Acid-Base Titration Lab
Introduction
A titration is a lab technique used to determine the volume of a solution that is needed to react with a given amount of another substance. In this activity, your goal is to determine the molar concentration and strength of two acid solutions by conducting titrations with a base solution of known concentration.
Before we get to the activity, let’s take a brief look at how acid-base titrations work. There are three different ways to define acids and bases, but for the purposes of this activity, we’ll use the Arrhenius definitions. Arrhenius acids are compounds that dissociate when added to water, resulting in free hydrogen ions (H⁺) in the water. Arrhenius bases also dissociate when added to water, but these compounds yield free hydroxide (OH⁻) ions in the water. The concentration of hydrogen ions can be quantified using the pH scale using the formula pH = -log[H⁺].
When free H⁺ and OH⁻ ions are both available in an aqueous solution, they will combine to make molecules of water (H2O). This is called neutralization. When an acid and a base combine, they produce an aqueous solution with a dissolved salt made up of the cation from the base and the anion from the acid. For example, a combination of HCl and NaOH will yield an aqueous solution of sodium chloride according to the following equation:
HCl + NaOH → H₂O + NaCl
An acid-base titration is essentially a controlled neutralization reaction between one aqueous solution of a known concentration (called a titrant or standard) and one aqueous solution of an unknown concentration (called an analyte). As shown in the experimental setup below, the analyte is placed in a flask and positioned under a buret containing the titrant. The titrant is then slowly added to the analyte, one drop at a time. If the analyte is an acidic solution, the titrant must be a basic solution and vice versa in order for the two solutions to neutralize one another.
As the titrant is added to the analyte, the free H⁺ and OH⁻ions in the resulting mixture form water molecules, and the pH value of the solution changes. If the analyte is an acid and the titrant is a base, the pH of the mixture will increase (become more basic). If the analyte is a base and the titrant is an acid, the pH of the mixture will decrease (become more acidic).
There are two ways to run an experimental titration. One uses an indicator that changes color based on the pH level of the analyte-titrant mixture. The other uses a pH meter. Each method offers different information.
Using a pH meter during a titration allows you to determine how the pH level changes as the titrant are added to the analyte. When you plot this data, it generates a titration curve with a signature S or reverse S shape, as shown below. This curve allows you to identify the equivalence point of the reaction, which is the point at which there are an equal number of moles of acid and base in the mixture. The dataset you will work with in this activity is the kind you could collect with a pH meter. You could of course upload your own dataset from this activity and compete the activity questions using your own data.
The Dataset
In this activity, you’ll analyze data from two different acid-base titrations. The titrant in both titrations is a 0.50 M sodium hydroxide (NaOH) solution. The analyte in Titration 1 is 25 mL of a hydrochloric acid (HCl) solution of unknown concentration. The analyte in Titration 2 is 25 mL of an acetic acid (CH3COOH) solution of unknown concentration.
The data for this activity was generated using the ChemReaX acid-base titrations simulator.
The Activity
1) Write balanced neutralization equations for both titrations. Include phase notations.
2) Make a graph showing the change in pH over time as sodium hydroxide is added to each acid solution.
Click on the Graph tab at the top of the screen to switch to graph view. Be sure that the Scatter/Box/Bar or Categorical Bubble icon is selected; this will ensure you make a scatter plot. Click the Show buttons beneath the variable names to show the independent variable on the X-axis and the dependent variable on the Y-axis of the graph. Be sure each variable is showing on the correct axis. If it’s not, you can correct that on the panel to the right side of your graph. Next, click on the Show button under the Analyte variable and select the Z axis (on the right side panel). This will show each titration’s data with its own color. Finally, check the Connect Dots and Hide Dots boxes.
3) Use the graph to describe the trend in pH over time for each titration.
4) Use the graph to approximate the equivalence point for each titration.
Using the volume of NaOH solution that was added to the analyte upon reaching the equivalence point, calculate the approximate concentration of the analyte used in each titration. Show your work.
5) Compare the actual molarity of the two solutions (which you will need to get from your instructor) to your calculated values. What was your percent error? Show the work you did to determine this.
6) Sometimes it’s not possible to monitor the exact pH of the analyte-titrant mixture during the entire course of a titration because a pH meter is not available. In these situations, it is still possible to approximate the equivalence point of the titration by using a pH indicator solution whose color changes right around the expected equivalence point. The point at which the indicator changes color is called the endpoint. Depending on the properties of the chosen indicator, the endpoint may or may not closely match the equivalence point. Look at the following chart of indicators and choose the indicator that will change color most closely to the equivalence for each of the titrations in this activity. Justify each of your choices.
7) At what volume of added base does pH = pKₐ for the CH₃COOH solution titration represented in this activity?
The Kₐ of CH₃COOH is 1.8 x 10⁻⁵.
AP Chem Extension
8) Assume you use the same NaOH solution from this activity to titrate a 20 mL sample of 0.25 M hydrofluoric acid (HF).
The Kₐ of HF is 6.6 x 10⁻⁴. Using this information, calculate the following values:
The initial pH of the HF solution
The pH after 5.0 mL of NaOH solution is added
The pH at the equivalence point
The pH after adding 5.0 mL of NaOH solution beyond the equivalence point