pH Buffer Solution Preparation

Preparing Buffer Solutions for pH meter Calibration

Buffer solutions are used for the calibration of pH meters as described in the post entitled "Calibrating a pH meter using buffers" . In most pH measurements, a single glass electrode-reference electrode probe assembly is transferred between two solutions. E_x is the measured potential of the solution tested and E_s is the measured potential of a standard pH buffer solution.

By definition the pH of a solution x is given by:

pH_x = pH_s +( E_x - E_s)/k

where pH_s is the pH of the standard pH buffer solution, pH_x the pH of the tested solution, E_s and E_x the corresponding potentials and k is a constant at a specific temperature.

The instrument if it is properly calibrated must be capable of reproducing pH values to 0.02 pH units.

To prepare the standard pH buffer solutions (recommended by NIST), the indicated weights of the followig pure materials should be dissolved in water of specific conductivity not greater than 5 micromhos:

• Potassium Tetraoxalate (0.05M) - Dissolve 12.61 g of KH₃(C₂O₄)₂.2H₂O in water to make 1000 ml
• Potassium Biphthalate (0.05M) - Dissolve 10.12 g of KHC₈H₄O₄ previously dried at 110 C for 1 hour, in water to make 1000 ml
• Equimolar Phosphate (0.05M) - Dissolve 3.53 g of Na₂HPO₄ and 3.39 g of KH₂PO₄, each previously dried at 120 C for 2 hours, in water to make 1000 ml
• Sodium Tetraborate (0.01M) -Dissolve 3.80 g of Na₂B₄O₇.10H₂O in water to make 1000 ml. Protect from atmospheric carbon dioxide.
• Calcium Hydroxide, saturated at 25 C - Dissolve an excess of calcium hydroxide with water, and decant at 25 C before use. Protect from atmospheric carbon dioxide.

The above solutions must be stored in hard glass or polyethylene bottles fitted with a tight closure or a carbon dioxide-absorbing tube (soda lime). Fresh solutions should be prepared at intervals not to exceed 3 months using carbon dioxide-free water.

Carbon dioxide-free water is distilled water that has been boiled vigorously for at least 5 minutes and allowed to cool without contact with the atmosphere.

Table 1 indicates the pH of the buffer solutions as a function of temperature.

Table 1: pH Values of Buffer Solutions for Standardization

Temperature(°C)	Potassium Tetraoxalate 0.05 M	Potassium Biphthalate 0.05 M	Equimolar Phosphate 0.05M	Sodium Tetraborate 0.01M	Calcium Hydroxide saturd. 25 °C
0	1.666	4.003	6.984	9.464	13.423
5	1.668	3.999	6.951	9.395	13.207
10	1.670	3.998	6.923	9.332	13.003
15	1.672	3.999	6.900	9.276	12.810
20	1.675	4.002	6.881	9.225	12.627
25	1.679	4.008	6.865	9.180	12.454
30	1.683	4.015	6.853	9.139	12.289
35	1.688	4.024	6.844	9.102	12.133
38	1.691	4.030	6.840	9.081	12.043
40	1.694	4.035	6.838	9.068	11.984
45	1.700	4.047	6.834	9.038	11.841
50	1.707	4.060	6.833	9.011	11.705
55	1.715	4.075	6.834	8.985	11.574
60	1.723	4.091	6.836	8.962	11.449
70	1.743	4.126	6.845	8.921
80	1.766	4.164	6.859	8.885
90	1.792	4.205	6.877	8.850
95	1.806	4.227	6.886	8.833

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Relevant Posts

Measuring the pH of a Solution with a pH meter

Calibrating a pH meter using buffers

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References

1. CRC Handbook of Chemistry and Physics, 52nd edition, The Chemical Rubber Co., (1971)
2. U.S. Pharmacopeia, USP 37
3. R. G. Bates, J. Res. Natl. Bur. Stand. (U.S.), 66A:179 (1962)
4. David W. Oxtoby, H.P. Gillis, Alan Campion, “Principles of Modern Chemistry”, Sixth Edition, Thomson Brooks/Cole, 2008
5. Steven S. Zumdahl, “Chemical Principles” 6th Edition, Houghton Mifflin Company, 2009

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Key Terms

pH buffer solutions, calibrating a pH meter, buffers, buffer solutions,

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Measuring the pH of a Solution with a pH meter

pH meter: Measuring the pH of a solution

The pH of a solution can be measured quickly and accurately with a pH meter (see Figure 1).

How does a pH meter work?

A pH meter has to somehow measure the concentration of the hydrogen ions [H⁺] in a solution. An acidic solution has far more positively charged hydrogen ions in it than an alkaline solution, so it has greater potential to produce an electric current under certain conditions - in other words, it is like a battery that can produce a greater voltage. A pH meter takes advantage of this and works like a typical voltameter: in brief, a pH meter consists of a pair of electrodes connected to a meter capable of measuring small voltages, on the order of millivolts. It measures the voltage (electrical potential) produced by the solution whose acidity we are interested in, compares it with the voltage of a known standard solution, and uses the difference in voltage (the potential difference) between them to calculate the difference in pH.

What are the parts of a pH meter?

A typical pH meter consists of two parts: i) one special measuring probe (a glass electrode) or two measuring probes that are inserted into the solution whose pH is required and ii) an electronic meter that measures and displays the pH reading. A glass electrode is in a sense two electrodes combined in one. It consists of a long glass tube with a thin walled glass bulb at the end. Special glass of high electrical conductance and low melting point is used for the purpose. This glass can specifically sense hydrogen ions H⁺ up to a pH ≈ 9 (with special glass electrodes pH ranges from 1-13 can be measured). The bulb contains 0.1 M HCl and a Ag/AgCl electrode (used as an internal reference electrode) is immersed into the solution and connected by a platinum wire for electrical conduct.

The main advantages of the glass electrode are:

• It can be used in the presence of strong oxidizing or reducing substances and metal ions
• Accurate results are obtained in the range pH 1-9. However, by using special glass electrodes pH 1-13 can be measured
• It is simple to operated. It can be attached to portable instruments and is used quite often in chemical, biological, industrial and agricultural laboratories

The main limitations of the glass electrode are:

• It does not function properly in some organic solvents (i.e. ethanol)
• It does not function properly above pH > 9 since it is sensitive to Na⁺ ions so a correction has to be made

In case that the pH meter has two probes (two electrodes): i) one of them is a glass electrode (has silver wire suspended in a solution of KCl that is contained in a special glass bulb coated with silica and metal salts) and ii) the other is the reference electrode and has a KCl wire suspended in a solution of KCl (see Figure 3).

• A higher voltage means more H⁺ ions in the solution and therefore a higher acidity. The pH meter shows in such a case a lower pH value since the solution is more acidic
• It does not function properly above pH > 9 since it is sensitive to Na⁺ ions so a correction has to be made

When the probe(s) are immersed into the solution some of the H+ ions in the solution move toward the glass electrode (Figure 3, labeled as 2) and replace some of the metal ions in its special surface. This creates a tiny current (voltage) that the silver electrode passes to the measuring device, the voltameter. The voltameter measures the voltage generated and shows a corresponding pH measurement as follows:

• A higher voltage means more H⁺ ions in the solution and therefore a higher acidity. The pH meter shows in such a case a lower pH value since the solution is more acidic
• It does not function properly above pH > 9 since it is sensitive to Na⁺ ions so a correction has to be made

The reference electrode (Figure 3, labeled as 6) acts as a reference for the measurement.

How accurate measurements can be made with a pH meter?

A calibrated instrument must be used for accurate measurements. A procedure for calibrating pH meters is given in the post “Calibrating a pH meter”

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Relevant Posts

Calibrating a pH meter using buffers

pH Buffer Solution Preparation

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References

1. U.S. Pharmacopeia, 68, USP 36
2. David W. Oxtoby, H.P. Gillis, Alan Campion, “Principles of Modern Chemistry”, Sixth Edition, Thomson Brooks/Cole, 2008
3. Steven S. Zumdahl, “Chemical Principles” 6th Edition, Houghton Mifflin Company, 2009

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Key Terms

pH meter, using a pH meter to measure pH, measuring the pH of a solution,

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Measuring the pH of a solution / Acid-Base Indicators

The pH of a solution can be measured as follows:

• By Acid-Base Indicators (less precise)
• By a pH-meter

What kind of substances are acid-base indicators?
Acid-base indicators are usually weak organic acids or weak organic bases. They tend to have different color depending on the pH of the solution in which they are in.

How acid-base indicators are prepared in the lab?
These are usually solid substances that are dissoved in a solvent (i.e. ethanol). Few drops of the solution of the indicator is added to the solution that we would like to determine the pH.

How simple acid-base indicators work?

An acid - base indicator is a colored substance that itself can exist in either an acid or base form. The acid form has a different color than the base form. Thus, the indicator turns one color in an acidic solution and another color if placed in a basic solution. If you know the pH at which the indicator turns from one form to the other, you can determine whether a solution has a higher or lower pH than this value.

For example methyl orange is one of the indicators commonly used in titrations. It gradually changes color from red to yellow over the pH interval from 3.1-4.4.  In a solution with a pH > 4.4  exists as a species with negative charge (anion, Meo^- ) and has a yellow color. In a solution with a pH < 3.1 exists in its neutral form and haw a red color (ΗMeo).

In reality what happens is that the two forms of the indicator participate in an equilibrium:

 ΗMeo   +  H₂O          Meo^-   +  H₃O⁺           [1]     

If acid is added the position of the above equilibrium shifts to the left according to Le Chatelier's Principle and turns the indicator red (the solution takes a red color).
If base is added the position of the equilibrium shifts to the right according to Le Chatelier's Principle and turns the indicator yellow (the solution takes a yellow color).

The Ηenderson-Hasselbach equation can be used in order to determine the pH range an indicator changes color. Let's apply this for the methyl orange case:

 pH = pk_a + log [Meo^-] / [ΗMeo]      [2]

where k_a is the ionization constant of methyl orange.

 It has been determined experimentally that when 90% or more of the indicator is in the ΗMeo form (that means when the ratio  [Meo^-] / [ΗMeo] ≈ 0,1) then the color of the solution is red. If 90% or more of the indicator  is in the Meo^-  form (that means [Meo^-] / [ΗMeo]  ≈ 10) the the color of the solution becomes yellow. By subsituting the above ratios to the Ηenderson-Hasselbach equation the pH range an indicator changes color can be determined:

 pH = pk_a + log  [Meo^-] / [ΗMeo] = pk_a + log(0,1) = pk_a – 1       [3]

  και

 pH = pk_a + [Meo^-] / [ΗMeo] = pk_a + log(10) = pk_a + 1        [4]

When [Meo^-] = [ΗMeo] the color of the indicator is a mixture of yellow and red and the solution takes an orange color.

From equation [3] and [4 ] it can be determined that the indicator changes color over a range of  two pH units (when the pH is between  pk_a + 1 and  pk_a - 1). 

As a conclusion for monoprotic indicators of the general  structure ΗIn with an equilibrium constant k_a:

 If  pH < pk_a – 1, then the color of the solution takes the color  of the ΗI form (unionized form)

 If  pH > pk_a + 1, then the color of the solution takes the color of the I^- (ionized form)

If  pH = pk_a then the color of the solution is a «mixture» of the colors of  ΗI and I^-.

 In the video below the color change of an indicator is shown as the pH of the solution in which is in changes from neutral (pH of distilled water) to basic and then back from basic to acidic: