ABT: Acid-Base Titration

 Purpose: Titration is an important quantitative analysis technique where a solution of accurately known concentration (the titrant) is added incrementally (usually by a buret) to another solution of unknown concentration (the analyte) until the chemical reaction between the two solutions is complete (i.e. no more analyte remains).  This is known as the equivalence point.  The titrant concentration and volume used to reach the equivalence point, along with the volume of the analyte titrated are used to calculate the concentration of the analyte.  This technique is commonly used in biomedical diagnostic labs.  Since there is often no visible change at eh equivalence point, an indicator must be added.  Indicators may involve a change of color near the equivalence point, pH (if the reaction is an acid/base), or some other physical property.  This lab will familiarize you with the following:

      1. The titration procedure between an acid and a base. The reaction completion point will be determined through the use of both a color indicator and a pH meter.
      2. Determining the concentration of an unknown solution by titration.
      3. The use and care of a buret and pH meter.
      4. Calibration of a pH meter.
      5. The proper handling procedures for dilute acid and base solutions.
      6. The use of a spreadsheet to graph a set of data.

 Background reading:

CHEM 107 Web Page

            Graphing with Excel

Cooperative Chemistry Laboratory Manual

Page 45-46                                          Containers

Pages 47-48                                         Measuring Liquids by Volume

Pages 67-69                                         Titration

Page 91                                                pH Meter- Its Care and Use

Chemistry & Chemical Reactivity

Section 5.3-5.4                                    Acids and Bases (pp 185-193)

Sections 5.8                                         Measuring Concentration of Compounds in Solution (pp 205-208)                   

Section 5.9                                           pH, a Concentration Scale for Acids and Bases (pp 212-214)

Section 5.10                                         Stoichiometry of Reactions in Aqueous Solution (pp 214-217) 

 

Chemical Principles, The Quest for Insight

            Pages F72-F76                                                Acids and Bases

            Pages 402-405                                                 The pH Scale

           

Equipment and Materials:

You will work in-groups of two for this experiment. Each group will need the following:

      1. One buret (obtain from the common equipment area)
      2. One stand (common equipment cabinet next to your locker.)
      3. One buret clamp (common equipment area)
      4. One 25-mL volumetric pipet (common equipment area)
      5. 100 mL of the unknown acid and 100 mL of the NaOH solution (common equipment area)
      6. A pH meter (The pH meters are already set up in the lab. Do not move the meters. Take your materials and equipment to the meter.)
      7. Two Erlenmeyer flasks and two beakers.
      8. Buffer solutions pH = 4.00, 7.00 and 10.00 (Use the buffer solutions in the jars that are provided. Do not pour the buffer solutions into another container.) (common equipment area)

Notes on titrating and using burets:

  1. Burets are expensive devices and must be used carefully. Do not try to clean a buret by putting it under the faucet. To clean a buret, securely clamp it to a stand, and using funnel pour water into it from the top of the buret. NEVER TRY TO POUR LIQUIDS INTO AN OPENING ABOVE YOUR HEAD.
  2. The burets used in this experiment are 50-mL burets, accurate to 2 decimal places. Remember from your reading assignment that a buret measures the quantity of liquid dispensed, and thus the zero mark is at the top of the buret.
  3. Review how to read a meniscus from last week so that you will be able to properly read the liquid volume in a buret to calculate the volume delivered.

Chemicals:

Dilute hydrochloric acid solution

Dilute sodium hydroxide solution

Phenolphthalein indicator

Distilled water

Buffer solution

 Safety Precautions:

Sodium hydroxide and hydrochloric acid solutions are corrosive chemicals and skin-irritants. If you splash either on your skin, rinse the area with lots of running water. You must wear your safety goggles at all times. Remember: even though you may be finished in lab does not mean everyone else is finished and hence an accident could still happen. Any spills must be cleaned up immediately. Wash everything you used that contained sodium hydroxide and/or hydrochloric acid with water before putting it away. Remember, chemical residues may linger for weeks.

 Overview:

The purpose of this lab will be to determine the concentration of a dilute hydrochloric acid solution. In order to do this, you will use an experimental technique called titration. A titration involves using two solutions: first the "titrant" which in this case will be a solution of sodium hydroxide of accurately known concentration, and the second solution, refereed to as the "unknown", is a solution of hydrochloric acid of unknown concentration. The sodium hydroxide solution is placed in a buret and added incrementally to the unknown acid solution until the moles of NaOH equals the moles of HCl.  This is called the equivalence point.  The overall reaction is: NaOH + HCl ® NaCl + H2O.  At the equivalence point, the solution contains only sodium chloride and water.  By knowing the volume of the unknown acid solution used in the titration and the concentration and volume of NaOH added to the unknown acid solution, the concentration of the unknown acid solution may be calculated (M = moles acid/volume solution). Now, all we need is a method to determine when the reaction is complete, i.e. when the equivalence point is reached. 

Two techniques will be used to monitor the reaction.  First, an acid/base indicator can by used to visually monitor the end point of the reaction.  Phenolphthalein is an indicator that is colorless in acidic solutions and pink in basic solutions. Therefore, you will monitor the reaction between NaOH and HCl by monitoring the color change that takes place (as the reaction changes from an acidic solution to a basic solution). When the amount of OH- supplied by the NaOH equals the amount of H+ supplied by the acid being analyzed, the solution changes color.  Secondly, a pH meter can be used to monitor the reaction. As you add the NaOH to the HCl, you will monitor the reaction for changes in pH. You will record the volume of titrant added and the pH of the unknown solution. Once sufficient data has been collected, you will use Excel to generate a titration curve and then you will be able to determine the concentration of the unknown acid.  A titration curve is generated by plotting pH vs. volume of NaOH added.  A titration curve is not linear showing an abrupt change in the pH at the equivalence point.  The pH at the equivalence point, the point when equal numbers of moles of NaOH and HCl have reacted, is taken as the mid-point in the vertical portion of the curve.  Once the equivalence point is found, the volume of NaOH added is obtained from the graph.

 Experimental Procedure:

In this experiment, you will be using a pH meter.  A pH meter must be calibrated before it can be used to collect data.  Be sure your pH meter is calibrated correctly before collecting data. 

Calibration of a pH Meter.

  1. Obtain the buffer solutions from the common equipment area.
  2. Directions for using the pH meters are provided with each meter.
  3. If you have problems using the pH meter, ask you instructor for help.
  4. Calibrate the pH meter with buffers 4.00 & 7.00 and then use buffer 10.00 to check the calibration.
  5. Each pH meter has an electrode that will need to be cleaned often with distilled water as outlined in the directions for the pH meters. Use a wash bottle with distilled water to rinse the electrode. Do not take the electrode to the sink to rinse.
  6. When you are finished with it for the day, rinse and dry the electrode and return the electrode to the KCl solution for storage.

 Acid-Base Titration using a pH meter.

  1. Clean the appropriate clean, dry glassware for the experiment.
  2. Obtain 100 mL of the NaOH solution and 100 mL of the unknown acid solution and label the beakers.
  3. Fill the buret with the NaOH solution to just below the 0.00-mL mark.  Be sure the tip of the buret is filled with NaOH. Read the initial buret volume and record the volume in the appropriate space in the data sheet.
  4. Using a clean 25-mL volumetric pipet, transfer 25-mL aliquots of the unknown acid to two beakers. Add about 75 ml of distilled water and 2 to 3 drops of the phenolphthalein indicator to each beaker. The volume of water is not critical.
  5. Measure the pH of this solution.
  6. Titrate the acid, collecting sufficient data (volume of titrant added and pH of the solution) to plot a titration curve. (Refer to Figure 18.4, page 863 of your textbook).  
  7. You will be monitoring two factors: pH and the color of the solution.  First, you will need to record the total volume of NaOH that was added for each pH measurement.  Secondly, by adding the phenolphthalein you will be able to monitor the endpoint of the reaction visually.  Therefore, you will need to record the total volume of NaOH that was added when the solution changed from a colorless solution to a very pale pink color.
  8. Repeat the procedure for the second sample.
  9. When you are finished with this step, it is a good idea to generate your titration curve before continuing.

 

Data Analysis:

A. Calculation of the Molarity of Unknown Acid

 Example:

 Five mL of HCl (aq) was titrated to the endpoint with 23.78 mL of a 0.1000-M solution of NaOH. Calculate the molarity of the HCl solution.

 Solution: Start with the balanced equation for the reaction between hydrochloric acid and sodium hydroxide:

 HCl(aq) +

NaOH(aq) ®

NaCl(aq) +

H2O(l)

Unknown Concentration

Known Concentration

 

 

Known volume

Known volume

 

 

 

Since the unknown acid was titrated to the endpoint (the point at which the unknown acid has been completely consumed) with the base and there is a one to one mole ratio between HCl and NaOH in the balanced equation, it follows that the moles of HCl in the unknown that reacted equals the moles of NaOH added from the buret.

1.  Calculate the number of moles of NaOH added to the HCl

2.  Use the balanced equation to determine the number of moles of HCl that was titrated.

3.  Calculate the concentration of HCl in moles/L.

1.  (0.1000  mol NaOH/L)(0.02378 L NaOH) = 0.002378 mol NaOH

2.  (0.002378 mol NaOH)(1 mol HCl/1 mol NaOH) = 0.002378 mol HCl

3.  MHCl = 0.002378 mol HCl / 0.00500 L HCl) = 0.4756 M

 

 *To summarize, if we need to find the concentration of an unknown acid, titrate a known volume of the unknown acid with a known concentration of base. Measure the volume of base required to get to the endpoint, which is monitored by using a color indicator or a pH meter.

  1. Creating a Titration Curve

 Set up a spreadsheet with your data from Sample 2 only as follows:

 

 

A

B

1

mL NaOH added

pH

 

  1. Plot a titration curve with mL NaOH added as the x-axis and pH as the y-axis.  This is a non-linear plot, so it does not require a trendline.  However, connect the data points.

(The endpoint is volume of base added where the slope of the pH curve is greatest.)

  1. Recalculate the concentration of the acid, as described above, using the pH curve endpoint. Generally, titration curves are more accurate in determining the endpoint, but require the use of a pH meter and generating a graph. Modern labs use computers and robot arms to generate titration curves.

 

 Data Sheet 1

 Data Sheet 2