# 21.3: Buffers

Difficulty Level: At Grade Created by: CK-12

## Student Behavioral Objectives

The student will:

• define and give an example of a buffer.
• explain the effect of a strong acid or base on a buffer system.
• explain the mechanism by which a buffer solution resists changes in pH.
• given appropriate information, calculate the pH of a buffer.
• describe how to make a buffer solution.

## Timing, Standards, Activities

Timing and California Standards
Lesson Number of 60 min periods CA Standards
Buffers 2.0 5g

### Activities for Lesson 3

Laboratory Activities

1. None

Demonstrations

1. None

Worksheets

1. Buffers Worksheet

1. None

## Answers for Buffers (L3) Review Questions

• Sample answers to these questions are available upon request. Please send an email to teachers-requests@ck12.org to request sample answers.

## Multimedia Resources for Chapter 21

This website provides a lesson on buffer solutions.

This website provides a list of household acid-base indicators.

This website provides a pH calculation problem generator.

This website provides an animation of the autoionization of water.

The following link shows a video of a neutralization reaction.

This video shows the technique for performing a titration using an indicator.

The following link is to a video about acid-base neutralization and titration.

The video at the link below shows the lab techniques needed for titration.

This video is a ChemStudy film called “Acid Base Indicators.” The film is somewhat dated but the information is accurate.

To see a short animated video showing concentration changes as strong acid or base is added to a buffer, follow the link below.

## Teacher's Pages for Hydrolysis of Salts

Lab Notes

Preparation of Solutions:

\begin{align*}200 \ mL\end{align*} of each solution should be more than enough to complete this lab. To prepare each solution, mass the specified amount of reagent, dissolve it in \begin{align*}150 \ mL\end{align*} of water, and dilute the resulting solution to \begin{align*}200 \ mL\end{align*}:

\begin{align*}CuSO_4 \cdot 5H_2O \ : \ 5.0 \ grams\end{align*}

\begin{align*}Ca(NO_3)_2 \ : \ 3.3 \ grams\end{align*}

\begin{align*}K_3PO_4 \ : \ 4.2 \ grams\end{align*}

\begin{align*}KCl \ : \ 1.5 \ grams\end{align*}

\begin{align*}NaBr \ : \ 2.1 \ grams\end{align*}

\begin{align*}Na_2S \ : \ 1.6 \ grams\end{align*}

\begin{align*}(NH_4)_2CO_3 \ : \ 1.9 \ grams\end{align*}

\begin{align*}Na_2CrO_4 \ : \ 3.2 \ grams\end{align*}

\begin{align*}MgBr_2 \ : \ 3.7 \ grams\end{align*}

\begin{align*}NaCl \ : \ 1.2 \ grams\end{align*}

Use care when opening the container of \begin{align*}(NH_4)_2CO_3\end{align*}. It undergoes decomposition over time, and outgases \begin{align*}NH_3\end{align*}, which collects in the container. Upon opening the \begin{align*}NH_3\end{align*} diffuses out rapidly, and is very irritating to eyes and skin. Live and learn.

1. Write dissociation equations for the following salts:
1. \begin{align*}CuSO_4 \rightarrow Cu^{2+} + SO_4^{2+}\end{align*}
2. \begin{align*}Ca(NO_3)_2 \rightarrow Ca^{2+} + 2 \ NO_3^-\end{align*}
3. \begin{align*}K_3PO_4 \rightarrow 3 \ K^+ + PO_4^{3-}\end{align*}
4. \begin{align*}KCl \rightarrow K^+ + Cl^-\end{align*}
5. \begin{align*}KBr \rightarrow K^+ + Br^-\end{align*}
6. \begin{align*}Na_2S \rightarrow 2 \ Na^+ + S^{2-}\end{align*}
7. \begin{align*}(NH_4)_2CO_3 \rightarrow 2 \ NH_4^+ + CO_3^{2-}\end{align*}
8. \begin{align*}Na_2CrO_4 \rightarrow 2 \ Na^+ + CrO_4^{2-}\end{align*}
9. \begin{align*}MgBr_2 \rightarrow Mg^{2+} + 2 \ Br^-\end{align*}
10. \begin{align*}NaCl \rightarrow Na^+ + Cl^-\end{align*}
2. \begin{align*}NaC_6H_5CO_2 + H_2O \rightarrow HC_6H_5CO_2 + Na^+ + OH^-\end{align*} solution will be basic

## Lab – Hydrolysis of Salts

Background Information

A salt is an ionic compound containing positive ions other than hydrogen and negative ions other than hydroxide. Most salts will dissociate to some degree when placed in water. In many cases, ions from the salt will react with water molecules to produce hydrogen ions, \begin{align*}H^+\end{align*}, or hydroxide ions, \begin{align*}OH^-\end{align*}. Any chemical reaction in which water is one of the reactants is called a hydrolysis reaction. Salts are usually formed from the neutralization reaction between an acid and a base. A salt formed from a strong acid and a strong base will not undergo hydrolysis. The resulting solution is neutral. An example of such a salt is \begin{align*}KBr\end{align*}, formed from a strong acid, \begin{align*}HBr\end{align*}, and a strong base, \begin{align*}KOH\end{align*}.

Salts formed from the reaction of a strong acid and a weak base hydrolyzes to form a solution that is slightly acidic. In this kind of hydrolysis, the water molecules actually react with the cation from the weak base. For example, when ammonium chloride, \begin{align*}NH_4Cl\end{align*}, hydrolyzes, water molecules react with the \begin{align*}NH_4^+\end{align*} ion:

\begin{align*}NH_4^+ + H_2O \rightarrow NH_4OH + H^+\end{align*}

The formation of the \begin{align*}H^+\end{align*} ion from this reaction makes the solution acidic.

Salts formed from the reaction of a weak acid and a strong base hydrolyze to form a solution that is slightly basic. In this kind of hydrolysis, it is the anion from the weak acid that actually reacts with the water. For example, when sodium acetate, \begin{align*}NaC_2H_3O_2\end{align*}, hydrolyzes, water molecules react with the acetate ion:

\begin{align*}C_2H_3O_2^- + H_2O \ \rightarrow \ HC_2H_3O_2 + OH^-\end{align*}

The formation of the \begin{align*}OH^-\end{align*} ion from this reaction makes the solution basic. Salts formed from a weak acid and weak base produce solutions that may be slightly acidic, slightly basic, or neutral, depending on how strongly the ions of the salt are hydrolyzed.

In this experiment you will test several different salt solutions with \begin{align*}pH\end{align*} paper and phenolphthalein solution to determine their acidity or basicity.

Purpose

To determine the relative acidity or basicity of various salt solutions, and thus predict whether hydrolysis occurred, and if so, what the reaction products are.

Pre-Lab Questions

1. Write dissociation equations for the following salts:
1. Copper(II) sulfate
2. Calcium nitrate
3. Potassium phosphate
4. Potassium chloride
5. Potassium bromide
6. Sodium sulfide
7. Ammonium carbonate
8. Sodium chromate
9. Magnesium bromide
10. Sodium chloride
2. Sodium benzoate is the salt formed in the neutralization of benzoic acid with sodium hydroxide. Benzoic acid is a weak acid. Write the hydrolysis reaction for the dissolution of solid sodium benzoate, \begin{align*}NaC_6H_5CO_2\end{align*}, in water. Will sodium benzoate solution be acidic, basic, or neutral?

Apparatus and Materials

• 10 small or medium sized test tubes, or a micro reaction plate
• Test tube rack
• \begin{align*}0.1 \ M\end{align*} solutions of cupric sulfate, calcium nitrate, potassium phosphate, potassium chloride, sodium bromide, sodium sulfide, ammonium carbonate, sodium chromate, magnesium bromide and sodium chloride
• Universal \begin{align*}pH\end{align*} indicator paper, range 0-14
• Phenolphthalein indicator solution
• \begin{align*}10 \ mL\end{align*} graduate
• Stirring rod

Safety Issues

The solutions used may be slightly acidic or basic, and as a result can be corrosive or caustic. Use proper laboratory safety equipment and techniques.

Procedure

1. Obtain a clean, dry micro reaction plate, or 10 test tubes

2. To test tubes 1 through 10 or the reaction plate, add eight to ten drops of the following solutions:

• Tube or Well 1: Cupric sulfate
• Tube or Well 2: Calcium nitrate
• Tube or Well 3: Potassium phosphate
• Tube or Well 4: Potassium chloride
• Tube or Well 5: Sodium bromide
• Tube or Well 6: Sodium sulfide
• Tube or Well 7: Ammonium carbonate
• Tube or Well 8: Sodium chromate
• Tube or Well 9: Magnesium bromide
• Tube or Well 10: Sodium chloride

3. Add two drops of phenolphthalein solution to each of the occupied wells of the microplate or test tube. Record your observations in the data table.

4. Test each solution with \begin{align*}pH\end{align*} paper and record your results.

Data

Data Table
Well # Salt Effect on Indicator pH Original Acid Strong or Weak Acid Original Base Strong or Weak Base
1
2
3
4
5
6
7
8
9
10

Post-Lab Questions

1. Where a hydrolysis is likely to occur in each of the following, write a net ionic hydrolysis equation. If no hydrolysis is likely, write NR.

\begin{align*}Cu^{2+}+SO_4^{2-}+2 \ H_2O & \rightarrow\\ Ca^{2+}+2 \ NO_3^-+2 \ H_2O & \rightarrow \\ 3 \ K^++PO_4^{3-}+3 \ H_2O & \rightarrow \\ K^++Cl^-+H_2O & \rightarrow \\ Na^++Br^-+H_2O & \rightarrow \\ 2 \ Na^++S^{2-}+H_2O & \rightarrow \\ 2 \ NH_4^++CO_3^{2-}+H_2O & \rightarrow \\ 2 \ Na^++CrO_4^{2-}+2 \ H_2O & \rightarrow\\ Mg^{2+}+2 \ Br^-+2 \ H_2O & \rightarrow \\ Na^++Cl^-+H_2O & \rightarrow\end{align*}

2. How do your observations and \begin{align*}pH\end{align*} readings compare with the expected results based on the equations for the hydrolysis reactions?

3. What is a spectator ion? Name the spectator ions present in each hydrolysis reaction in this experiment.

4. A salt formed from a strong acid and a strong base produces a neutral solution. A salt of a weak acid and a weak base may or may not produce a neutral solution. Explain why.

5. Bases make effective cleaning agents, because they can convert grease and oils to a water soluble substance. Trisodium phosphate (TSP) is a common commercially available cleaner. Give the reaction TSP undergoes to create a basic solution.

## Teacher's Pages for pH Measurements Using Indicators

Lab Notes

Buffered solutions of various \begin{align*}pH\end{align*} values can be purchased in dropper bottles, as can dropper bottles of indicator solutions. If you have many chemistry classes and perform the experiment for several years, it may be more economical to prepare the solutions yourself.

The solutions used in this lab can be prepared as follows.

• \begin{align*}pH = 1\end{align*} solution: dilute \begin{align*}8.3 \ mL\end{align*} of concentrated \begin{align*}HCl\end{align*} \begin{align*}(12 \ M)\end{align*} to \begin{align*}1.00 \ liter\end{align*}
• \begin{align*}pH = 3\end{align*} solution: dilute \begin{align*}10. \ mL\end{align*} of \begin{align*}pH = 1\end{align*} solution (above) to \begin{align*}1.00 \ liter\end{align*}
• \begin{align*}pH = 5\end{align*} solution: dilute \begin{align*}10. \ mL\end{align*} of \begin{align*}pH = 3\end{align*} solution (above) to \begin{align*}1.00 \ liter\end{align*}
• \begin{align*}pH = 7\end{align*} solution: distilled water
• \begin{align*}pH = 13\end{align*} solution: dissolve \begin{align*}4.00 \ g\end{align*} of \begin{align*}NaOH\end{align*} in sufficient water to produce \begin{align*}1.00 \ liter\end{align*} of solution
• \begin{align*}pH = 11\end{align*} solution: dilute \begin{align*}10. \ mL\end{align*} of \begin{align*}pH = 13\end{align*} solution (above) to \begin{align*}1.00 \ liter\end{align*}
• \begin{align*}pH = 9\end{align*} solution: dilute \begin{align*}10. \ mL\end{align*} of \begin{align*}pH = 11\end{align*} solution (above) to \begin{align*}1.00 \ liter\end{align*}
• methyl orange indicator: dissolve \begin{align*}0.1 \ g\end{align*} of methyl orange powder in \begin{align*}100 \ mL\end{align*} of water and filter
• bromthymol blue indicator: dissolve \begin{align*}.01 \ g\end{align*} of bromthymol blue in \begin{align*}100 \ mL\end{align*} of 50% water and 50% ethanol solution and filter
• phenolphthalein solution: dissolve \begin{align*}1.0 \ g\end{align*} of phenolphthalein powder in \begin{align*}100 \ mL\end{align*} of ethanol

To prepare the unknown solutions for Part IV, select three of the known \begin{align*}pH\end{align*} solutions used in the lab and label them as unknown. Be sure to keep a record of which \begin{align*}pH\end{align*} values were selected as unknowns.

1. How is universal indicator made?

Several indicators are mixed.

2. What distinguishes weak organic acids that are useful as acid-base indicators from weak organic acids that will not function as acid-base indicators?

The undissociated molecule of the acid and the anion of the dissociated acid must be different colors.

## pH Measurements Using Indicators

Background Information

The nature of acids and bases have been known to man for quite sometime. Chemically speaking, acids are interesting compounds because a large number of common household substances are acids or acidic solutions. For example, vinegar contains ethanoic acid, also called acetic acid, \begin{align*}HC_2H_3O_2\end{align*}) and citrus fruit contain citric acid. Acids cause foods to have a sour taste and turn litmus red. (Note: You should never taste substances in the laboratory.) Also, many common household substances are bases. Milk of magnesia contains the base magnesium hydroxide, \begin{align*}Mg(OH)_2\end{align*} and household ammonia is a common cleaning agent. Bases have a slick feel to the fingers and turn litmus blue. (Note: You should never feel chemicals in the laboratory.)

Indicator dyes, of which litmus is one, turn various colors according to the strength of the acid or base applied to it.

Pure water, which is neutral in terms of acid-base, exists mostly as \begin{align*}H_2O\end{align*} molecules but does, to a very slight extent dissociate into hydrogen and hydroxide ions.

\begin{align*}HOH_{(L)} \rightarrow H^+_{(aq)} + OH^-_{(aq)}\end{align*}

The extent of this dissociation is \begin{align*}1.0 \times 10^{-7} \ moles/liter\end{align*} (at \begin{align*}25^\circ C\end{align*}). Therefore, in all neutral water (and neutral water solutions), the concentration of hydrogen ions is \begin{align*}1.0 \times 10^{-7} \ M\end{align*} and the concentration of hydroxide ions is \begin{align*}1.0 \times 10^{-7} \ M\end{align*}. The dissociation constant for this process is \begin{align*}K_w = [H^+][OH^-] = (1.0 \times 10^{-7})(1.0 \times 10^{-7} = 1.0 \times 10^{-14})\end{align*}.

In 1909, a Danish chemist (Soren Sorenson), developed a mathematical system for referring to the degree of acidity of a solution. He used the term \begin{align*}pH\end{align*} for “power of hydrogen” and established the equation, \begin{align*}pH = -log \ [H^+]\end{align*}.

In a neutral solution, the hydrogen ion concentration is \begin{align*}1.0 \times 10^{-7} \ M\end{align*} and therefore, the \begin{align*}pH\end{align*} is 7. If the concentration of hydrogen ions is \begin{align*}1.0 \times 10^{-5} \ M\end{align*}, then the \begin{align*}pH\end{align*} is 5. A solution is neutral when the \begin{align*}pH\end{align*} equals 7, it is acid if the \begin{align*}pH\end{align*} is less than 7, and it is basic if the \begin{align*}pH\end{align*} is more than 7. In commonly used solutions, \begin{align*}pH\end{align*} values usually range from 1 to 14.

Living matter (protoplasm) contains a mixture of variously dissociated acids, bases, and salts and usually has a \begin{align*}pH\end{align*} very near neutral. The \begin{align*}pH\end{align*} of human blood is generally 7.3 and humans cannot survive if the blood becomes more basic than \begin{align*}pH \ 7.8\end{align*} or more acidic than \begin{align*}pH \ 7.0\end{align*}. Life of any kind exists only between \begin{align*}pH \ 3\end{align*} and \begin{align*}pH \ 8.5\end{align*}. Buffer solutions regulate the \begin{align*}pH\end{align*} of the body by neutralizing excess acid or base. The chief buffers of the body are proteins, carbonates, phophates, and hemoglobin. The kidneys play a role by eliminating excess electrolytes.

Some Typical
pH Substance Acidity/Basicity
0 Sulfuric Acid (Battery Acid) Very Highly Acidic
1 \begin{align*}0.10 \ M\end{align*} Hydrochloric Acid Highly Acidic
2 Stomach Acid Acidic
3 Vinegar Acidic (\begin{align*}\frac{1}{100}^{th}\end{align*} as strong as \begin{align*}pH \ 1\end{align*})
4 Tomato Juice
5 Black Coffee and Vitamin C Weakly Acidic
6 Cow's Milk Very Weakly Acidic
7 Distilled Water Neutral
8 Sea Water Very Weakly Basic
9 Baking Soda, \begin{align*}NaHCO_3\end{align*} Weakly Acidic
10 Detergents
11 Basic
12 Household Cleaning Ammonia
13 \begin{align*}0.10 \ M\end{align*} \begin{align*}NaOH\end{align*} Strongly Basic
14 \begin{align*}1.0 \ M\end{align*} \begin{align*}NaOH\end{align*} (Lye) Very Strongly Basic

Some acids dissociate completely into ions when dissolved in water. Such acids are called strong acids (\begin{align*}HCl\end{align*}, \begin{align*}HI\end{align*}, \begin{align*}HBr\end{align*}, \begin{align*}HNO_3\end{align*}, \begin{align*}H_2SO_4\end{align*}, \begin{align*}HClO_4\end{align*}). Some bases dissociate completely when dissolved in water. Such bases are called strong bases (\begin{align*}NaOH\end{align*}, \begin{align*}LiOH\end{align*}, \begin{align*}KOH\end{align*}, \begin{align*}RbOH\end{align*}). There are other acids and bases that dissociate only slightly (although completely soluble) when dissolved in water. Such acids and bases are called weak acids or weak bases and some examples are \begin{align*}HF\end{align*}, \begin{align*}HC_2H_3O_2\end{align*}, \begin{align*}NH_4OH\end{align*}.

An important method of determining \begin{align*}pH\end{align*} values in the lab involves the use of substances called “acid-base indicators”. These are certain organic substances (almost always weak organic acids) that have the property of changing color in solutions of varying hydrogen ion concentration. In order for a weak organic acid to be useful as an acid-base indicator, it is necessary that the undissociated molecule and the indicator anion be different colors. For example, phenolphthalein is a colorless substance in any aqueous solution in which the hydrogen ion concentration is greater than \begin{align*}1 \times 10^{-9} \ M \ (pH < 9)\end{align*} but changes to a red or pink color when the hydrogen ion concentration is less than \begin{align*}1 \times 10^{-9} \ M \ (pH > 9)\end{align*}. Such substances can be used for determining the approximate \begin{align*}pH\end{align*} of solutions. Electrical measurements can determine the \begin{align*}pH\end{align*} even more precisely. This lab will use three acid base indicators and what is called a “universal indicator”.

Some Indicator Color Changes
Indicator pH Color Change Range Color Change
Methyl Orange \begin{align*}3.1 - 3.4\end{align*} Red to Yellow
Bromthymol Blue \begin{align*}6.0 - 7.6\end{align*} Yellow to Blue
Phenolphthalein \begin{align*}8.3 - 10.0\end{align*} Colorless to Red

The universal indicator (one type is called Bogen’s Universal Indicator), is made by mixing a number of indicators that all change color at different \begin{align*}pH’s\end{align*}. As you slowly change the \begin{align*}pH\end{align*} of the indicator from 1 to 14, it goes through a series of subtle color changes. The indicator is provided with a photographic chart that shows the color of the indicator at every different \begin{align*}pH\end{align*} and the \begin{align*}pH\end{align*} is identified by matching the indicator color to the chart.

Pre-Lab Questions

1. How is universal indicator made?
2. What distinguishes weak organic acids that are useful as acid-base indicators from weak organic acids that will not function as acid-base indicators?

Purpose

The purpose of this lab is to have the student experience the color changes involved with acid-base indicators and to identify the approximate \begin{align*}pH\end{align*} of an unknown solution using acid-base indicators.

Apparatus and Materials

• Well Plates, at least 12 wells (1 per lab group)
• drop controlled bottles of \begin{align*}pH = 1\end{align*}
• drop controlled bottles of \begin{align*}pH = 3\end{align*}
• drop controlled bottles of \begin{align*}pH = 5\end{align*}
• drop controlled bottles of \begin{align*}pH = 7\end{align*}
• drop controlled bottles of \begin{align*}pH = 9\end{align*}
• drop controlled bottles of \begin{align*}pH = 11\end{align*}
• drop controlled bottles of \begin{align*}pH = 13\end{align*}
• drop controlled bottles of methyl orange
• drop controlled bottles of bromthymol blue
• drop controlled bottles of phenolphthalein
• drop controlled bottles of universal indicator
• drop controlled bottle of unknown #1
• drop controlled bottle of unknown #2
• drop controlled bottle of unknown #3

Safety Issues

All solutions are irritating to skin, eyes, and mucous membranes. Handle solutions with care, avoid getting the material on you, and wash your hands carefully before leaving the lab.

Procedure for Part I: Determining the effect of \begin{align*}pH\end{align*} on indicator dyes.

1. Place the Chemplate on a sheet of white paper.
2. Place one drop of methyl orange into cavities #1 and #2.
3. Place one drop of bromthymol blue in cavities #5 and #6.
4. Place one drop of phenolphthalein in cavities #9 and #10.
5. Carefully add one drop of \begin{align*}pH \ 1\end{align*} to cavities #1, #5, and #9.
6. Carefully add one drop of \begin{align*}pH \ 13\end{align*} to cavities #2, #6, and #10.

Data for Part I

1. What color is the original methyl orange solution? ______________
2. What color is methyl orange in a strong acid?______________
3. What color is methyl orange in a strong base? ______________
4. What color is the original bromthymol blue solution? ______________
5. What color is bromthymol blue in a strong acid? ______________
6. What color is bromthymol blue in a strong base? ______________
7. What color is the original phenolphthalein solution? ______________
8. What color is phenolphthalein in a strong acid? ______________
9. What color is phenolphthalein in a strong base? ______________

Rinse the Chemplate in tap water and dry with a paper towel.

Procedure for Part II: Determining the \begin{align*}pH\end{align*} color change range of indicator dyes.

1. Place one drop of methyl orange in each cavity numbered 1 – 7.
2. Carefully add one drop of \begin{align*}pH \ 1\end{align*} to cavity #1, \begin{align*}pH \ 3\end{align*} to cavity #2, \begin{align*}pH \ 5\end{align*} to cavity #3, \begin{align*}pH \ 7\end{align*} to #4, \begin{align*}pH \ 9\end{align*} to #5, \begin{align*}pH \ 11\end{align*} to #6, and \begin{align*}pH \ 13\end{align*} to #7.
3. Repeat the rinsing, drying, and steps 1 and 2 except using bromthymol blue and then repeat the entire process again using phenolophthalein.

Data for Part II

1. Describe the color changes and the \begin{align*}pH’s\end{align*} around the color change \begin{align*}pH\end{align*} for methyl orange.
2. Describe the color changes and the \begin{align*}pH’s\end{align*} around the color change \begin{align*}pH\end{align*} for bromthymol blue.
3. Describe the color changes and the \begin{align*}pH’s\end{align*} around the color change \begin{align*}pH\end{align*} for phenolphthalein.

Rinse the Chemplate in tap water and dry with a paper towel.

Procedure for Part III: Determining a color standard for universal indicator.

1. Place one drop of universal indicator in each cavity numbered 1 – 7.
2. Carefully add one drop of \begin{align*}pH \ 1\end{align*} to cavity #1, \begin{align*}pH \ 3\end{align*} to cavity #2, \begin{align*}pH \ 5\end{align*} to cavity #3, \begin{align*}pH \ 7\end{align*} to #4, \begin{align*}pH \ 9\end{align*} to #5, \begin{align*}pH \ 11\end{align*} to #6, and \begin{align*}pH \ 13\end{align*} to #7.

Keep these solutions for Part IV.

Data for Part III

1. Describe the color of the universal indicator at each \begin{align*}pH\end{align*} used.

Cavity #1 \begin{align*}(pH = 1)\end{align*}, color = __________________

Cavity #2 \begin{align*}(pH = 3)\end{align*}, color = __________________

Cavity #3 \begin{align*}(pH = 5)\end{align*}, color = ___________________

Cavity #4 \begin{align*}(pH = 7)\end{align*}, color = ___________________

Cavity #5 \begin{align*}(pH = 9)\end{align*}, color = ___________________

Cavity #6 \begin{align*}(pH = 11)\end{align*}, color = ___________________

Cavity #7 \begin{align*}(pH = 13)\end{align*}, color = ___________________

Procedure for Part IV: Determining the \begin{align*}pH\end{align*} of some unknown solutions.

1. Place one drop of universal indicator in cavities #10, #11, and #12.
2. Place one drop of unknown #1 in cavity #10.
3. Place one drop of unknown #2 in cavity #11.
4. Place one drop of unknown #2 in cavity #12.
5. Compare the color in each cavity with the colors in cavities # 1 – 7 that you made in activity 3.

Data for Part IV

Color of unknown #1 in universal indicator ___________________

Color of unknown #2 in universal indicator ___________________

Color of unknown #3 in universal indicator ___________________

Post-Lab Questions

1. What is the \begin{align*}pH\end{align*} of unknown #1? _____________
2. What is the \begin{align*}pH\end{align*} of unknown #2? _____________
3. What is the \begin{align*}pH\end{align*} of unknown #3? _____________

## Teacher's Pages for Acid-Base Titration

Lab Notes

Students will need two days to do both parts of the lab.

Preparation of Solutions, \begin{align*}KHP\end{align*}, and unknown acids

\begin{align*}6 \ M \ NaOH\end{align*} solution: Boil \begin{align*}600 \ mL\end{align*} of distilled water to drive off any dissolved \begin{align*}CO_2\end{align*}. (The \begin{align*}CO_2\end{align*} produces carbonic acid, which drives down the concentration of \begin{align*}NaOH\end{align*}.) Add \begin{align*}120. \ g\end{align*} of \begin{align*}NaOH\end{align*} to a \begin{align*}500 \ mL\end{align*} volumetric flask, and add some of the freshly boiled distilled \begin{align*}H2O\end{align*} to the flask. Swirl to dissolve. Cool the resultant solution in a cold water or ice water bath, let the solution and flask return to room temperature, and dilute the resulting solution to \begin{align*}500 \ mL\end{align*}. Store this solution in a tightly capped bottle, preferably Nalgene or other base-resistant bottle. This will provide enough solution for (75) two-student teams with a 20% excess for spills and endpoint over-runs.

Phenolphthalein indicator solution: Dissolve \begin{align*}0.1 \ g\end{align*} of phenolphthalein in \begin{align*}50 \ mL\end{align*} of 95% ethanol, and dilute to \begin{align*}100 \ mL\end{align*} by adding distilled water. This will provide enough solution for (75) two student teams.

\begin{align*}KHP\end{align*} – Potassium Hydrogen Phthalate – Dry \begin{align*}200 \ g\end{align*} of \begin{align*}KHP\end{align*} in a laboratory oven for at least one hour prior to titration. Store the \begin{align*}KHP\end{align*} in a dessicator. \begin{align*}KHP\end{align*} is slightly hygroscopic. This is sufficient \begin{align*}KHP\end{align*} for (75) two – student teams.

Unknown Acids

The following acids are suggestions for use. Their number of ionizable hydrogens vary, are stable chemically, and many schools have them on hand. Twenty grams of each acid are required for (75) two – student teams:

Acid Number of Ionizable Hydrogens Molar Mass (g/mol)
Lactic 1 \begin{align*}90.1\end{align*}
Malonic 2 \begin{align*}104.1\end{align*}
Maleic 2 \begin{align*}116.1\end{align*}
Succinic 2 \begin{align*}118.1\end{align*}
Benzoic 1 \begin{align*}122.1\end{align*}
Salicylic 1 \begin{align*}138.1\end{align*}
Tartaric 2 \begin{align*}150.1\end{align*}

NOTE: The molecular weights listed are for the anhydrous acids. If hydrates are used, use the molecular weights as shown on the reagent bottle. The number of ionizable hydrogens does not change.

1. Pre-rinsing the buret will remove any water or residual \begin{align*}NaOH\end{align*} solution within the buret. If there were water present, it would dilute the \begin{align*}NaOH\end{align*} added. If there were residual \begin{align*}NaOH\end{align*}, the concentration would increase upon addition of solution due to the crystalline or concentrated solution of \begin{align*}NaOH\end{align*} present.

2. The \begin{align*}KHP\end{align*} and the \begin{align*}NaOH\end{align*} react with each other in a 1:1 stoichiometric ratio. Thus, the number of moles of \begin{align*}NaOH\end{align*} required to react with the \begin{align*}KHP\end{align*} will be equal to the number of moles of \begin{align*}KHP\end{align*} originally present.

\begin{align*}Moles \ of \ NaOH \ \text{added} & = M \times L = (0.1 \ M)(0.040 \ L) = 0.004 \ moles \ NaOH \ \text{added: thus}\\ Moles \ KHP \ \text{needed}& = 0.004 \ moles\\ Grams \ of \ KHP \ \text{needed} & = (moles)(g/mole) = (0.004 \ mol)(204.23 \ g/mol) = 0.8 \ grams \ KHP\end{align*}

3. \begin{align*}Molarity \ of \ KHP = \frac{moles \ KHP}{liters \ of \ solution}\end{align*}

\begin{align*}Moles \ of \ KHP = \frac{grams \ KHP}{g/mol \ of \ KHP} = \frac{0.759 \ g}{204.2 \ g/mol} = 0.00372 \ mol \ KHP\end{align*}

\begin{align*}Molarity = \frac {0.00372 \ mol \ KHP} {0.0500 \ L} = 0.0743 \ M \ KHP\end{align*}

4. Again, the amount of water needed to dissolve the \begin{align*}KHP\end{align*} is not needed to solve this problem. All you need is the weight of the \begin{align*}KHP\end{align*} and it’s molecular weight :

\begin{align*}Moles \ of \ KHP = \frac{0.521 \ g} {204.2 \ g/mol} = 0.00255 \ mols \ KHP\end{align*}

Since this will be equal to the number of moles of \begin{align*}NaOH\end{align*} reacted, the number of liters of titrant you need to add can be calculated by re-arranging the molarity formula:

\begin{align*}M & = \frac {moles} {L}\\ \text{so} \ L & = \frac{moles}{M} = \frac {0.00255 \ mol}{0.102 \ mol/L} = 0.0213 \ L = 21.3 \ mL\end{align*}

## Acid-Base Titration Lab

Background Information

Standardization of the sodium hydroxide solution through titration is necessary because it is not possible to directly prepare a known molarity solution of sodium hydroxide with high accuracy. Solid sodium hydroxide readily absorbs moisture and carbon dioxide from the atmosphere and thus it is difficult to obtain a precise amount of the pure substance. A sodium hydroxide solution will be made close to \begin{align*}0.1 \ M\end{align*} and then the actual molarity of the solution will be determined by titration of a primary standard. A primary standard is a substance of very high purity that is also stable in air. Because the substance remains pure, it is possible to mass a sample of the substance with a high degree of accuracy. The primary standard used in this experimental procedure is potassium hydrogen phthalate, \begin{align*}KHC_8H_4O_4\end{align*} (molar mass \begin{align*}204.2 \ g/mole\end{align*}).

This primary standard, \begin{align*}KHP\end{align*}, is used to standardize the secondary standard, sodium hydroxide. The standardized sodium hydroxide solution can then be used to determine the molar mass of an unknown acid through titration.

In both steps a titration is performed in which a buret is used to dispense measured increments of the sodium hydroxide solution into a second solution containing a known mass of \begin{align*}KHP\end{align*} (\begin{align*}NaOH\end{align*} standardization). For the second reaction, a mass of acid whose molecular weight is unknown is then titrated with the solution of \begin{align*}NaOH\end{align*} whose concentration was determined by the standardization with \begin{align*}KHP\end{align*}. The stoichiometry of the reaction depends on the number of ionizable hydrogens within the acid. \begin{align*}KHP\end{align*} is a weak monoprotic acid that will react with sodium hydroxide in a 1:1 mole ratio:

\begin{align*}KHC_8H_4O_4 + NaOH \rightarrow KNaC_8H_4O_4 + H_2O\end{align*}

The unknown acid may be monoprotic, diprotic, or triprotic dependent on the number of acidic hydrogens present in the molecule. A monoprotic acid, \begin{align*}HA\end{align*}, has one acidic hydrogen, a diprotic acid, \begin{align*}H_2A\end{align*}, two acidic hydrogens, and a triprotic acid, \begin{align*}H_3A\end{align*}, three acidic hydrogens. The stoichiometries of the reactions are shown below. You will be told whether your unknown acid is monoprotic, diprotic, or triprotic.

\begin{align*}\text{Monoprotic} \ HA + NaOH & \rightarrow NaA + H_2O\\ \text{Diprotic} \ H_2A + 2 NaOH & \rightarrow Na_2A + 2 H_2O\\ \text{Triprotic} \ H_3A + 3 NaOH & \rightarrow Na_3A + 3 H_2O\end{align*}

The indicator phenolphthalein is used as a signal of the equivalence point. Phenolphthalein is a weak organic acid that will change from colorless to pink near the equivalence point of the titrations. The actual point at which the indicator changes color is the end point. The endpoint and the equivalence point are not the same. The difference between the two is the titration error. Obviously, for a titration to be of value, care must be taken to select an indicator for which the difference between the equivalence point and the endpoint is small. With this particular titration, it is very small.

Pre-Lab Questions

1. Why is it necessary to rinse out the buret with the \begin{align*}NaOH\end{align*} solution?
2. Calculate the approximate weight of \begin{align*}KHP\end{align*} required so that about \begin{align*}40 \ mL\end{align*} of \begin{align*}0.1 \ M\end{align*} sodium hydroxide is used in a titration. (MW of \begin{align*}KHP = 204.23 \ g/mol\end{align*})
3. Calculate the molarity of a \begin{align*}KHP\end{align*} solution when \begin{align*}0.759 \ g\end{align*} of \begin{align*}KHP\end{align*} is dissolved in \begin{align*}50.0 \ mL\end{align*} of water.
4. \begin{align*}0.521 \ g\end{align*} of \begin{align*}KHP\end{align*} is dissolved in \begin{align*}40 \ mL\end{align*} of water, and titrated with a \begin{align*}0.102 \ M\end{align*} \begin{align*}NaOH\end{align*} solution. Calculate the number of \begin{align*}mL\end{align*} of the \begin{align*}NaOH\end{align*} solution added.

Purpose

The purpose of this experiment is to determine the concentration of a titrating solution, \begin{align*}NaOH\end{align*}, using a stable compound, \begin{align*}KHP\end{align*}. Once the concentration of the \begin{align*}NaOH\end{align*} solution is known, it can then be used to determine the molecular weight of an acid whose formula is unknown.

Apparatus and Materials

• \begin{align*}50 \ mL\end{align*} buret
• buret stand
• buret clamp
• \begin{align*}125 \ mL\end{align*} Erlenmeyer flask
• phenolphthalein
• \begin{align*}NaOH\end{align*} solution
• \begin{align*}KHP\end{align*}, solid
• unknown acid, solid
• \begin{align*}10 \ mL\end{align*} graduate
• \begin{align*}400 \ mL\end{align*} beaker
• \begin{align*}100 \ mL\end{align*} graduate

Safety Issues

\begin{align*}NaOH\end{align*} is a caustic solution and will cause severe burns, especially to eye tissue. Wear goggles and aprons at all times. The solid acids cause considerable irritation if exposed to skin or mucous membranes. Avoid exposure.

Procedure for Part I

Part 1. Standardization of Sodium Hydroxide Solution

Obtain the primary acid standard from your instructor. Record the name of the acid, its molecular formula, and number of acidic hydrogens per molecule. Prepare about \begin{align*}300 \ mL\end{align*} of approximately \begin{align*}0.1 \ M\end{align*} sodium hydroxide by diluting \begin{align*}6 \ M\end{align*} \begin{align*}NaOH\end{align*} with distilled water. (Calculate, ahead of time, how much water and how much \begin{align*}6 \ M\end{align*} \begin{align*}NaOH\end{align*} will be needed.)

WARNING: Concentrated sodium hydroxide is corrosive and causes severe burns. Handle with care. Dilute and wash up spills with plenty of water. Wash affected skin with water until it no longer feels slippery, but feels “squeaky” clean.

Store the solution in your plastic bottle, label it “\begin{align*}0.1 \ M\end{align*} \begin{align*}NaOH\end{align*}”, and keep it tightly capped. You will determine the exact molarity of this \begin{align*}NaOH\end{align*} solution by standardization.

Calculate the mass of the primary acid standard that would react with about \begin{align*}20 \ mL\end{align*} of \begin{align*}0.1 \ M\end{align*} \begin{align*}NaOH\end{align*}. Weigh approximately this amount into a clean \begin{align*}125-mL\end{align*} Erlenmeyer flask by taring the balance with the flask on the pan, and then adding the acid to the flask. Record the mass of the primary acid standard to the highest precision allowed by the electronic balance. Add \begin{align*}30\end{align*} to \begin{align*}40 \ mL\end{align*} of your purified water to the flask, and swirl to dissolve the primary acid standard. Add three or four drops of phenolphthalein indicator solution to the flask, and swirl to mix well. Label this flask and keep it tightly capped until ready for use.

Rinse the inside of a CLEAN buret three times with small quantities of your \begin{align*}0.1 \ M\end{align*} sodium hydroxide solution (called “rinsing in” with the solution to be used in the buret). Drain the rinses though the stopcock and tip. Do not forget to rinse liquid through the tip, to replace water there. Fill the buret above the \begin{align*}0.0 \ mL\end{align*} mark with \begin{align*}0.1 \ M\end{align*} \begin{align*}NaOH\end{align*}, and then drain it until the meniscus is slightly below

NOTE: Do not waste the time it takes to set the starting level to exactly \begin{align*}0.0 \ mL\end{align*}. It is more efficient and more accurate to set the level between \begin{align*}1\end{align*} and \begin{align*}2 \ mL\end{align*} and read the starting level precisely.

Tips on Technique

• To read the buret accurately, hold a white card with a black stripe behind the buret, with the black stripe below the meniscus, and the meniscus itself in front of the white region above the black stripe (see illustration). The meniscus will appear black against the white card. Keeping your eye level with the meniscus, read the buret.
• Remember to estimate one more digit than those marked on the scale.
• Remember that the buret scale reads increasing volume downward, not upward.

Tips on Technique

• Record the starting level to \begin{align*}0.01 \ mL\end{align*} precision (as always, estimate one more digit than marks indicate).
• Titrate the solution of primary acid standard with the \begin{align*}0.1 \ M \ NaOH\end{align*} until faint (see figure) phenolphthalein color appears and persists for \begin{align*}30 \ seconds\end{align*}. (Why might the color slowly disappear even after all acid is titrated?) Record the final buret reading to \begin{align*}0.01 \ mL\end{align*} precision.
• Mix the solution in the titration flask thoroughly after each addition of titrant, to ensure complete reaction before adding more.
• As you near the endpoint, wash the sides of the flask with distilled water to make sure that all delivered titrant is in solution.
• When you see that you are within a drop or two of the endpoint, split drops to avoid overshooting the endpoint.

Lab Procedure

Perform three titrations. For each, calculate the molarity of your \begin{align*}NaOH\end{align*} solution. (From the mass of acid, and its molecular weight, you can calculate the number of moles of acid, which is equal to the number of moles of base you delivered. The molarity is found from the number of moles and the volume.)

When you have three values for the molarity of your \begin{align*}NaOH\end{align*} solution, determine the average value.

Data Table for Part I
Trial 1 Trial 2 Trial 3
Initial Reading, \begin{align*}NaOH\end{align*} buret mL mL mL
Final Reading, \begin{align*}NaOH\end{align*} buret mL mL mL
Volume of \begin{align*}NaOH\end{align*} added mL mL mL
Grams of acid standard g g g
Moles of acid standard mol mol mol
Molarity of \begin{align*}NaOH\end{align*} M M M

Average molarity of \begin{align*}NaOH\end{align*} = _________ M

Procedure for Part II

Part 2. Finding the Molar Mass of an Unknown Acid

• Obtain a sample of a solid unknown from your instructor. Record its ID code in your report.
• Also, record the approximate amount of unknown to use in each titration, and the number of acid hydrogens per molecule. Your instructor will provide this information.
• Weigh the suggested amount into a clean \begin{align*}125-mL\end{align*} erlenmeyer flask, by taring the balance with the flask on the pan, and then adding the acid to the flask. Record the mass of the sample to the precision allowed by the balance. Add \begin{align*}30\end{align*} to \begin{align*}40 \ mL\end{align*} of distilled water and swirl to dissolve your sample. Add three to four drops of phenolphthalein indicator solution and swirl to mix well.
• Titrate your sample with your standardized \begin{align*}NaOH\end{align*} solution until faint phenophthalein color persists for \begin{align*}30 \ seconds\end{align*}. *If your titration requires \begin{align*}10\end{align*} to \begin{align*}25 \ mL\end{align*} of \begin{align*}NaOH\end{align*} solution, carry out a second titration with an unknown sample of about the same mass. Otherwise, adjust the sample mass to bring the expected end-point volume to between \begin{align*}10\end{align*} and \begin{align*}25 \ mL\end{align*} and do two more titrations.
• For each titration, compute the molar mass of the unknown acid, to the precision allowed by your data. Do three titrations and report the average molar mass for the solid acid.

Data for Part II

ID code of acid _______

Number of Acidic Hydrogens in Acid _______

Approximate mass of acid to be used _______g

Trial 1 Trial 2 Trial 3
Mass of unknown acid sample g g g
Volume of \begin{align*}NaOH\end{align*} used mL mL mL
Mols of \begin{align*}NaOH\end{align*} used moL moL moL
Moles of acid present\begin{align*}^1\end{align*} mol mol mol
Molar mass of acid\begin{align*}^2\end{align*} g/mol g/mol g/mol
Molarity of \begin{align*}NaOH\end{align*} M M M

\begin{align*}^1\end{align*} Moles of acid present = \begin{align*}\frac{moles \ NaOH} {number \ of \ H^+ions \ per\ acid\ molecule}\end{align*}

\begin{align*}^2\end{align*} Molar mass of acid = \begin{align*}\frac {grams \ acid} {moles\ acid}\end{align*}

Average Molar Mass = _______g/mol

## Demonstrations for Chapter 21

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