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Detection of Ions in Solutions Using Acid/Base Chemistry: A Quality Control Test
Objective: In this experiment you will standardize a NaOH solution to know the exact concentration. In follow up experiments, you will used the standardized NaOH to determine the unknown concentrations of HCl and HC 2 H 3 O 2 solutions. You will learn to detect endpoints using different indicators.
What is a Titration?
A titration is an analytical procedure used to determine the concentration of a sample by reacting it with a standard solution. One type of titration uses a neutralization reaction, in which an acid and a base react to produce a salt and water.
In equation 1, the acid is HCl (called hydrochloric acid) and the base is NaOH (called sodium hydroxide). When the acid and base react, they form NaCl (sodium chloride), which is also known as table salt. The titration proceeds until the equivalence point is reached, where the number of moles of acid is equal to the number of moles of base. This point is usually marked by observing a color change in an added indicator.
In a titration, the standard solution goes in a buret, which is a piece of glassware used to measure the volume of solvent to approximately 0.1 mL of accuracy. The solution that you are titrating goes in an Erlenmeyer flask, which should be large enough to accommodate both your sample and the standard solution you are adding.
What is an Indicator and What is it Used For?
An indicator is any substance in solution that changes its color as it reacts with either an acid or a base. Selecting the proper indicator is important because each indicator changes its color over a particular range of pH values. Indicators are either weak acids or weak bases. For example, phenolphthalein is a weak acid (which we will represent as HIn). In aqueous solution, the phenolphthalein dissociates slightly, forming an equilibrium.
An equilibrium occurs when the amount of reactants and the amount of products are constant. This means the rate of the forward reaction is the same as the rate of the reverse reaction. When a system is in equilibrium, it will stay there until something changes the conditions. A famous French chemist, named LeChatelier, developed a way to predict how changes in equilibrium affect the system. LeChatelier’s principle states that when an equilibrium is disturbed by applying stress, the equilibrium will shift to relieve the stress. In an acidic solution , there is an excess of H 3 O+^ ions so the equilibrium will shift to the left and favor the formation of HIn, thus we observe a clear solution. In basic solution , there is an excess of OH-^ ions that react with the H 3 O+^ ions to form water. This shifts the equilibrium to the right because water is being formed and H 3 O+^ ions are being removed, thus we observe a pink solution. We can use this color change to determine when the end of the titration has been achieved.
Table 1 lists common indicators and the pH range over which they change colors.
Standardizing a Sodium Hydroxide (NaOH) Solution
In a titration, it is critical to know the exact concentration of the titrant (the solution in the buret which will be added to the unknown) in order to determine the concentration of the solution being tested. We will standardize the ~0.1 M NaOH solution (the titrant) with potassium hydrogen phthalate (KHP, KC 8 H 4 O 4 H) using phenolphthalein as the indicator. KHP is a weak acid and reacts with base in the following way:
Dry primary standard potassium hydrogen phthalate, KHP, for two hours at 100 oC. This has been done for you. You will not need to dry the KHP before you do the experiment. The technique information remains for your learning pleasure. You should understand the necessity to dry the salt first.
Weigh ~0.75 g of dried KHP (MW = 204.23 g/mol) into an Erlenmeyer flask. Determine the exact mass by weighing by difference.
Dissolve your sample in 50-75 mL of distilled water. Determine the amount of water added by weighing by difference.
Record the amount of KHP and water used.
Add 4 drops of indicator into the flask.
Obtain approximately 100 mL of NaOH solution in a clean, dry beaker. Prepare the buret for titration by rinsing with 2 aliquots of solution. Fill the buret and record the initial volume to 0.1 mL.
Titrate to the first permanent appearance of pink. Near the endpoint, add the NaOH dropwise to determine the total volume most accurately. Record the final buret reading.
If time allows, repeat the titration.
Calculations and Questions
- Calculate the number of moles of KHP used in the titration.
- Calculate the molarity of the KHP solution after your addition of water.
- Determine the number of moles of NaOH required to reach the endpoint of the titration.
- Calculate the molarity of the NaOH standard solution.
- How would the amount of NaOH required to neutralize the KHP have changed if you added twice the volume of water to the KHP before the titration? Justify your answer.
- During the titration, splatter from the drops of NaOH splashed some of the KHP solution onto the sides of the Erlenmeyer flask. If the student did not rinse these droplets back down into the solution with distilled water, how would that have affected the calculation of concentration of NaOH?
