includes a discussion of the electrode system with equations for the potentiats of both electrodes. It is shown that the equation relating potential to pH is linear and that therefore a single scale on a calibrated and properly standardized voltmeter can be read either in volts or directly in pH units. The amplifier and voltmeter are treated as a unit with the details of the circuits omitted. The advantages of a pH meter are stressed, particularly its accuracy, convenience, and adaptability to applications involving recording and control. In the selection of experimental work, we imposed these requirements: (1) experimental data that could be compared to theoretical data calculated by principles discussed in class, (2) determinations that would be more difficult by other methods, and (3) an unknown, the results of which could he graded. The titration of a polyprotic acid and its salts fulfills these conditions. The procedures which follow were designed for use with a Beckman Model H-2 line operated pH meter. They are taken in part from the instructions furnished by the manufacturer.' With minor modifications in the standardization and reading, they can be used for any of the commercial instruments available. They are obviously designed to keep breakage and other troubles at a minimum. In four semesters' experience with about 175 students our only maintenance has been replacing potassium chloride in calomel electrodes and one broken glass electrode. More rugged electrodes have recently become available. Bulletin 190-B, Beckman Instruments, Inc., South Pasadena, California.
MAY, 1951 PROCEDURES
Care of Electrodes. The electrodes are fragile and easily mined if scratched or bumped. Never touch the glass parts of the electrodes with anything. The electrodes are always stored in distilled water. When transporting the instrument, raise the electrodes and turn them toward the meter. Procedure for Standardization. 1. 2. 3. 4. 5.
Immerse the electrodes in a beaker of distilled water. Set range switch to "start!' Connect power line plug t o 110-V. a,-c. outlet. Allow a five-minute warm-up period. During warm-up period prepare standard buffer by either of the following methods. (a) Dilute 5 ml. of Beckman concentrated buffersolution to 125 ml. using a volumetric flask. Note the temperature of the buffer, and its comct pH from the bottle label. ( b ) Prepare a mturated solution of potassium acid tartrate (solubility-4.6 g./100 ml.) The pH of this solution is 3.57 with a negligible temperature coefficient.= Meter needle should read exactly 7.00 OH. If not.. adjust . with small screw above needlepivot.Set "temp" control to temperature of buffer. Set range control to "neut" position. The rsnge mviteh should always be snapped smartly from one position to another and not stopped at an intermediate point. Rinse and fill beaker with buffer solution. Immerse electrodes. Set range switch to 0-8 pH range. Adjust the meter needle to read the pH of the buffer by turning the "~tmdardieatiou"control. Set m g e control to "neut" position. Set adjustable pointer to indicate position of meter needle. Pointer is controlled by small knob a t lower left of meter panel. The coincidence of the pointer and the needle with the range switch in "neut" position should be checked frequently as this indicates any drift of the amplifier. Adjust the meter needle if necessary with the "standardization" control. Always leave the range switch in the "neut" position between readings or when changing solutions, and whenever the electrodes are not immersed. Failure to observe this
Experimental Set-up for Titrations
4.
Arrange stirrer so that there is no danger of contact x-ith the electrodes. The solution should bc stirred continuously and the electrodes left immersed during the entire titration. I t may be necessary to ground the meter to prevent drifting of the needle due to electrical pick-up of the amplifier. 5. Mount the buret and commence the titration with mdium hydroxide solution. Measure and record the pH after each increment of barn. Take sufficient points so that an accurate titration curve can be plotted, e. g., every 3 to 5 ml. between end points and every 0.5 ml. or less near the end points. Do not continue the titration beyond pH 11 to 12. Above pH 11 significant corrections are necessary. At very low hydrogen ion concentrations the glass bseomea sensitive to sodium ion as well as hydrogen ion. 6. Plot the curve on graph paper. Plot the theoretical curve and compare, particularly, the pH at the equivalence points. Note the pH for the first two equivalence points of the experimental curve and use these values in subsequent titrations.
Standardization of Base.
absorbent tissue may be used.)
Measurement of pH. 1. With the range switch in the "neut" position, immerse electrode in sample. 2. Check coincidence of ~ o i n t e and r needle and adjust if neceswry. 3. Switch range coutml to proper pH range. 4. The needle now indicates the pH. 5. Return range switch to "neut" position before raising olcctrodes.
Titration Curwe of Phosphoric Acid. 1. Prepare a liter of 0.2 N sodium hydroxide (oarbonate-free). Pipet a 25-ml. partion of approximately 0.1 M phosphoric acid into a 250-ml. beaker. 3. Immerse electrodes in the acid, dilute with water if necessary to cover the bulb of the glass electrode (see the figure).
1. Obtain a sample of standard su1ful.i~acid from stockroom. 2. Pipet a 25-ml. portion of acid into tho beaker and titrate with base, following cbange of pH. The student should add a. drop or two of a suitable indicator snd follow its color change. Obtain a eufficient number oi points to plot a. titration ourve. Plot this curve on the same graph as the phosphoric seid curve and compare the two. 3. Calculate two normalities for tho basc, using the end points determined from the phosphoric acid curve. The normalities will differ depending on the amount of carbonate in the base.
Analysis of Mixed Phosphate Solutions. 1. Two separate samples are to be analyzed."ake a 25-ml. samnle and titrate to both end mints i ~ r o ~oeHr has been
2.
'LINGANE, J. J., Anal. Chem., 19, 810 (1947).
2.
Identify the oomponents of each unknown mixture and report the concentration of each.
a Various mixtures of H3POrNaHzP0, and Pu'aH2POrNasHP04 can be used.
