Gardeners, particularly those who grow hydrangeas, use more precise techniques in adjusting the ionic balance of solutions. They first determine the relative proportion of hydronium to hydroxyl ions in the moisture of the soil by means of a simple chemical test. If the test indicates that there are more hydronium than hydroxyl ions, the soil is acid; if the reverse is true, the soil is basic or alkaline. Pink hydrangeas that are grown in acid soil are likely to turn blue, and blue varieties that are grown in alkaline soil usually turn pink. Accordingly the gardener adds to the flower bed a carefully measured quantity of either aluminum sulfate, which introduces additional hydronium ions and therefore an acid condition, or hydrated lime, which produces hydroxyl ions and therefore a basic condition, depending on whether blue or pink hydrangeas are wanted.

The pigments in flower petals and green leaves are sensitive to the ratio of hydronium to hydroxyl ions in water, as can be demonstrated by a simple experiment. Crush the petals from a red rose to a fine pulp with a mortar and pestle. Add to the pulp about five milliliters of tap water and continue grinding for a minute or so. Filter the fluid through a sheet of paper toweling into a clean glass container. Discard the pulp. Similarly, filter the juice of a lemon into a clean glass container and discard the pulp.

Place five milliliters of tap water in a clean glass container and make a saturated solution of baking soda by adding soda until no more will dissolve. In another container of clean, clear glass (preferably a small test tube) place 10 milliliters of tap water and to it add 10 drops of the filtered extract from the rose petals. The extract will impart a pink tint to the tap water.

Now add five drops of the concentrated soda solution to the pink water. Swirl the test tube for a few seconds to mix the fluids. The color of the mixture will soon turn light blue. Add 10 drops of filtered lemon juice to the mixture and swirl the tube as before. The pink tint will reappear.

The color can be cycled from pink to blue and back again many times by alternately increasing the concentration of hydronium ions with lemon juice and decreasing it with soda solution. Deeper colors can be developed by using stronger reagents. The addition of a single drop of concentrated lye solution (sodium hydroxide) will change the pink color to a rich brown. Two drops of hydrochloric acid will restore the pink hue.

A tap-water solution of pigment extracted from the petals of blue hydrangeas turns green when it is made acid and yellow when it is made basic. The pigment of most green leaves be clear in acid and yellow in basic solutions. Substances that react in this way are known to chemists as "indicators." They are frequently used to determine the approximate proportions of hydronium to hydroxyl ions in solutions. Sam Epstein, chief chemist of the Federated Metals Division plant of the American Smelting and Refining Company in Los Angeles, explains how to use indicators for measuring the acid-base balance of solutions, why they change color and how to make an instrument for determining the relative concentrations of hydronium to hydroxyl ions at which they change color. Epstein writes:

"Part of the explanation is found in the nature of water. The chemical formula for a molecule of this unique liquid can be written in either of two forms, H₂O or HOH. I prefer the H₂O formula because it suggests that the structure of a molecule of water is not symmetrical. The hydrogen atoms (H) are grouped on one side of the oxygen atom (O), not on precisely opposite sides.

"For this reason the molecule is also electrically asymmetrical. When water is placed between a pair of charged electrodes, the oxygen sides of its molecules turn toward the positive electrode and the hydrogen sides turn toward the negative electrode. The molecule is said to have an electric dipole moment.

"At room temperature the water molecules of a solution dart about randomly and knock into one another continuously. Occasionally an impact is so violent that a molecule splits into a pair of fragments. One fragment consists of the nucleus of a hydrogen atom (a proton). The other is a clump of particles consisting of the remaining hydrogen atom, the oxygen atom and the electron formerly associated with the dislodged proton. Both fragments now carry an electric charge. One fragment is H + (because the proton lost its electron) and the other one is OH- (because the atoms of this fragment gained the unit charge carried by the acquired electron). "The force of electric attraction between the oppositely charged fragments causes some pairs to reunite promptly. Occasionally, however, a proton is knocked a substantial distance by the impact and wanders briefly among neighboring molecules. The positive charge carried by the wandering H+ ion soon responds to the negative side of a water molecule (H₂O) and unites with the molecule to form a new ion: H₃O+. This particle is known as a hydronium ion.

'Water molecules, because of their electrical character, also tend to form clumps. The weakly positive side of one molecule attracts the weakly negative side of a neighbor. For this reason the H+ ion may attach itself to a group, forming H₅O₂+, H₇O₃+, H₉O₄+ and so on. All can be considered hydronium ions, but hereafter hydronium will refer to the H₃O+ group.

"The relatively massive remaining fragment, the OH- or hydroxyl ion, tends to exist alone in solution because the polar forces of the molecules are too weak to bind it. In a liter of water at room temperature about 6 x 10¹⁶ molecules are ionized. As a result a liter of 'pure' water is actually a solution containing 6 x 10¹⁶ hydronium ions and the same number of hydroxyl ions. As large as this number is, it is only .0000002 percent of the number of molecules in a liter of water. This is a seemingly trivial percentage, yet it accounts for much of the chemical activity of water.

"Measuring the concentration of either ion in an aqueous solution automatically gives the concentration of the other one. Since a count of the electrified particles would result in an awkwardly large number, a convention has been adopted for expressing the hydronium-ion concentration of water solutions by means of an inverse logarithmic scale that ranges from 0 to 14 in units called pH. To make use of this scale the hydronium-ion eoncentration must first be known in terms of 'molarity.' One liter of a solution that is one 'molar' with respect to H₃O+ contains one mole of H₃O+, an amount whose weight in grams is-equal to the sum of the atomic weights of the atoms in the chemical formula, that is, 19 grams of H₃O + (3 x 1 + 16 = 19). The pH is defined as the logarithm of the reciprocal of hydronium-ion concentration. Pure water, or any neutral aqueous solution, contains 10-7 mole of H3O + per liter, therefore its pH is 7 (logarithm of 1/10^-7 = logarithm of 10⁷ = 7). A pH of less than 7 indicates that a solution contains an excess of hydronium ions and a correspondingly smaller amount of hydroxyl ions. Such solutions taste sour or acid and many of them are toxic. The pH of solutions that are deficient in hydronium ions and therefore contain an equivalent amount of excess hydroxyl ions is greater than 7. They taste bitter and feel soapy. Many basic solutions are also toxic.

"The pH of solutions can be measured most easily and simply by the use of organic substances that resemble the pigments of flowers and act chemically like such pigments. They impart one color to a solution of given pH, change to a different hue at a higher pH and serve as indicators. Many of them are highly purified synthetic compounds that are designed to change color within a predetermined range of pH. Two common ones are phenolphthalein, which changes from colorless to red in the pH range between 8.5 and 10, and methyl orange, which undergoes a change from red to yellow in the pH range between 3 and 4.

"Indicators are also available in the form of impregnated strips of paper that are particularly convenient to use. One product, known as 'pHydrion' paper, is stained by a series of indicators that assume a characteristic color at each unit of pH from 1 to 11. After the paper strip has been dipped in the test solution the hue is estimated to the nearest pH unit by comparing the paper with a color chart supplied with each roll. This is an easy way to measure the pH of water in a swimming pool. Gardeners will find pHydrion paper equally convenient for measuring the pH of soil. To test soil mix a sample with a comparable amount of tap water, let most of the dirt settle and check the liquid with the paper.

"The pH range over which there are changes in the color of homemade indicators, such as those extracted from flower petals, leaves, red cabbage and the colored juices of berries, can be determined with an electrical pH meter that can be made at home. Building a meter and doing the experiments needed to calibrate it will provide an interesting introduction to the properties of acidic and basic solutions.

"The pH meter measures the voltage developed between two electrodes that are immersed in the solution being tested. In essence the electrodes and the solution constitute a voltaic cell that works much like an ordinary flashlight battery. The potential between the solution and one electrode, which is called the indicator electrode, varies with the concentration of hydronium ions in the solution. The potential between the solution and the other electrode, which is called the reference electrode, remains constant.

"The output of the cell is compared with a known voltage by means of a potentiometer. The pH of the test solution is determined by reference to a graph on which pH is plotted against voltage. The instrument can measure a hydronium-ion concentration accurately to a fraction of a pH unit.

"The apparatus can use indicator electrodes of two types, one of platinum wire and the other of antimony [see Figure 2]. When the platinum electrode is used, a small amount of quinhydrone is added to the solution being tested. For this reason the platinum-wire electrode is known as a quinhydrone electrode. The potential between the solution and the platinum-wire electrode varies directly with the hydroniumion concentration, but the measurements are not reliable above pH 8. The antimony electrode works well in both acid and basic solutions, but it is more difficult to make.

"To make the electrodes, heat to redness a half-inch zone in the middle of a 12-inch length of 10-millimeter, soft glass tubing. The heat can be provided by a liquid-propane torch of the kind sold in hardware stores. Grasp the glass near the ends by both hands, pass the center of it through the flame a few times for gradual preheating and then rotate the glass directly in the flame until the midsection softens uniformly.

"Remove the glass from the fire, stretch the midsection, let the glass cool, nick the center of the constricted portion with the corner of a file and pull the ends apart. The tubing will break squarely where the glass was nicked. The finished pieces should have the form shown in the illustration.

"For the quinhydrone electrode slide a length of platinum wire into the tapered end of one tube so that half of the wire extends beyond the glass. Return the tip of the glass, with the wire, to the flame and apply heat until the glass melts uniformly around the wire. Rotate the assembly as required to prevent the glass from drooping. Remove the assembly from the flame and, while the glass is soft, push the exposed part of the wire sideways as necessary to align the wire with the axis of the tubing.

"Return the sealed end to the edge of the flame and, with the pointed end downward, feed enough rosin-core solder into the tube to immerse the inner end of the wire. Complete the electrode by pushing the cleaned end of a flexible, insulated copper wire into the molten solder. Remove the electrode from the flame and let it cool in the upright position.

"To make the antimony electrode, melt the constricted tip of the second glass tube until it closes and fill the tube to a depth of about two inches with small lumps (not powder) of 99.8 percent pure antimony metal. Antimony of this grade can be bought in lots of 1/4 pound for about $2 from the Fisher Scientific Company, 711 Forbes Avenue, Pittsburgh, Pa. 15219. Heat the tube until the antimony melts. Shake the tube lightly as necessary to prevent bubbles of air from becoming trapped in the metal.

"Let the metal cool until it solidifies. Then add a layer of rosin-core solder and insert a copper lead wire into the solder. The solder should fuse to the antimony as well as to the copper. When the assembly cools, nick and break off the tip of the glass to expose a small cylinder of antimony, which will act as the electrode. Polish the exposed metal with crocus cloth.

"The reference electrode, which is also known as the saturated calomel electrode, must be connected to the test solution through two tubes of agar moistened with potassium chloride solution [see Figure 1]. The tubes are known as salt bridges. The complete electrode assembly consists of 40-milliliter centrifuge tubes or heavy-walled test tubes, U loops of quarter-inch glass tubing and also the associated chemicals and hardware. The mercury, calomel paste, solid potassium chloride and saturated potassium chloride solution are prepared and placed in sequence in one of the centrifuge tubes. The relative proportions are not critical and can be judged from the illustration.

"Care must be taken, however, in preparing the agar-potassium chloride bridges, which protect the saturated calomel electrode from contamination by the test solutions. To prepare the bridges soak four grams of agar in a small beaker containing 100 milliliters of distilled water. Preferably the soaking should continue overnight.

"Place the small beaker in a larger one of boiling water and apply heat until the agar is fully dissolved. Add 30 grams of potassium chloride and, with the beaker still in the boiling water, stir the mixture until the potassium chloride has dissolved. If necessary, add just enough water to dissolve the salt completely.

"Invert the U tubes, fit the ends with short sleeves of rubber tubing, clamp the tubing and fill the assembly with the agar solution. The solution must completely fill the bridge tubes and the flexible sleeves. If it does not, air bubbles may be trapped when the tubes are inverted in the solutions.

"In spite of the best care the agar eventually becomes contaminated, producing erratic results. To replace it disassemble the apparatus, put the bridges in boiling water to melt the agar, rinse them thoroughly and refill them. The agar tube that dips into the test solution should be immersed in saturated potassium chloride solution when the calomel electrode is not in use.

"The voltage developed across the electrodes by reaction with the test solution must be measured by a potentiometer of the null balancing type, which is an instrument that draws no electric current during the interval when the unknown voltage is measured. In instruments of this type voltages of opposing polarity, one known and the other unknown, are simultaneously applied to a meter [see Figure 5]. The magnitude of the known voltage is adjusted until the meter reads zero. At this point it equals the unknown voltage. The instrument must be calibrated each time it is used.

"To make the calibration for measuring potentials up to one volt, plug in the leads from the dry cell and turn the range switch to the one-volt position. Rotate the dial of the precision potentiometer to the 135 position. Move the 'measure-calibrate' switch to the calibrate position. Set the knob of the high-current potentiometer to the center of its travel. Depress the 'coarse-read' button momentarily and note the meter deflection. Turn the knob of the high-current potentiometer slightly in either direction, depress the button again and note the response. If the pointer of the meter moves more violently in the same direction, readjust the high-current potentiometer in the other direction and depress the button again.

"The object is to adjust the high-current potentiometer to a position that causes a meter deflection of less than one scale division. The purpose of the coarse button is to protect the meter from possible damage by excessive current when the instrument is grossly unbalanced. Complete the calibration by alternately depressing the fine button and adjusting the high-current potentiometer until the meter deflection is negligible. When the instrument is so balanced, the potential of unknown voltage that is developed by the test solution is determined by dividing the number on the dial of the precision potentiometer by 1,000. For example, a dial reading of 250 indicates a potential of .250 volt (250/1,000 = .250).

"The instrument is similarly calibrated to measure the potential range from 0 to .2 volt (0 to 200 millivolts). To calibrate this range turn the range switch to the .2-volt position and the dial of the potentiometer to 675. Balance the instrument by alternately operating the coarse and the fine button and adjusting the low-current potentiometer until the meter shows negligible deflection. The dial readings when now multiplied by .0002 indicate the potential. For example, a dial reading of 250 indicates a potential of .05 volt (250 X .0002 = .05). The calibration is valid for only one range at a time. Calibration of the .2-volt range destroys the calibration of the one-volt range and vice versa.

"The electrode system must also be calibrated. Prepare with distilled water a solution of 1,000 milliliters that contains 28.4 grams of reagent-grade anhydrous disodium phosphate and another solution of the same volume that contains 21 grams of reagent-grade citric acid. Weigh the chemicals directly from freshly opened bottles and store the solutions in clean plastic containers.

"Put 39.2 milliliters of the citric acid solution and .8 milliliter of the phosphate solution, which is basic, in a clean beaker. The solutions can be transferred conveniently by using burettes, which are calibrated glass tubes that can discharge solution one drop at a time, or by means of a calibrated pipette. The pH of the mixture is 2.2.

"Immerse the antimony electrode and the reference electrode in this mixture. Connect the antimony electrode to the negative terminal of the potentiometer and the reference electrode to the positive terminal. Mix the solution in an electrically operated agitator for a few minutes to saturate it with air. Measure the voltage developed by the test solution on the one-volt range. When making the test, set the measure-calibrate switch in the 'measure' position. If the potentiometer fails to balance, reverse the leads to the electrodes. Record the voltage and also the corresponding pH

"Next, replace the antimony electrode by the platinum electrode and connect it to the positive terminal of the potentiometer. Connect the reference electrode to the negative terminal. Dissolve as much solid quinhydrone in 10 milliliters of rubbing alcohol as the alcohol wili hold, thus making a saturated solution. This chemical can be obtained from the Fisher Scientific Company in lots of 1/4 pound for about $4. Add five drops to the 2.2-pH test solution. While stirring the mixture, measure and record the voltage.

"Repeat this procedure, alternately using the antimony and quinhydrone electrodes, to determine the voltage developed by solutions that contain the following volumes (in milliliters) of citric acid and sodium phosphate: 24.6 and 15.4, 14.7 and 25.3, 1.1 and 38.9. The pH of the three mixtures is respectively 4, 6 and 8.

"The polarity of the quinhydrone electrode reverses at a pH of approximately 7.6. For this reason the leads from the electrodes to the potentiometer must be reversed when the quinhydrone electrode is used for measuring solutions of concentration higher than 7.6 and the potentiometer must be operated on the .2-volt range. Indeed, accuracy is improved by using this range for measuring all potentials of less than .1 volt.

"Plot the recorded voltages and the corresponding pH in the form of two calibration graphs [see Figure 6]. Because of slight irregularities, including errors of observation, the plotted points may tend to scatter slightly on both sides of a straight line. Draw the graphs through the center of the scattered group of points so that equal numbers of points fall on each side of the line. Any point that is obviously out of line with the others should be remeasured.

"The range of pH through which a unknown indicator changes color ca now be determined easily. For example prepare an extract of pigment from rose petals or other flowers. Add a few drop of the extract to 35 milliliters of distilled water, enough to tint the water a viewed against a sheet of white paper. Insert the antimony and reference electrodes in the solution.

"Pour into a clean container a few milliliters of the citric acid solution use during the calibration, into another clean container pour a like amount of phosphate solution. Add the phosphate solution drop by drop to the test solution until the pigment just changes color. Stir the mixture continuously. Measure the pH. Add citric acid drop by drop until the color changes again. Measure the pH;

"These measurements establish the limits of the pH range through which the indicator is effective. The range of hues that indicate intermediate values of pH can be determined by beginning at either extreme and measuring the voltage of the system after the addition of each increment of acid or phosphate solution. The corresponding pH values are obtained from the antimony-electrode graph.

"Why do indicators change color when pH changes? Most indicators are weak acids-substances that can donate hydronium ions to solutions. The molecules of the acid display a characteristic color when the acid is placed in a solution whose pH is below a certain level, which will depend on the particular indicator involved. In solutions of higher pH the molecule ionizes, or splits. The ionized fragments have a different color. Thus it is possible to switch the color of an indicator back and forth a number of times by changing the pH of the solution

"The potentiometer can be used to check the pH of aqueous solutions such as vinegar, baking soda, borax, Epsom salts, carbonated beverages, household ammonia and fruit juices. The instrument is also useful for measuring the pH of turbid or colored solutions that obscure the true color of an indicator.

"Warning: Many chemical compounds are toxic, particularly those containing mercury. If chemicals come in contact with the skin, wash the affected area promptly, preferably in running water. When making experiments, wash your hands frequently and keep them away from your mouth."

Bibliography

STANDARD METHODS OF CHEMICAL ANALYSIS. Wilfred W. Scott. D. Van Nostrand Company, Inc., 1939.

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