7 .4 An Introduction to Titrations

Warm Up

1. Balance the following neutralization equation: ___ H2SO4(aq) + ___ NaOH(aq) ➝ _____H2O(l) + ____ Na2SO4(aq)

2. Write the balanced equation for the reaction between aluminum hydroxide and hydrobromic acid to form aluminum bromide and water.

3. Complete and balance the following equation: NH4OH(aq) + H2SO4(aq) ➝ _______________ + _______________

4. Write the formulas for the acid and base that will react to give the salt K2CO3 and water.

Titration

A titration is a quantitative analysis method used to determine the concentration of an unknown solution by reacting it with another substance of known concentration. The most common types of titrations make use of an acid-base neutralization reaction or a reduction-oxidation reaction. In this section, we will focus on acid-base titrations. In a neutralization reaction, an acid reacts with a base to form a salt and water:

HCl(aq) + NaOH(aq) ➝ NaCl(aq) + H2O(l) acid + base ➝ salt + water

Solving Titration Problems

During this course, you have learned how to perform calculations based on balanced chemical equations. A titration problem is essentially a simple stoichiometric calculation. For any titration, the balanced chemical equation is a good starting point. In a titration, one reactant completely uses up the other reactant. By knowing the amount of one substance used in a reaction, the amount of the unknown substance is calculated by using the mole ratio from the balanced equation.

Sample Problem — Calculating the Unknown Concentration of an Acid

A student completely reacted 10.00 mL of HCl with 18.25 mL of 0.100M NaOH. Calculate the [HCl].

What to Think about

1. Write a balanced chemical equation for the acid-base reaction.

2. Calculate the moles of substance that you know.

3. Using the balanced equation, convert the amount of substance you know to the amount of substance you don’t know.

4. Calculate the concentration by dividing the moles of HCl by liters of HCl.

The same type of calculation can be used to calculate the unknown concentration of the base.

How to Do It

HCl(aq) + NaOH(aq) ➝ NaCl(aq) + H2O(l) moles NaOH = 0.100 mol NaOH 1 L NaOH × 0.01825 L NaOH = 0.001825 mol NaOH moles HCl = 0.001825 mol NaOH × 1 mol HCl 1 mol NaOH = 0.001825 mol HCl

[HCl] = 0.001825 mol HCl 0.01000 L HCl = 0.183 M

Practice Problems — Simple Titration Calculations

1. A student titrated 25.00 mL of HCl with 15.62 mL of 0.30M NaOH. Calculate the [HCl].

2. In a titration, a 10.00 mL sample of NaOH was titrated with 24.25 mL of 0.20M H2SO4. Calculate the [NaOH].

3. A student used 22.68 mL of 0.015M Sr(OH)2 to titrate 5.00 mL of HNO3. Calculate the [HNO3].

4. What volume of 0.50M H2C2O4 is required to titrate 10.00 mL of 0.12M NaOH?

(a)

Practical Aspects of Titration

In a titration, there are two solutions. The solution whose concentration is known is called the standardized solution. The other solution is of unknown concentration. Since titration is a quantitative method of analysis, the precision of the measuring tools used will determine how precise your final answer will be. Recording your data to the correct number of significant figures, and paying close attention to significant figures in calculations is very important in a titration. Special glassware enables accurate titration. It is important to know the names of this glassware.

• Burette (Figure 7.4.1): contains the standardized solution. The numbering on the burette allows us to measure the initial volume before the titration and final volume after the titration. Subtracting these two volumes gives us the volume of standardized solution added to the Erlenmeyer flask. Notice the scale reads from 0.00 mL at the top to 50.00 mL at the bottom.

(a) (b) Figure 7 .4 .1 (a) A burette (b) The reading on this burette is 42.30 mL.

• Pipette (Figure 7.4.2): used to measure and deliver a precise volume of the solution of unknown concentration to the Erlenmeyer flask. Pipettes come in a number of sizes: 1.00 mL, 5.00 mL, 10.00mL, 25.00 mL, etc. They may be graduated like a burette, or designed to deliver a specific volume. A pipette bulb is used to suck the solution up into the pipette. Never suction the solution by mouth.

• Erlenmeyer Flask (Figure 7.4.3): contains the measured volume of solution of unknown concentration and a few drops of an indicator at the beginning of the titration. As the titration proceeds, solution from the burette is added drop by drop to the solution in the Erlenmeyer flask where the reaction takes place. The indicator changes color when the reaction between the moles of acid and moles of base is complete.

stand

• The equivalence point occurs when the moles of H+ equal the moles of OH–. When the indicator changes color, the endpoint is reached. This is when you record the final burette reading.

Figure 7 .4 .4 A titration apparatus

1. How is a pipette different from a burette?

2. What solution goes into the burette?

3. Which piece of glassware is the indicator added to?

4. What is “equivalent” at the equivalence point?

7.4 Activity: Titrating CH3COOH with NaOH

Question

What is the unknown concentration of a sample of CH3COOH?

Background

In previous science classes, you learned about chemical indicators and how they are used to test for the presence of acids or bases. An indicator will be one color in the presence of an acid, and a different color in the presence of a base. One such indicator is phenolphthalein. In this activity you will see how phenolphthalein can be used to indicate when an acid has completely reacted with a base. When there is an excess of acid, the phenolphthalein will be colorless. Once the acid has been completely used up and there is an excess of base, the phenolphthalein will turn pink. This color change occurs within one drop of NaOH being added.

Procedure

1. The steps involved in a titration of CH3COOH using NaOH are described below. Answer the questions following each step to help you to understand what is happening at each part of the titration.

2. Discuss each step with your partner.

Results and Discussion

1. Write a balanced equation for the reaction between CH3COOH(aq) and NaOH(aq).

2. Consider a solution of CH3COOH of unknown concentration. It was titrated with a standardized solution of 0.15M NaOH. A student used a pipette to deliver 10.00 mL of CH3COOH to the Erlenmeyer flask. A few drops of phenolphthalein were added. What would be the color of the solution in the flask? ______________

3. The standardized 0.15 M NaOH was placed in the burette. The initial reading on the burette was 0.50 mL. A few drops of NaOH from the burette were allowed to drip into the flask. The solution in the flask did not change color. Why does the solution in the flask not change color when the NaOH is added?

4. After adding more NaOH drop by drop, the solution became pink. At this point, no more NaOH was added. The final reading on the burette was 12.30 mL. What volume of NaOH was added to the flask? _____________________

5. When the color of the indicator changed, the endpoint had been reached. Why did the color of the solution become pink?

6. Once the titration was completed, the contents of the flask were rinsed down the drain, and the flask was rinsed out with water. To ensure that the data was accurate, a second and third titration was completed. We call these trials. An example of the data collected would look like the table shown below.For each trial, complete the subtraction and calculate the volume of NaOH added. The average volume of NaOH is calculated by averaging the two trials within 0.1 mL. Fill in the table with the appropriate values.

7. Using the concentration of NaOH and average volume of NaOH, calculate the moles of NaOH used in the titration.

8. Use the mole ratio from the balanced equation to calculate the moles of CH3COOH present in the flask.

9. Calculate the molarity of the CH3COOH by dividing the moles of CH3COOH by the volume of CH3COOH in the flask.

10. Compare your answer to that of another group.

7.4 Review Questions

1. Describe how you can tell from a formula if a substance is an acid or a base. Use examples.

2. Complete and balance the following acid-base neutralization equations:

(a) HI(aq) + LiOH(aq) ➝

(b) Ca(OH)2(aq) + HNO3(aq) ➝

(c) CH3COOH(aq) + NH4OH(aq) ➝

(d) H2SO4(aq) + KOH(aq) ➝

3. A student titrates a standardized solution of H2C2O4 to determine the concentration of NaOH. The indicator used was phenolphthalein. Draw an apparatus used for the titration. Label all glassware and their contents. Describe what color the solution would be in the flask at the beginning of the titration and at the end of the titration.

4. Define the terms “standardized,” “equivalence point,” and “endpoint.”

5. A student titrated 10.00 mL HCl with 0.050 M Sr(OH)2. The table below shows the data collected. Calculate the [HCl].

6. A 25.00 mL sample of 0.20 M H2CO3 was titrated with 0.50 M NaOH. What volume of NaOH was required to reach the equivalence point?

7. A 10.00 mL sample of vinegar (CH3COOH) was titrated with 18.20 mL of 0.50 M NaOH. Calculate the [CH3COOH].

8. In order to standardize a solution of NaOH, 0.18 g of oxalic acid, H2C2O4∙ 2H2O, was dissolved to make 250.00 mL of solution. A 25.00 mL sample of this solution was titrated against 15.25 mL NaOH. Calculate the [NaOH].

9. Aspirin has the formula C9H8O4. Only one of the H atoms is released when it acts as an acid. An aspirin tablet with a mass of 0.50 g was dissolved in water, and titrated with 18.30 mL of 0.10 M NaOH. Calculate the mass of aspirin in the tablet, and then the percent by mass of aspirin in the tablet.

10. A 250.00 mL sample of Ca(OH)2 was titrated with 7.25 mL 0.10 M HCl. Calculate the mass of Ca(OH)2 present in the solution.

11. A student dissolved 0.1915 g of an unknown acid HA in 10.00 mL of water. This solution was then titrated with 0.100 M NaOH. The table below shows the data collected. Calculate the molar mass of the acid HA.