Titration Procedure

Standardizing a Base

Suppose that you needed a 0.1000 M solution of NaOH to do a titration. You might think that you could just weigh 4.000g of solid NaOH and dissolve it in enough water to make 1.000 L of solution.  4.000 g is after all 0.1000 mole of NaOH, so dissolving it in one litre of water should produce a one litre solution.

However, it is very difficult to prepare a solution of NaOH of an exact concentration for two reasons:

• Solid NaOH is very hygroscopic, that is it absorbs water from the air, so it is very hard to weigh exactly.
• Solutions of NaOH react with carbon dioxide from the air. The overall reaction is
          CO₂ (g) + 2 NaOH (aq) Na₂CO₃ + H₂O
  This decreases the concentration of the OH^- ions in the solution.  Unless the bottle is tightly sealed, the molar concentration of the OH^- solution will change on a daily basis.

In order to have accurate solutions of NaOH, you must standardize the solution. The best way of doing this is to use a primary standard, which is a solid acid that can be weighed. To have a good primary standard it should not absorb anything readily from the air.  It should also have a high molar mass (so that you can weigh out reasonable quantities of it without getting too many moles to do a good titration.)

In this lab you will calibrate the molarity of some NaOH by titrating it against potassium hydrogen phthalate (chemical formula is KHC₈H₄O₄). This molecule is a solid acid that produces one ionizable hydrogen.

Procedure

1. Wear your goggles at all times. NaOH is highly corrosive, especially to the eyes.
2. Set up your buret. Get about 150 mL of NaOH in a clean beaker (this NaOH will have a molarity of about 0.2 M). Use several rinses with small quantities of this base to clean your buret.
3. On the balance, tare a piece of weighing paper and measure out about 1.00 – 1.25 g of KHC₈H₄O_4. Don’t use less than nor more than this amount.
4. Add the KHC₈H₄O₄ to a clean erlenmeyer flask and dissolve it in about 50 mL of water. Don't use a graduated cylinder to measure the water, just use the calibration marks on the flask. You are not trying to prepare a solution of any particular molarity, just adding enough water to dissolve the mass of KHC₈H₄O₄ that you weighed. Using another container to measure the water just introduces another possible source of contamination.
5. Add some phenolphthalein indicator to the acid. Fill the buret, measure the starting volume, and titrate to the endpoint. Note: it will take more than 20 mL to do the titration. Measure the final volume.  Make sure to estimate between the calibration marks on the buret so that you get a reading to the nearest 0.00 (hundredths) mL.
6. Rinse out the flask, then repeat steps 3 to 5. After each trial calculate the molarity of the base. Repeat until you have results for the molarity that are consistent to no more than 2 % error.  Calculate the average molarity of the NaOH (this value can then be used in further titrations of other acids).

Calculations

Using the measured mass, and the molar mass of KHC₈H₄O₄, calculate the number of moles of solid acid used. Since the acid was a monoprotic solid, and the base is a solution, you will have to make use of the relationship between moles of solid acid and the molarity of the base and its volume in litres:

mol_acid = (M_base)(L_base)

Rearrange this equation to solve for the molarity of the base:

If you need some help doing the calculations, click on the question marks.

2. Complete the following table for each trial that you did (there is room for four trials, but if you did not do that many, then leave the extra columns blank).  Make sure to record the correct number of significant digits for each measurement you made.   The lines in blue are results you calculated.

	Trial 1	Trial 2	Trial 3	Trial 4
Mass of KHC8H4O4 (g)
Moles of KHC8H4O4 (mol)
Initial volume of NaOH (mL)
Final volume of NaOH (mL)
Molar concentration of base