2609 pH Titration of Acids and Bases

1. INTRODUCTION

A titration is a procedure for determining the concentration of a solution (the analyte) by allowing a carefully measured volume of this solution to react with another solution whose concentration is known (the titrant). In this experiment, the analyte is NaOH, and the titrant is an acid called KHP (potassium hydrogen phthalate). The point in the titration where enough of the titrant has been added to react exactly with the analyte is called the equivalence point, and it occurs when moles of titrant equal moles of analyte according to the balanced equation between the analyte and titrant. There are many types of titrations. In this experiment, you will be performing an acid-base titration.

1.1 Objectives

After completing this experiment, the student will be able to:

• determine the concentration of an NaOH solution using data from titrations involving (visual) indicators and a pH meter.
• determine which indicator(s) provides the best titration data.

1.2 Background

In general, an acid, HA, and a base such as sodium hydroxide react to produce a salt and water by transferring a proton (H+):

HA(aq) + NaOH(aq) 🡪 NaA(aq) + H2O(l) (Equation 1)

Because sodium hydroxide is hygroscopic, it draws water from its surroundings. This means one cannot simply weigh out a sample of sodium hydroxide, dissolve it in water, and determine the number of moles of sodium hydroxide present using the mass recorded, since any sample of sodium hydroxide is likely to be a mixture of sodium hydroxide and water. Thus, the most common way to determine the concentration

of any sodium hydroxide solution is by titration. Determining the precise concentration of NaOH using a primary standard is called standardization.

To find the precise concentration of the NaOH, it must be titrated against a primary standard, an acid that dissolves completely in water, has a high molar mass, that remains pure upon standing, and is not hygroscopic (tending to attract water from the air).

In this experiment you will be given the acid named potassium hydrogen phthalate (KHP). Warning: the “P” in KHP is phthalate, not phosphorus! This acid is available as a very pure solid, and therefore, it is very convenient for use in titrations because the number of moles of KHP can be accurately calculated from careful measurement of its weight.

The structure of KHP (the acidic hydrogen is circled).

After preparing your own dilute solution of NaOH, you will use it to perform several titrations with KHP. Although the dilution equation, M1V1 = M2V2, can approximate the concentration (molarity) of your NaOH from your original preparation (it will be approximate because the concentration of NaOH stock solution has only one or two significant figures), one goal of this experiment is to determine more precisely the concentration of your NaOH solution using an appropriate set of data obtained from titrations using three different acid-base (visual) indicator solutions and a pH meter.

Visual indicators change color over a relatively narrow pH range, known as the endpoint. When using a visual pH indicator, it is important to match the endpoint of the indicator with the expected pH of the equivalence point of the titration being observed. Recall from the ‘Introduction’ (Part 1.0) that the equivalence point is the point in a titration when the moles of acid and base present in the reaction match the stoichiometry of an appropriately balanced chemical equation. The endpoint is the pH when the visual indicator changes color. Not all of the indicators in this experiment have endpoints that match the equivalence point!

The following three indicators will be used in this experiment. The color changes to expect when going from an acidic to a basic solution (i.e., increasing pH) are:

When using a pH meter, during a titration, the pH of the analyte solution will be recorded after each increment of a specific volume of the titrant solution is added. After which, a graph of the pH of the analyte solution vs the volume of NaOH added will be created. This is called a titration curve. (Figure 1) To find the exact volume of NaOH needed to reach the equivalence point, a tangent line that touches the steepest part of the titration curve is drawn. The equivalence point is exactly halfway between the two points where the titration curve deviates from the tangent line.

Figure 1. Titration Curve: pH of analyte vs volume of NaOH added during titration. (Warning: this graph is for illustration purposes only. Your actual data may differ.)

References and further reading

Technique G: Buret Use

Experiment 2501 Using Excel for Graphical Analysis of Data of the laboratory manual

2. SAFETY PROCEDURES AND WASTE DISPOSAL

3.0 CHEMICALS AND SolutionS

Chemical	Concentration	Approximate Amount	Notes
Potassium hydrogen phthalate (KHP)	Pure	2.5 g
NaOH solution	6 M	10 mL
Phenolphthalein solution		Several drops
Bromothymol blue solution		Several drops
Methyl red solution		Several drops

4.0 GLASSWARE AND APPARATUS

pH meter (fragile!)	50 mL buret
125 mL or 250 mL Erlenmeyer flask	Buret clamp
10 mL graduated cylinder	Ring stand
600 mL beaker	Funnel (to fill buret)

5.0 PROCEDURE

5.1 PREPARATION OF DILUTE NaOH Solution

Into a clean 600 mL beaker (or another container), add approximately 500 mL of laboratory water. Then add about 10 mL of 6M NaOH stock solution and mix thoroughly. Note that it is not necessary to accurately measure volumes at this time because the final concentration of your dilute NaOH solution will be calculated from titration data. Label and cover the beaker with a watch glass.

Be careful not to contaminate your dilute NaOH solution at any time during the lab as this may change its concentration.

5.2.0 TITRATION OVERVIEW WITH HELPFUL TIPS

In this experiment, you will collect six (6) sets of titration data:

• two (2) titrations using phenolphthalein
• two (2) titrations using bromothymol blue
• one (1) titration using methyl red
• one (1) titration using a pH meter

Each titration requires a different amount of KHP; however, the choice of indicator and/or use of the pH meter can be in any order. Lab partners should alternate so that every student has an opportunity to use the buret to complete at least one titration.

Tip #1: To save time, you may use any one indicator and the pH meter at the same time to collect both sets of data simultaneously. Discuss this option with your instructor.

Tip #2: If pH meters are in short supply, some lab groups should use a pH meter at the beginning of the experiment whereas other groups should plan to use a pH meter at a later time.

Tip #3: To speed up your titrations, initially add large amounts (1-2 mL at a time) of solution from the buret. Then, slow down when you are 2-3 mL before the equivalence point and add solution from the buret dropwise. Estimate the equivalence point of each titration by using the balanced chemical equation, the mass of KHP, and an estimate of the concentration of your NaOH solution obtained from the dilution equation (see Part 7.0 DATA ANALYSIS below for a step-by-step process). Note: Different amounts of KHP are used for each titration which means the equivalence point will be different for each titration!

Warning: Do not share titration data with other lab groups because each group has an NaOH solution that is slightly different from everyone else. Remember that the goal is to accurately determine the concentration of NaOH solution in your group!

5.2.1 TITRATION PROCEDURE WITH INDICATORS

1. Obtain a buret and make sure that it is clean and does not leak. If necessary, clean the buret with a buret brush and soapy water and rinse with laboratory water. Then, rinse the buret several times with a few milliliters of your NaOH solution, making sure that the stopcock and buret tip are also thoroughly rinsed with NaOH. (Why rinse with NaOH? To remove water that would otherwise dilute and change the concentration of your NaOH solution!) Fill the buret with your NaOH solution making sure that there are no air bubbles and no leaks in the stopcock or tip. Record the initial volume of the buret (reading the meniscus to 2 decimal places).

2. The buret does not need to be cleaned or rinsed between titrations. Simply refill the buret with NaOH prior to beginning a new titration. Be sure to record the new initial volume of the buret (reading the meniscus to 2 decimal places).

3. Clean a 125 mL or 250 mL Erlenmeyer flask and rinse well with laboratory water. The flask may remain wet.

4. On a weighing paper, weigh about 0.5 g KHP and record the exact mass in the column of the data table for Phenolphthalein #1. Completely transfer the KHP to the flask, and then add approximately 50 mL of laboratory water and 2-3 drops of the phenolphthalein indicator. Swirl to mix thoroughly and dissolve the KHP completely.

5. (Optional) Obtain an estimate of the equivalence point by completing the data table in 7.0 DATA ANALYSIS below. If you know where the endpoint is, you will have a better chance of not passing it during the titration. This step (Step 5) is optional if you want to devote more lab time to careful titrations. If you choose to skip this step now, be sure to complete these calculations after the lab in order to earn all points on the lab write-up.

6. Begin the titration by carefully opening and closing the buret stopcock to allow the NaOH solution to drain into the KHP/phenolphthalein solution while swirling the flask. Stop the flow of NaOH frequently to assess the titration by observing its color. You may wash the sides of the flask with your squirt bottle containing laboratory water during the titration. As the endpoint approaches, the indicator may change color from colorless to light pink initially as NaOH is added. However, the solution may change back to colorless as you swirl the flask. Slowly continue the titration if the new color does not persist. Tip: place a sheet of white paper under the titration flask to help you detect the faintest of pink color. As the endpoint gets closer, add NaOH one drop at a time while swirling the reaction mixture well before adding another drop. Stop the addition of NaOH as soon as one drop causes the solution to change permanently to the new color (about 30 seconds) — this is the endpoint! Record the final volume of the buret (reading the meniscus to 2 decimal places).

Tip: Identifying the endpoint is not easy. When you suspect that you have reached the endpoint, proceed to record the final buret volume. Then, add one more drop of NaOH and assess the titration. If the color does not change or if the color becomes darker without changing hue, then the endpoint was reached before that drop was added and your recorded volume was correct. But if the color changes, then the true endpoint has now been reached—cross out the previously recorded buret volume and write in the new final volume of the buret. You may repeat this process if you are still uncertain of the endpoint.

7. Repeat the titration procedure (steps 2-6) to acquire all sets of titration data with the other indicators, using different quantities of KHP as follows:

Phenolphthalein #2: Use about 0.7 g KHP (record exact amount) and 2-3 drops of phenolphthalein. The color change for this indicator is colorless 🡪 pink. Stop the titration when the solution turns very light pink.

Bromothymol blue #1: Use about 0.3 g KHP (record exact amount) and 2-3 drops of bromothymol blue. The color change for this indicator is yellow 🡪 green 🡪 blue. Stop the titration when the solution turns green.

Bromothymol blue #2: Use about 0.6 g KHP (record exact amount) and 2-3 drops of bromothymol blue. The color change for this indicator is yellow 🡪 green 🡪 blue. Stop the titration when the solution turns green.

Methyl red: Use about 0.4 g KHP (record exact amount) and 2-3 drops of methyl red. The color change for this indicator is red 🡪 orange 🡪 yellow. Stop the titration when the solution turns orange.

Note: If you think that you overshot the endpoint of any titration, you may repeat that titration at the end of the lab, if time permits.

5.2.2 TITRATION PROCEDURE WITH pH METER

Warning: The pH meter probe (end or tip) is extremely fragile! Do not allow the tip to touch the flask or other solid object.

Complete the titration as described in Part 5.2.1 steps 2-6 with the following modifications:

1. Use a pH meter instead of an indicator (or in addition to an indicator, as mentioned in Part 5.2.0)

2. Use an Erlenmeyer flask with an opening large enough for both the pH meter and the buret tip.

3. Weigh about 0.4 g KHP and record the exact mass.

4. The procedure below describes how to use a portable Flinn pH meter (model AP8673). If the pH meter is not measuring properly (unstable, or inaccurate), refer to the manual for troubleshooting. This same procedure may be applied to other brands of pH meter.

1. Check out a portable pH meter from the stockroom.
2. Remove the protective cap on the electrode. Clean any salt build-up off by rinsing with laboratory water.
3. Press the ON/OFF button once.
4. Rinse the electrode with laboratory water and blot dry with filter paper.
5. Immerse the electrode in the flask containing the analyte solution. Once the display stabilizes (approx. 1 min.), record the exact pH.
6. Remove the pH meter from the solution.
7. Repeat steps (d) – (g) after a certain volume of the titrant is added.

5. During the titration, record both the buret volume and the pH measurement periodically: every 1 mL increment when the pH is lower than 8.0, every 1-2 drops when the pH is between 8 and 10, and every 1 mL increment when the pH is higher than 10.

6. Continue the titration after the endpoint, recording both the buret volume and the pH measurement in 1 mL increments until the pH is higher than 12.5 or the NaOH solution in the buret reaches the 50 mL mark.

7. When finished, rinse the pH meter electrode with laboratory water and blot dry with filter paper. Replace cap and return the pH meter to the stockroom.

6.0 DATA RECORDING SHEET

Last Name

First Name

Partner Name(s)

Date

Table 1. Titrations with Indicators

	Phenolphthalein #1	Phenolphthalein #2	Bromothymol blue #1	Bromothymol blue #2	Methyl Red
Mass of weighing paper
Mass of KHP + weighing paper
Mass of KHP
Moles of KHP
Initial buret reading
Final buret reading
Total NaOH volume

Table 2. Titration with pH Meter

Mass of weighing paper
Mass of KHP + weighing paper
Mass of KHP
Moles of KHP

Volume of NaOH added (Buret reading)	pH reading

Volume of NaOH added (Buret reading)	pH reading

1. DATA ANALYSIS

1. Estimate the titration equivalence point by completing the following data table using the balanced chemical equation, the mass of KHP, and an estimate of the concentration of your NaOH solution obtained from the dilution equation (see Parts 1.0 Introduction and 1.2 Background).

	Phenolphthalein #1
Mass of KHP
Moles of KHP (molar mass = 204.2236 g/mol)
Moles of NaOH required to react in a 1:1 ratio with KHP
Approximate concentration of your NaOH solution Use the dilution equation M1V1 = M2V2 and find the unknown M2 when M1 = 6 M (or use the concentration written on the bottle of 6 M NaOH), V1 = 10 mL, and V2 = 500 mL.
Total NaOH volume in liters required for titration (also the estimate for the equivalence point). Divide Moles of NaOH by the Approximate concentration of your NaOH solution
Convert liters to mL by multiplying total NaOH volume in liters by 1000

2. Using a spreadsheet program such as Microsoft Excel or Google Sheets, construct a graph of pH (y-axis) vs. volume of NaOH added in mL (x-axis).

• Attach the graph to your lab report.
• Determine the equivalence point (in mL of NaOH) of this pH titration.

3. Using a spreadsheet program such as Microsoft Excel or Google Sheets, construct a graph of volume of NaOH solution in mL at the endpoint of each titration (y-axis) vs. mass of KHP for each corresponding titration (x-axis).

• Attach the graph to your lab report.
• Determine the line of best fit (linear regression line) and in the space below, write the equation of this line in the form y = mx + b.

4. In the space below, describe how well your data points match the line of best fit (from #3 above). Do one or more points appear to be outliers? For these outliers, propose a reason why they do not match the line of best fit: experimental error/inaccurate use of the buret or perhaps the indicator’s endpoint poorly matched the titration equivalence point?

5. What are the advantages and disadvantages of using a pH meter instead of a visual indicator?

6. Calculate the concentration of your NaOH solution using each set of titration data.

• Fill in the following table with your results.
• For one set of titration data, show the calculations and/or explain how you obtained each of the values you entered in that row.
• Comment on the quality of all your titrations: Do they all give reliable results, or should one or more titrations not be included in the average concentration of NaOH? Explain.

Table 2: NaOH Concentration Calculation

Titration data from:	Moles of KHP	Moles of H+	Moles of NaOH	Liters of NaOH solution	Concentration of NaOH
Phenolphthalein #1
Phenolphthalein #2
Bromothymol blue #1
Bromothymol blue #2
Methyl Red
pH meter

Table note: Remember, KHP is the monoprotic acid we are using to standardize the NaOH. Litres of NaOH should be the same in this table as total volume added of NaOH from Table 1.

8.0 POST-LAB QUESTIONS

1. Draw the balanced chemical equation for the titration reaction using Lewis Dot structures (for example, instead of “NaOH” in your balanced chemical equation, draw “Na+ -O-H” and include 3 lone pairs of electrons around the oxygen atom in your drawing). Do not include indicator molecules.

2. What volume of 0.812 M HCl is required to titrate 1.33 g of KOH to the endpoint? Show your work (attach a separate sheet of paper, if necessary).

3. What volume of 1.346 M H2SO4 is required to titrate 1.54 g of KOH to the endpoint? Show your work, including the balanced chemical equation and a short explanation on how you used the balanced equation when calculating for the endpoint.