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Overview

Not sure what titration is or what you can do with it? Then you are in the right place! In this tutorial, you will find information on titration, including the chemicals that are commonly used and the chemical reactions that make titration work, as well as how titration is performed and some tips to get better results.

What is Titration?

Titration is a method commonly used in chemistry to figure out the amount of a chemical in a solution. This chemical is called the titrand. To figure out the amount of titrand in the solution, a known amount of a different chemical is added to the titrand's solution; this chemical— called the titrant, or titrating solution—reacts with the titrand. By measuring how much of the titrating solution is needed to react with all of the titrand in the solution, you can calculate how much titrand was in the solution. Simply put:

The point at which all of the titrand has reacted is called the endpoint, or equivalence point. How do you know when the endpoint has been reached? At the endpoint, there is usually a color change in the titrand's solution. This is caused by an indicator, which is included in the titrand's solution just so you can see when you have reached the endpoint. The unknown amount of titrand in the titrand's solution can usually be determined by setting up a proportion with the known amount of titrating solution that was added. How this is done depends on the exact titrating reaction that is being used.

You can watch the video below, made by the Massachusetts Institute of Technology (MIT)'s Department of Chemistry, to see titration in action. Note: This video uses an indicator that turns light pink at the endpoint, but different indicators turn different colors at their endpoints. The next section contains more information about indicators.

Titration | MIT Digital Lab Techniques Manual
Watch this titration video made by the Massachusetts Institute of Technology (MIT)'s Department of Chemistry.

Indicators

There are many different types of indicators used in titration experiments. Which indicator is used depends on the chemistry of the reaction taking place between the titrand and the titrating solution. This titration tutorial will cover two commonly used indicators—starch and phenolphthalein—along with their associated reactions.

Starch as an Indicator

Starch is an indicator that reacts with iodine. When there is iodine present, starch reacts with it to create a blue chemical complex. This means the solution turns blue! How is this used in titration? Iodine is included in the titrating solution, and as it is added to the titrand's solution (which includes the titrand and starch), the titrand reacts with the iodine to turn it into iodide ions (which do not react with the starch). However, as soon as all of the titrand has reacted with the iodine and the endpoint is reached, the addition of any more iodine will finally react with the starch and turn the titrand's solution blue!

An example of titration using a starch indicator is the titration of vitamin C, which is technically ascorbic acid. Ascorbic acid reacts with iodine to make dehydroascorbic acid and iodide ions. (This reaction is technically an oxidation-reduction reaction, also called a redox reaction for short.) When ascorbic acid and starch are both in a solution, iodine will react with the ascorbic acid. So when titrating ascorbic acid, a titrating solution containing iodine is added to the titrand's solution, which contains starch (the indicator) and ascorbic acid (the titrand), and when all of the ascorbic acid has reacted with the iodine, any more iodine added will react with the starch and turn the titrand's solution blue! Figure 1, below, shows a picture of the endpoint of an ascorbic acid titration using starch and iodine. Because there is a known concentration of iodine in the titrating solution, by keeping track of how much solution is added, you can determine how much titrand there was.

Note: Because elemental iodine is needed to react with the starch, pure potassium iodine does not work as an indicator in this titration reaction and Lugol's solution (which contains elemental iodine and potassium iodide) must be used instead.

The following Science Buddies Project Ideas involve titration with a starch indicator:

Black liquid at the bottom of a flask
Figure 1. The titrand's solution turns blue-black when the endpoint has been reached in a titration using starch as an indicator (to react with iodine).

Phenolphthalein as an Indicator

Phenolphthalein is an indicator that changes color depending on the pH of the solution it is in. The pH of a solution is a measure of how acidic or basic it is. (For a refresher, see the Science Buddies resource on Acids, Bases, & the pH Scale.) Specifically, phenolphthalein is colorless when the pH of a solution is acidic or neutral, but when the solution becomes slightly basic, phenolphthalein turns slightly pinkish, and then darker pink as the solution becomes more basic. How is this used in titration? A base is included in the titrating solution, and it is added to the titrand's solution, which contains an acidic titrand and phenolphthalein. As more base is added to the titrand's solution, the pH changes, becoming more basic, and the solution changes color. Usually, with this indicator, when the titrand's solution just starts to turn pink, you have reached the endpoint.

An example of titration usng phenolphthalein is the titration of vinegar, which is technically acetic acid. When titrating acetic acid, a titrating solution containing a base—normally sodium hydroxide—is added to the titrand's solution, which contains phenolphthalein (the indicator) and acetic acid (the acidic titrand). (The acetic acid reacts with the sodium hydroxide in an acid-base reaction.) When the titrand's solution becomes basic enough due to the addition of the basic titrating solution, the phenolphthalein turns the titrand's solution slightly pink. Phenolphthalein is specifically colorless at a neutral or acidic pH, and becomes light pink as the pH becomes more basic (first turning slightly pink around a pH of 8.3). Figure 2, below, shows a picture of the endpoint of an acetic acid titration using phenolphthalein and sodium hydroxide. Because the number of moles of sodium hydroxide used to titrate the acetic acid equals the number of moles of acetic acid in the titrand solution, by keeping track of how much titrating solution is added, you can determine how much titrand there was.

The following Science Buddies Project Idea involves titration with a phenolphthalein indicator:

Purple translucent liquid in a glass jar underneath a buret
Figure 2. The titrand's solution turns slightly pink when the endpoint has been reached in a titration using phenolphthalein as an indicator (to show the change in pH).

Setup & Procedure

There are many steps that should be taken to ensure that a titration is successful and that the results produced are accurate. Check out the video of best practices in titration. Here are some key points to follow and keep in mind when doing a titration:

A clamp holds a buret on a ring stand
Figure 3. This picture shows a general titration setup. Note that the buret clamp is firmly attached to the ring stand. The buret shown here slides into place between the prongs of the buret clamp. The buret is held firmly in place, but can be moved up and down if needed.

An erlenmeyer flask filled with a light orange solution is placed under a buret
Figure 4. The titrand's solution is placed in an Erlenmeyer flask and set right below the bottom of the buret. An Erlenmeyer flask is used because its shape allows a person to swirl the solution to mix it without spilling.
A valve on a buret is turned to the closed position
Figure 5. Make sure the buret is in the closed position, as shown here with the red stopper on this buret, before filling it with titrating solution.

Small funnel inserted into the top of a buret
Figure 6. Place a small funnel in the top of the buret to slowly fill your buret with the titrating solution.
The meniscus is measured in a red solution contained within a graduated cylinder
Figure 7. When reading the level of liquid in the buret, read from the bottom of the meniscus, which is being pointed to with a black arrow in this picture. For example, the level of the liquid in this buret should be read as 21.85 mL. (Note: The long white line at the top of the buret is the mark for 21 mL.)
Cloudy liquid appears at the bottom of an erlenmeyer flask
Figure 8. This example of titrating ascorbic acid using an iodine titrating solution shows how you may see a temporary (dark blue) color change when adding a drop of the iodine solution. The temporary color change goes away when the iodine solution is mixed in the flask by swirling it.

Interpreting Results

The unknown amount of titrand in the titrand solution can usually be determined by setting up a proportion with the known amount of titrating solution that was added. How this is done depends on the exact titrating reaction that is being used.

For example, if you are titrating ascorbic acid using an iodine titrating solution, you need to titrate an ascorbic acid standard, which is a sample of ascorbic acid with a known amount of ascorbic acid in it. Knowing how much iodine titrating solution is needed to titrate this known amount of ascorbic acid, you can set up a proportion to determine how much ascorbic acid is in other samples that were titrated. See Equation 1, below, for how this can be done:

Equation 1.

Equation 1 can be rearranged to directly solve for the unknown amount of ascorbic acid in the sample ("Ascorbic Acid2"). This gives us Equation 2, below:

Equation 2.

As an example, if you were using 20 mL of a 1 mg/mL ascorbic acid standard solution, this means you had 20 mg of titrand in your standard solution ("Ascorbic Acid1"). If it took 8.5 mL of iodine titrating solution ("Iodine1") to titrate this 20 mg, but to titrate an unknown sample of ascorbic acid it took 6.8 mL of iodine titrating solution ("Iodine2"), you can use the equation to determine that the amount of ascorbic acid in the unknown sample ("Ascorbic Acid2") equals 16 mg.

In the example titration using phenolphthalein in the titration of acetic acid, the unknown amount of acetic acid (the titrand) can again be determined by setting up a proportion with the known amount of sodium hydroxide (the titrating solution). Specifically, the number of moles of sodium hydroxide used to titrate the acetic acid equals the number of moles of acetic acid in the titrand's solution. For example, if you added 12.5 mL (0.0125 liters [L]) of a 0.1 molar (M, which is moles/L) sodium hydroxide to titrate the acetic acid, the number of moles of both sodium hydroxide and acetic acid would be 0.0125 L x 0.1 moles/L = 0.00125 moles. You could divide by the amount of the sample (in liters) to determine the molar concentration of the acetic acid. For example, if your sample volume was 1.5 mL (0.0015 L), it would have a molarity of 0.00125 moles / 0.0015 L = 0.833 M.

References

If you have read through this titration resource and still have unanswered questions, the references below might be helpful.

The following are resources from Science Buddies on acids, bases, and the pH scale, as well as chemistry safety:

The following science project ideas on titration are available from Science Buddies:

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