How to Determine the pKa of an Unknown Acid

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The pKa of an acid is a measure of its strength: the lower the pKa, the stronger the acid. Hydrochloric acid has a pKa of -8, for example, while acetic acid (vinegar) has a pKa of 4.8. Determining the pKa of an unknown acid is a common assignment in introductory lab classes; it's also possible you might need to make this measurement as part of a more advanced experiment.

Things You'll Need

  • Calibrated pH meter
  • Goggles
  • Gloves
  • Buret
  • 0.1 molar sodium hydroxide (NaOH) solution
  • Beaker
  • Magnetic stir bar and plate
  • Graduated cylinder
  • Funnel
  • Make sure you're wearing safety goggles and gloves before you start.

  • Measure out 40 mL of your unknown acid solution using the graduated cylinder; try to be as accurate as possible. Pour it into the beaker and set the beaker on the magnetic stir-bar plate. Be very careful with the acid; if you don't know what it is, you have no guarantee that it's not corrosive or toxic.

  • Rinse your buret with the 0.1 molar sodium hydroxide (NaOH) solution, being careful not to spill any of the NaOH -- it is highly corrosive. A good way to rinse it is to close the stopcock, pour a little NaOH in at the top with the funnel, place a waste beaker beneath the buret and open the stopcock to allow the solution and any air to run out. Repeat this process, but this time, close the stopcock before all the solution has run out of the buret. Fill the buret at least halfway with NaOH solution. Take special care during this procedure that you do not attempt to pour NaOH from close to or above eye level -- the top of the buret should be below your eye level when you do this.

  • Measure the volume of NaOH in the buret using the markings on the glass. Make sure you read the level from the bottom of the meniscus (the concave depression in the surface of the water). Write down the volume in your notebook; use Excel or other graphing program to create a spreadsheet.

  • Add the stir bar to the acid solution. Turn the stir-bar plate on and slowly increase the setting -- increase it too quickly and the stir bar won't work properly. The stir bar should spin around slowly to keep the fluid in the beaker circulating.

  • Insert the pH meter into the acid solution, but do not allow it to come into contact with the stir bar -- pH meters are expensive, and it's worth your while to avoid damaging them. Measure and record the pH of the acid solution.

  • Draw up a table using your spreadsheet, listing pH in one column and volume of NaOH in another.

  • Slowly open the stopcock and allow NaOH solution to drip into the beaker until you've added 2 mL. Stop and record the pH and amount of NaOH added (2mL). Repeat this procedure until the pH change associated with a single 2 mL addition is 0.2 pH points or greater.

  • Add 1 mL of NaOH then stop and record pH. Repeat this procedure until the pH change associated with a single 1 mL of titrant (re-agent solution of precisely determined concentration) is 0.2 pH points or greater, then switch to adding 0.5 mL of NaOH. Once the pH change associated with a single 0.5 mL of titrant is 0.2 pH points or greater, switch to adding 0.1 mL at a time.

  • Continue to add NaOH and record pH until the solution in the beaker reaches pH 11.

  • Clean up your glassware, equipment and work area, then remove and discard your gloves. Remember that corrosive liquids (including both your unknown acid solution and NaOH solution) should never be poured down the drain -- they belong in your school lab's hazardous waste bucket instead. Be sure to neutralize any acid or base spills with sodium bicarbonate/citric acid as appropriate.

  • Using Excel or another graphing program, enter the data you recorded during the experiment and graph it. Assuming your acid is monoprotic (i.e., donates only one proton), which is usual for undergrad lab class titrations, you'll find the graph will be roughly sigmoid or S-shaped. Note the part of the graph where the pH changed very rapidly (i.e., the line becomes very steep) and find the center of this steep part. This is your inflection point.

  • Note the volume of NaOH added to reach the inflection point. Again, assuming your acid is monoprotic, the number of moles of NaOH added are equal to the number of moles of unknown acid originally present. A mole measures the amount of substance; one mole is equal to 6.022 x 10^23 molecules, ions or atoms.

  • Write down the volume of NaOH required to go one-third of the way, one-half of the way and two-thirds of the way to the inflection point. Find the pH at these points on your graph.

  • Calculate the pKa at each of these three points using the Henderson-Hasselbach equation:

    pH = pKa + log [B] / [A], in which "B" is conjugate base and "A" is the acid.

    For the one-third mark: At this point, one-third of the original acid has been converted into conjugate base. [B] / [A] = (1/3)/(2/3) = 0.5. Take the log of 0.5, subtract it from the pH, and you have your pKa.

    For the one-half mark: [B] / [A] = (1/2) / (1/2) = 1. The log of 1 is 0, so the pH at this point is equal to the pKa.

    For the two-thirds mark: [B] / [A] = (2/3) / (1/3) = 2. Take the log of 2 and subtract it from the pH to get the pKa.

  • Average all three results by adding them together then dividing by three. This will give you a fair estimate of the pKa of your unknown acid.

References

  • Photo Credit NA/Photos.com/Getty Images
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