Acid-Base Titration

Acid-base Titration Definition

Acid-base Titration Theory

Before we dive into the experiment itself, let's do a recap of acid-base reactions. Acid-base titrations hinge on the fact that the pH of a solution changes when an acid and base are reacted together. When a base is added, the pH increases, the opposite is true for acids. When the pH of a solution is equal to 7, it is at the equivalence point, which is the point where the concentration of the acid is equal to the concentration of the base. The formula for this is:

M₁V₁ = M₂V₂

where, M₁, is the molarity of solution 1, M₂, is the molarity of solution 2, V₁, is the volume of solution 1, and V₂, is the volume of solution 2.

Acid-Base Titration Example

Let's look at an example:

Now that we know the basics, let's look at how we perform the titration.

Acid-base Titration Procedure

Let's look at how we would perform an acid-base titration in the lab. For our first step, we need to pick our titrant. Since this is an acid-base reaction, if our analyte is an acid, the titrant has to be a base and vice-versa. We take our titrant and pour it into a buret (a long tube with a dropper at the bottom). The buret is clamped above a flask which will be filled with the analyte (make sure to note the volume of both the titrant and analyte). The next thing we need to do is add the indicator to the analyte solution.

Many indicators will change color at certain pH ranges. When choosing an indicator, you want to pick one that will change color at a pH close to the endpoint. Here are some common indicators:

Once we have picked our indicator, we will add a few drops of it to our analyte solution. Next, we will turn the buret open, so drops of the titrant can flow out. When a flash of color appears, we close the buret slightly to slow down the flow. When the color stays for longer, we swirl it around until it returns to its original color. Once the indicator has changed color and stayed that way for several seconds, the titration is finished.

The setup for the titration. The pink splash is Phenolphthalein beginning to change color, indicating we are near the endpoint. Pixabay

We note the final volume of the titrant, then repeat the experiment a few times for accuracy. Once we have our average volume of titrant used, we can use that to calculate the concentration of the analyte.

Acid-base Titration Curves

The way we visualize these titrations is through titration curves.

A titration curve can help us figure out the volume of the titrant at the equivalence point. The equivalence point is always at pH = 7 since the solution will be neutral when there are equal amounts of acid and base. The shape of the curve is dependent on the strength of the acid/base and whether the analyte is an acid or base. Let's look at an example:

There are 4 different possible shapes for the curve when an acid is the analyte. StudySmarter Original

There are 4 different possible shapes for the curve when a base is the analyte. StudySmarter Original.

You'll notice that there are technically 4 shapes, as the base analyte curves (in blue) are mirrors of the acid analyte curves (in red). For example, the weak acid/strong base curve for the acid analyte is the reverse of the strong acid/weak base curve. To help pick out an indicator, you need to know the identity of the titrant and analyte as well as their strengths, then you can match up the pair to the curve.

Polyprotic Acid-base Titrations Examples and Curves

The titrations we've looked at previously have all been with monoprotic acids, but these titrations can also be done with polyprotic acids. These are acids that have more than one proton to donate. The titration curves for these look different since there are multiple equivalence points: one for each proton donated. Let's first look at one of these curves:The titration curve of a polyprotic acid (analyte) with a strong base shows the different equivalence points for each step of the reaction. StudySmarter Original

There's a lot going on in this curve, so let's break it down piece by piece. Let's start by looking at the equations for these reactions:

$$H_2SO_{3\,(aq)} +NaOH_{(aq)} \rightarrow HSO_{3\,(aq)}^{-} + H_2O_{(l)}+Na^+$$

$$HSO_{3\,(aq)}^- +NaOH_{(aq)} \rightarrow SO_{3\,(aq)}^{2-} + H_2O_{(l)}+Na^+$$

Sulfurous acid, H₂SO₃, has 2 protons it can donate, so it has two equivalence points, as shown by the circles on the graph. Their equations are:

$$[HSO_3^-]=[NaOH]\,\,\text{(equivalence point 1)}$$

$$[SO_3^{2-}]=[NaOH]\,\,\text{(equivalence point 2)}$$

The other key points on this graph are the half-equivalence points, triangles on the graph. These are when the concentration of the acid is equal to the concentration of its conjugate base. Their equations are:

$$[H_2SO_3]=[HSO_3^-]\,\,\text{(half-equivalence point 1)}$$

$$[HSO_3^-]=[SO_3^{2-}]\,\,\text{(half-equivalence point 2)}$$

One thing to note is that polyprotic acids are always weak acids. As you can see in the graph, the acid gets weaker as it loses more protons, so the "spike" at the equivalence point gets smaller. But what if our analyte is a base?

The titration curve for a base that becomes a polyprotic acid. This curve is a mirror of the polyprotic acid analyte curve. StudySmarter Original

In this reaction, Na₂SO₃ is our base. Let's look at the reactions:

$$Na_2SO_{3\,(aq)} + HCl_{(aq)} \rightarrow NaHSO_{3\,(aq)}^- + NaCl_{(aq)}$$

$$NaHSO_{3\,(aq)}^- + HCl_{(aq)} \rightarrow H_2SO_{3\,(aq)} + NaCl_{(aq)}$$

So instead of having a polyprotic acid donate multiple protons, we have a base gaining those protons to form the polyprotic acid. It can do this since HCl is a much stronger acid than H₂SO_3.