Titration is a very useful technique that allows a chemist to identify the concentration of an unknown solution by reacting it with a known solution.  Since we often measure concentrations in Molarity (mol/L) we can easily find the number of moles in a sample of solution by

n = MV

where n is the number of moles, M is Molarity of the solution, and V is the volume in liters.  Commonly, the solutions used are acids and bases.  The idea is to titrate a base into an acid until the last bit of acid is neutralized.  At that point, the moles of acid equal the moles of base, so if we know how much of one we used, we can figure out how much of the other was present.

We need to overcome a few potential problems before the technique becomes useful.  First, how will we know that the reaction is complete?  Since all the solutions are clear liquids, we might add more base than we needed to neutralize the acid if we are not careful.  Luckily there are substances, known as indicators, which are a different color in acid solution than they are in base solution.  For example, phenolphthalein is colorless in acid solution but bright pink in base solution.  This will tell us when the reaction is complete; we keep adding base to the acid until it turns pink.

A second problem is that we need a way to accurately measure the volume of base added to the acid.  We can’t just pour the base from a beaker because we would end up adding way too much before we reacted to the pink color and stopped pouring.  Chemists have invented a device called a buret for this purpose.  A buret is a long, thin tube, marked off in fractions of a milliliter, with a fine tip that can dispense the solution in small quantities.  Now we can add the base slowly to the acid, sometimes drop-by-drop, and stop as soon as we see a pink color.

Part A: Standardization of Solutions

We have one last problem to solve before titration becomes practical.  How will we know the Molarity of the base solution we are using to titrate the acid?  We can easily do a mole calculation to tell us how many grams of NaOH to dissolve in order to make a given volume of the desired Molarity, but how can we be sure the solution is exactly the Molarity we calculated?  After all, the NaOH crystals probably aren’t 100% pure, especially since NaOH tends to absorb water from the air.  To solve the problem we need to find a primary standard, an acidic substance that can be reliably purified to 100% and can be stored without changing its purity over time.  That substance is potassium hydrogen phthalate (KHP), an acidic crystalline salt.  We can measure out a desired number of moles of KHP and titrate our NaOH solution against it.  Using the first equation (n = MV), we can calculate the true Molarity of the base solution.  We say that we have standardized the base solution.  We can now use the standardized solution to titrate our unknown acid solution.

Methodology

1.	Determine the number of grams of NaOH needed to make 250 mL of a 0.100 M solution.  Obtain the correct amount of NaOH and dissolve it with distilled water in a 250 mL volumetric flask.

2.	Fill a buret with the NaOH solution, taking care to flush out any air bubbles in the tip of the buret.

3.	Obtain approximately 0.40 g of pure, dry KHP and record its exact mass.  Pour the KHP into an Erlenmeyer flask and dissolve it in approximately 100 mL of distilled water.  Add two drops of phenolphthalein indicator.

4.	Titrate the NaOH solution into the KHP until a pink color persists.  Record the volume of NaOH used and calculate the true Molarity of the solution.

5.	If time permits, repeat steps 3 and 4 and average your results.

Analysis

To standardize the NaOH solution, you will need to convert the mass of KHP that you used to moles.  The molecular mass of KHP is 204.23 g/mol.  Once you have converted to moles, you can plug into the equation

n = MV

to find the Molarity of the NaOH solution.  Remember that the volume must be expressed in liters in this equation!  You have now standardized your NaOH solution, which means that you know its correct Molarity.  You'll need to use that Molarity value in part B of the experiment.

Part B: Determining the Molarity of an Acid

One important use of titration is to find the Molarity of an acid solution of unknown concentration.  Consider for example the neutralization reaction between hydrochloric acid and sodium hydroxide:

HCl + NaOH  NaCl + H2O

There is a one-to-one stoichiometric ratio of reactants, which means that 1 mole of NaOH exactly neutralizes 1 mole of HCl.  Also remember that we can rearrange the Molarity equation to solve for moles as follows:

n = MV

where n is the number of moles, M is Molarity of the solution, and V is the volume in liters.  Since we know that the number of moles of HCl must equal the number of moles of NaOH when the reaction is complete, we can write

MaVa = MbVb

where the a subscript means acid and b means base.  So if we take a specific volume of an unknown acid solution, we can react it with a base solution of known Molarity, measure the volume of base used, and calculate the Molarity of the unknown acid.

Luckily, we happen to have a base of known Molarity, freshly standardized in part A.  We can use our standardized NaOH solution to find the Molarity of an unknown HCl solution.

Methodology

1.	Place exactly 20.0 mL of an unknown HCl solution in an Erlenmeyer flask.  Use a volumetric pipet to get the exact volume.

2.	Add approximately 50 mL of distilled water to increase the volume.  Add two drops of phenolphthalein.

3.	Titrate your standardized NaOH solution into the HCl solution until a pink color persists.  Record the volume of NaOH used and calculate the Molarity of the acid solution.

4.	If time permits, repeat steps 1-3 and average your results.

Analysis

The Molarity of the acid solution can be found using

MaVa = MbVb

which is often called the titration equation.  Note that you are free to express volumes in any unit you like in this equation since it represents a simple ratio.  Va is just the volume of HCl you placed in the flask.  Mb is the Molarity of the base, which you found in the standardization part of the experiment.  Vb is the volume of NaOH that you titrated into the HCl solution.  Plug those numbers into the equation and solve for Ma.

Questions

1.	Why do we have to standardize out NaOH solution?  In other words, why can't we trust the calculated Molarity of the solution?

2.	Why don't chemists make burets thicker so they can fit more NaOH in them?

3.	If a student uses too much NaOH to neutralize the KHP in the standardization run, how will that affect his results?

4.	In part B, we used a volumetric pipet to measure out the HCl sample.  Why didn't we just use a graduated cylinder?

5.	Indicators like phenolphthalein are actually weak acids.  If you were to add a lot of phenolphthalein to your sample, how would it affect your titration?