## Lab 7 Introduction

### Learning Objectives

By the end of this section, you will be able to:

• Calculate the concentration of a secondary standard through titration with a primary standard.
• Titrate a sample of carbonated soda with a standard solution to determine the concentration of acid in the beverage.
• Distinguish between endpoint and equivalence point with the addition of indicator to your sample.

## Introduction

### Neutralization Reactions

Acid –Base reactions are an important portion of chemistry. In aqueous solutions, a compound that produces H+ ions upon dissolution is termed an acid. A compound that produces OH– ions when dissolved in water is called a base. The reaction of an acid and base is a neutralization reaction, the products of which are a salt and water. This is shown below:

Equation 1

$\text{HCl}_{(aq)}\text{+NaOH}_{aq}\rightarrow\text{NaCl}_{aq}+\text{H}_{2}\text{O}_{(I)}$

In an aqueous solution, virtually all of the OH– ions present will react with all of the H+ ions that are present. Thus the net ionic equation of a strong acid reacting with a strong base is:

Equation 2

$\text{HCl}_{(aq)}+\text{OH}^{-}_{aq}\stackrel{\rightarrow}{\text{H}_{2}\text{O}}\text{H}_{2}\text{O}_{(I)}$

This reaction is essentially quantitative, which allows us to use these titrations to determine the concentration of an acid or base in an aqueous solution with high accuracy.

It is important to note that not all hydrogen atoms are considered acidic. Only acidic protons will disassociate from the compound and contribute to the reaction. Monoprotic acids produce one acidic proton per molecule. An example of a monoprotic acid is hydrochloric acid (HCl) which disassociates according to:

Equation 3

$\text{HCl}_{(aq)}\stackrel{\rightarrow}{\text{H}_{2}\text{O}}\text{H}^{+}_{(aq)}+\text{CL}^{-}_{(aq)}$

As shown in Equation 1, hydrochloric acid reacts with sodium hydroxide in a one to one ratio. We can tell from this balanced equation that when a solution of hydrochloric acid, HCl, is exactly neutralized with a solution of sodium hydroxide, NaOH, the number of moles of NaOH used will equal the number of moles of HCl originally present. This is also called the equivalence point.

Equation 4

$\text{mol}_\text{NaOH}=\text{mol}_\text{HCl}$

For this experiment we can also expand this equation to

Equation 5

$\text{M}_\text{NaOH}\text{V}_\text{NaOH}=\text{M}_\text{HCl}\text{V}_\text{HCl}$

where M = concentration in molarity and V= volume in liters. If three of the above quantities are known, the fourth can be calculated.

### Classification of Acids

Not all acids are monoprotic. Diprotic acids produce two acidic protons in aqueous solution. We will be titrating a soda that contains citric acid, a triprotic acid. Acids are often added to carbonated beverages as an ingredient to give a sour “bite” or taste. Dark sodas typically add phosphoric acid while light colored and clear sodas tend to use citric acid. Citric acid is a triprotic acid which contains three acidic protons.

Citric acid reacts with sodium hydroxide according to the following reaction where A is assumed to be the citrate ion:

Equation 6

$\text{H}_{3}\text{(aq)}+3\text{NaOH}_{aq}→\text{Na}_{3}\text{A}_{aq}+3\text{H}_{2}\text{O}_{(I)}$

Notice that because citric acid is a triprotic acid the mole to mol ratio between the acid and base is 1:3 respectively. This will be important for your calculations.

In order to determine when a solution has been exactly neutralized, an acid-base indicator is used that changes color in a certain pH range (pH is a scale used to measure acidity). This color change is termed the endpoint of the titration. (Ideally we would detect the equivalence point, but no indicator changes color at exactly the equivalence point, so instead titrations usually proceed 1-2 drops past the equivalence point to the endpoint which is when you can detect a color change). Because the pH of a neutral solution is 7, an indicator that changes color near this pH should be used for an acid-base titration. Phenolphthalein indicator changes color in the range pH = 8.3 – 10.0 and can be used to determine when the correct amount of base has been added to an acidic solution to exactly neutralize it.

### Preparation of Sodium Hydroxide Solution

You will need ~ 200 mL of 0.04 M NaOH for this experiment. You will use a small amount of 2.5 M NaOH to make the more dilute solution for use in the procedure. In your pre-lab, you will calculate how to prepare this solution.

### Standardization of a Sodium Hydroxide Solution

It is difficult to accurately prepare a solution of sodium hydroxide since it is hygroscopic (absorbs water readily from air). Thus the solution of NaOH must be standardized (meaning the exact molarity must be determined) using a primary standard. A primary standard must satisfy the four following criteria:

1. Can be obtained as a solid compound that is not hygroscopic and can be easily handled.
2. Is available in very pure form.
3. Is chemically stable over time.
4. Has a medium to high molecular weight.

We will be standardized using potassium hydrogen phthalate,(HKC8H4O4) which is a monoprotic acid. The molecular weight of (HKC8H4O4) is 204.23 g/mole, and it has one acidic proton, which will react quantitatively with OH: Occasionally potassium hydrogen phthalate is abbreviated as KHP even though the formula is HKC8H4O4.

Equation 7

${\text{HKC}_{8}\text{H}_{4}\text{O}_{4}}_{(aq)}+\text{NaOH}_{(aq)}$→${\text{NaKC}_{8}\text{H}_{4}\text{O}_{4}}_{(aq)}+\text{H}_{2}\text{O}_{(I)}$

OR

${\text{HKC}_{8}\text{H}_{4}\text{O}_{4}}_{(aq)}+\text{OH}^{-}_{(aq)}$→${\text{KC}_{8}\text{H}_{4}\text{O}_{4}^{-}}_{(aq)}+\text{H}_{2}\text{O}_{(I)}$

For the highest accuracy,a sample size is chosen such that it will consume as large a volume of the base as possible without exceeding the capacity of the buret.If a 50 mL buret is used,the amount of KHP is chosen such that it will require approximately 15 mL of 0.04 M NaOH solution to reach the endpoint (so that ~ 3 trials can be completed without refilling the buret). Thus,

0.015 L NaOH ${{\frac{{0.04}\text{ NaOH}}{{1}\text{L}}}\text{ } {\frac{{1}\text{ mol KHP}}{{1}\text{ mol NaOH}}}\text{ }{\frac{{204.23}\text{ g KHP}}{{1}\text{ mol KHP}}}}= {0.122}\text{ g KHP}$

About 0.122 g of is needed. At the endpoint, the number of moles of NaOH equals the number of moles of KHP used:

Equation 8

$\text{M}_\text{NaOH}={\frac{\text{mol KHP}}{\text{ L NaOH ( )}}}$

Equation 9

$\text{M}_\text{NaOH}={\frac{\text{g KHP}}{{204.23}\text{ g/mol}}}\text{ * }\frac{{1000}\text{ mL/L}}{\text{mL NaOH}}$

### Determination of the Acid Content of the Soda

In this experiment you will determine the concentration of citric acid in a soda. In order to do this we will titrate the soda with a standardized solution of sodium hydroxide. Most soda contains acid to provide a tangy taste. The soda used in this lab has citric acid (a triprotic acid) as a main ingredient. You can determine the concentration of H+ and citric acid in your soda sample using stoichiometry. Note that it is important that your sample of soda be flat. The presence of carbon dioxide gas in the solution will greatly affect your numbers. Soda should be heated to remove gas prior to its use.

### Confirmation that Citric Acid is Triprotic

Citric acid (H3C6H5O7) has three hydrogen atoms that can be dissociated in solution. You can titrate a small sample of citric acid with your standardized NaOH to determine if the acid really is triprotic (1 mol citric acid : 3 mol NaOH).