Introduction to titrations

A titration is a procedure used by chemists to determine the concentration of a particular solute in a solution.

To understand how titrations work, let’s do an example.

Let's say we are given an aqueous solution of hydrochloric acid.

We do not know the concentration of the solution, but we do know its volume, which is 20 milliliters.

To perform a titration, we need a piece of equipment called a burette. Into this burette, we add a 0.100 molar solution of aqueous sodium hydroxide.

Since we know the concentration of our sodium hydroxide solution, we call it our “known” solution.

The known solution is also called the “titrant.”

Since we don't know the concentration of our HCl solution, we call it our “unknown” solution.

To do a titration, we are going to add the known solution to the unknown.

Before we start doing this, we need to carefully record the initial volume of the solution in our burette.

The numerical value of the initial volume does not matter, it just matters that we know the exact starting point so that we can compare it to our ending volume later.

We also need to add into the hydrochloric acid solution a few drops of phenolphthalein, which is an acid base indicator.

Phenolphthalein is colorless when it is in an acidic environment.

But in a basic environment, the indicator will turn the solution pink.

Next, we open up the stopcock on the burette and allow the sodium hydroxide to gradually drip into the hydrochloric acid solution.

As the sodium hydroxide is added, the aqueous solution of sodium hydroxide reacts with the aqueous solution of hydrochloric acid in an acid-base neutralization reaction to form an aqueous solution of sodium chloride and water.

The net ionic equation leaves out all of the ions that don't participate in the reaction.

Remember that those ions are called are spectator ions.

Since sodium ions and chloride ions are present in solution both before and after the acid-base reaction, they are the spectator ions.

Therefore, we can take out the spectator ions to give the net ionic equation of OH minus plus H plus goes to H2O.

So, as we add hydroxide ions, we are neutralizing the H plus ions in the unknown.

As soon as we have neutralized all of the H plus ions, the unknown solution will no longer be acidic.

Any additional hydroxide ions will form a basic solution and the phenolphthalein indicator will turn the solution pink.

This color change indicates the endpoint of the titration and means that all of the acid present in the unknown has been completely neutralized.

In the titration, when we see this color change, we should immediately close the stopcock on the burette.

At this point, our next step is to record the final volume of the known solution in our burette.

We subtract this final volume from the initial volume to find the total volume of base that was needed to neutralize the acid.

For this problem, let's say it took 48.6 mL of the base to neutralize our acid sample.

Given this information, we are now ready to find the concentration of our hydrochloric acid.

Since we know that the sodium hydroxide solution was 0.100 molar and we used 48.6 mL of it to neutralize the acid, we can calculate the total number of moles of base that we added.

Molarity is equal to moles over the volume of the solution (in liters), and so the number of moles equals the molarity times volume.

We first convert 48.6 milliliters into 0.0486 liters, and we multiply this number by 0.100 molar to get 0.00486 moles of sodium hydroxide.

Since the mole ratio of sodium hydroxide to hydroxide ions is one to one, 0.00486 is also the number of moles of hydroxide ion in solution.

Next, we look at the net ionic equation for the neutralization reaction.

Since all of the mole ratios are one to one, for every mole of hydroxide that we added, we consumed one mole of H plus ions, and we produced one mole of water molecules.

So, if we added 0.00486 moles of hydroxide, we must have had 0.00486 moles of H plus ions present in the unknown.

The point in a titration where the number of moles of acid is equal to the number of moles of base is called the “equivalence point.”

The goal of a titration is to try to match the equivalence point to the endpoint.

Therefore, a titration is stopped the moment the color change occurs.

If we continue adding too much base after the equivalence point, we have “overshot” the titration.

To find the concentration of H plus ions we use the molarity equation.

Since we had 0.00486 moles of H plus ions and the original volume was 20.0 ml or 0.0200 liters, the concentration is equal to 0.00486 divided by 0.0200 which is 0.243 molar.

Since hydrochloric acid is a strong acid, the concentration of H plus ions is equal to the concentration of the HCl solution.

Therefore, the starting concentration of the hydrochloric acid solution was 0.243 molar.

We could have also done this problem using the MV is equal to MV equation, which states that molarity of acid times the volume of acid is equal to the molarity of base times the volume of base.