Calibrating a pH meter using buffers

How to calibrate a pH meter using buffers

A pH meter requires proper calibration in order to give accurate pH readings. The meter must accuratetely translate voltage measurements into pH measurements.

How a pH meter is calibrated to give accurate pH readings?

A pH meter is calibrated by immersing its electrode(s) into buffers (test solutions of known pH) and by adjusting the meter accordingly. Since pH measurements are affected by temperature, the temperature must remain constant during the time of the calibration. Modern pH meters have build-in thermometers and automatically correct their own pH measurements as the temperature changes. See the following video to get an idea how a pH meter is calibrated using buffer solutions:

The following procedures are used to calibrate pH meters (calibration procedure):

A 2 or 3-point calibration : Two or three buffers solutions are used respectively. They are usually sufficient for initial calibration. After this initial calibration the meter can accurately measure pH values in between.

A 1-point calibration: A buffer with a pH close to the expected sample pH is used. The pH measurements are not as accurate as with the 2 or 3 point calibration.

Procedure for an 1-point calibration:

1. Place the pH buffer solution (normally pH = 4.01 solution) into a small beaker. Place a magnetic stirrer and a temperature probe into the buffer solution in the beaker.
2. Measure the temperature of the buffer solution if the pH meter does not have its own temperature probe. Adjust the temperature control of the pH meter according to the buffer’s measured temperature (remember pH measurements are temperature dependent).
3. Remove the electrode protective cap(s). Rinse the electrodes and the temperature probe with distilled water using a squeeze bottle.
4. Dab dry the bottom of the glass bulb with a tissue paper. Do not wipe the glass bulb (scratches and static charges affect the electrode’s response).
5. Place the electrode(s) and temperature probe into the well stirred pH buffer solution. The porous frit must be covered with the buffer solution
6. Adjust the slope/sensitivity control to read the true pH of the buffer solution (modern pH meters do this automatically)

Procedure for an 2-point calibration:

1. Place the pH buffer solution (normally pH = 7.0 solution) into a small beaker. Place a magnetic stirrer and a temperature probe into the buffer solution in the beaker.
2. Follow steps 2-6 of the 1-point calibration.
3. Follow steps 1-6 of the 1-point calibration using the pH = 4.01 buffer solution
4. Keep repeating steps 2 and 3 until practically no adjustments are required

Procedure for a 3-point calibration:

1. Follow steps 1-6 of the 2-point calibration procedure
2. Use a third pH buffer, whose pH value is as close as possible to the suspected pH of the sample, and follow the 1-point calibration. The third point is used to confirm that the pH meter has been calibrated correctly.

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Relevant Posts

Measuring the pH of a Solution with a pH meter

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, calibrating a pH meter, buffers, calibrating ph meter using buffers,

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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: