by: michael wild, matt huber, jasmine gilbert and dr...

Acid Chemistry

Acid Chemistry By: Michael Wild, Matt Huber, Jasmine Gilbert and Dr. Faith Yarberry

In this module the student will:

Understand the concept of an Acid.

Discover the differences between strong acids and weak acids.

Learn the meaning of the term pH with respect to acid / base strength.

Uncover what an indicator is and how it works.

Identify the strength of an acid and a base in various household products via pH.

Be able to define a buffer, prepare a buffer, and understand the purpose of a buffer.

Acid Chemistry

Lesson 1: Strengths of Acids

The Bronsted/Lowry Definition of an acid is:

Acid – a substance that donates a proton.

Proton – the description used for a hydrogen cation (H+) because once an electron is

removed from a hydrogen atom, all that is left is a proton)

Example:

Hydrochloric acid HCl → H+ + Cl

-

When an acid is dissolved in water, hence aqueous (aq), the hydrogen cation is donated to the

water forming a hydronium ion (H3O+).

Hydrochloric acid HCl (aq) → H3O+ + Cl

-

Not all acids are created equally.

Strong acid – an acid that dissociates entirely in water to produce a hydronium ion and an

anion. The diagram below shows this relationship.1

Acid Chemistry

Weak acid – an acid that only partially dissociates in water to produce a hydronium ion

and an anion. The diagram below shows this relationship.1

Weak acids reach a state of equilibrium.

Equilibrium – The point where the concentrations of the reactants and the concentration of

the products cease to change, but the reaction itself continues.

(1) Initially all that is present are reactants.

Acid Chemistry

(2) The reaction moves in a forward direction building the concentration of products.

(3) As the concentration of the products build the reverse reaction begins to occur.

(4) The forward and the reverse reactions occur simultaneously. The forward reaction

continues to make product while the reverse reaction continues to make reactants. At

some point the concentration of the reactants and those of the products will cease to

change. This point is called equilibrium.

Acid equilibrium constants (Ka) – indicate the degree to which an acid will dissociate. The

diagram below indicates the relationship between Ka and dissociation.1

Acid Chemistry

In the laboratory you are going to prove that the stronger the acid, the greater its reactivity due

to the quantity of free H+. To do so, you will be reacting magnesium metal with hydrochloric

acid and with acetic acid. You will be collecting the amount of hydrogen produced by each

reaction. From the amount of hydrogen produced you will be able to identify the strength of the

acid. (LAB 1)

Following the lab put up the following table to prove that what they observed was not just a

result for acetic acid and hydrochloric acid. The value for trichloroacetic acid is off due to the

fact that this material is hygroscopic. Being hygroscopic the concentration of the original

solution not likely to be near the 5.0 M region as with the other materials.

Acid Ka Volume of H+

Hydrobromic Acid 1 x 109 55.5 mL

Hydrochloric Acid 1 x 104 57.5 mL

Sulfuric Acid 1 x 10

3

1.02 x 10-2 55.5 mL

Trichloroacetic Acid 2 x 10-1

43.5 mL

Iodic Acid 1.7 x 10-1

----

Chloroacetic Acid 1.4 x 10-3

46.0 mL

Ascorbic Acid 8 x 10-5

----

Acetic Acid 1.8 x 10-5

38.0 mL

Propionic Acid 1.32 x 10-5

12.0 mL

Water 1 x 10-7

----

Phenol 1.3 x 10-10

----

Sodium Hydroxide 1.58 x 10-14

----

Possible questions:

(1) What reason can you give as to why sulfuric acid generates nearly twice the volume of

H+ compared to hydrochloric acid and hydrobromic acid even though all three are strong

acids? (Structure)

(2) What reason can you give as to why this method may not have worked for determining

the H+ content in iodic acid even though its Ka resides between that of trichloroacetic acid

and acetic acid? (Structure)

(3) What do the weak acids have in common that did react with magnesium? (Structure)

Acid Chemistry

Compound Compound Structure Ka Conc. of

Compound

Volume of

Compound

Used

Volume of H2

Generated in 4

minutes

Conc of H+

Hydrobromic Acid HBr 1 x 109

0.50 M 10.00 mL 55.5 mL 0.437 M

Hydrochloric Acid HCl 1 x 104 0.50 M 10.00 mL 57.5 mL 0.453 M

Sulfuric Acid S

O

OH

O

OH

1 x 103

1.02 x 10-2 0.50 M 5.00 mL 55.5 mL 0.874 M

Trichloroacetic Acid C C

Cl

O

OH

ClCl

2 x 10-1

0.50 M 10.00 mL 43.5 mL ----

Acid Chemistry

Iodic Acid I

O

OHO

1.7 x 10-1

0.50 M 5.00 mL ---- ----

Chloroacetic Acid CH2 C

Cl

O

OH

1.4 x 10-3

0.50 M 10.00 mL 46.0 mL 0.362 M

Ascorbic Acid O O

OHOH

OH

OH

8 x 10-5

0.50 M 5.00 mL ---- ----

Acetic Acid CH3

O

OH

1.8 x 10-5

0.50 M 10.00 mL 38.0 mL 0.299 M

Acid Chemistry

Propionic Acid CH2 C

O

OH

CH3

1.32 x 10-5

0.50 M 10.00 mL 12.0 mL 0.0945 M

Water H2O 1 x 10-7

0.50 M 5.00 mL ---- ----

Phenol

OH

1.3 x 10-10

0.50 M 5.00 mL ---- ----

Sodium Hydroxide NaOH 1.58 x 10-14

0.50 M 5.00 mL ---- ----

Acid Chemistry

Lesson 2: pH and Indicators

A method for describing the concentration of H+ in solution is pH.

pH – a term used to describe the acidity of a solution through the detection of H+

concentration in solution

A neutral solution has a pH of 7.0. pH’s below 7.0 indicate that the solution is acidic and pH’s

above 7.0 indicate that the solutions is basic. See the diagram below.1

The pH of a solution can be determined using an indicator or pH probes. Of these methods the

pH probe will give the greatest amount of accuracy regarding pH, but pH probes can be

expensive. A much less expensive method for determining the pH of a solution requires the use

an indicator.

Indicator – A substance that changes colors depending on the pH of the solution.

pH indicators can be substances such as chemical compounds, litmus papers etc… that when

added in small quantities to a solution visually allow us to determine the relative acidity or

basicity of a solution by undergoing a change in color. Indicators are usually, themselves, either

a weak acid or base that detect the hydrogen ion concentration in solution. Below is a table of

indicators, the color of that indicator below the lowest pH indicated, the range for which a color

change will be detectable, and the color of that indicator above the highest pH indicated.2

Acid Chemistry

Indicator Color Below

Lowest pH pH Range

Color Above

Highest pH

Methyl violet Yellow 0.5 to 2.0 Violet

Thymol Blue Red 1.2 to 2.8 Yellow

Yellow 8.2 to 9.1 Blue

Methyl Orange Red 3.1 to 4.4 Yellow

Methyl Red Red 4.2 to 6.3 Yellow

Bromocresol Green Yellow 3.8 to 5.4 Blue

Alizarin Yellow 5.7 to 7.1 Red

Red 11.0 to 12.4 Purple

Bromothymol Blue Yellow 6.0 to 7.6 Blue

Phenol Red Yellow 6.4 to 8.0 Red

Cresol Red Yellow 7.0 to 8.8 Red

Phenolphthalein Colorless 8.0 to 9.8 Red

Thymolphthalein Colorless 9.3 to 10.5 Blue

By using a combination of indicators it is possible to determine the relative pH. Lab #2.

There are many plants that contain chemicals that can be used as natural pH indicators such as

the anthocyanin compound family. Red cabbage is part of that anthocynanin containing family.

These compounds will usually turn acidic solutions a red color, basic solutions a green-yellow

color, and neutral solutions a purple color. Due to this result it is possible that one can use red

cabbage to find the pH of a solution as seen in the photograph below.

.

Acid Chemistry

At this point have the students bring household products for class the following day to test the

pH of the solution in Lab 3.

We have stated that the pH of the solution allows us to determine the concentration of the H+ in

solution. What is the connection?

pH = -log [H+]

Your assignment tomorrow, is to determine the approximate H+ concentration in each of the

household solutions studied today.

10-pH

= [H+]

The next day the students will come in with the H+ concentration determined. Ideally you should

come in with the real H+ concentration of the solution (pH probe). That way the students can see

that the indicator does a fairly good job of predicting the H+ concentration.

Acid Chemistry

Lesson 3: Buffers

Some solutions, called buffers, are remarkably resistant to changes in pH. Water is not a

buffer, since its pH is very sensitive to the addition of acidic or basic species. A good buffer,

again, is a system that will prevent drastic changes in pH. Buffers are usually solutions of a weak

acid and their conjugate base. A conjugate base results when the weak acid loses its acidic

hydrogen. Below are a couple of examples.

Acid Conjugate Base

Acetic Acid – HC2H3O2 Acetate ion – C2H3O2-

Dihydrogen phosphate – HPO42-

Phosphate – PO43-

In lab today you are going to prove that a buffer is used to prevent pH changes. Lab 4 and/or 5.

Acid Chemistry

Overheads

Acid Chemistry

Lesson 1 Definitions

Acid – a substance that donates a proton.

Proton – the description used for a hydrogen cation (H+) because

once an electron is removed from a hydrogen atom, all that is left is

a proton)

Strong acid – an acid that dissociates entirely in water to produce a

hydronium ion and an anion.

Weak acid – an acid that only partially dissociates in water to

produce a hydronium ion and an anion.

Equilibrium – The point where the concentrations of the reactants

and the concentration of the products cease to change, but the

reaction itself continues.

Acid equilibrium constants (Ka) – indicate the degree to which an

acid will dissociate.

Acid Chemistry


pH – a term used to describe the acidity of a solution through the

detection of H+ concentration in solution

Indicator – A substance that changes colors depending on the pH of

the solution.

pH = -log [H+]

10-pH

= [H+]

Acid Chemistry


Buffer – A material that prevents the pH of a solution from

changing drastically upon addition of an acid or a base.

Acid Chemistry

Acid Chemistry

Acid Ka Volume of H+

Hydrobromic Acid 1 x 109

55.5 mL

Hydrochloric Acid 1 x 104 57.5 mL

Sulfuric Acid 1 x 10

3

1.02 x 10-2 55.5 mL

Trichloroacetic Acid 2 x 10-1

43.5 mL

Iodic Acid 1.7 x 10-1

----

Chloroacetic Acid 1.4 x 10-3

46.0 mL

Ascorbic Acid 8 x 10-5

----

Acetic Acid 1.8 x 10-5

38.0 mL

Propionic Acid 1.32 x 10-5

12.0 mL

Water 1 x 10-7

----

Phenol 1.3 x 10-10

----

Sodium Hydroxide 1.58 x 10-14

----

Acid Chemistry

Compound Compound Structure Compound Structure

Hydrobromic Acid HBr Iodic Acid I

O

OHO

Propionic Acid CH2 C

O

OH

CH3

Hydrochloric Acid HCl Chloroacetic Acid CH2 C

Cl

O

OH

Water H2O

Sulfuric Acid S

O

OH

O

OH

Ascorbic Acid O O

OHOH

OH

OH

Phenol

OH

Trichloroacetic

Acid C C

Cl

O

OH

ClCl

Acetic Acid CH3

O

OH

Sodium

Hydroxide NaOH

Acid Chemistry

Acid Chemistry

Indicator Color Below

Lowest pH pH Range

Color Above

Highest pH

Methyl violet Yellow 0.5 to 2.0 Violet

Thymol Blue Red 1.2 to 2.8 Yellow

Yellow 8.2 to 9.1 Blue

Methyl Orange Red 3.1 to 4.4 Yellow

Methyl Red Red 4.2 to 6.3 Yellow

Bromocresol Green Yellow 3.8 to 5.4 Blue

Alizarin Yellow 5.7 to 7.1 Red

Red 11.0 to 12.4 Purple

Bromothymol Blue Yellow 6.0 to 7.6 Blue

Phenol Red Yellow 6.4 to 8.0 Red

Cresol Red Yellow 7.0 to 8.8 Red

Phenolphthalein Colorless 8.0 to 9.8 Red

Thymolphthalein Colorless 9.3 to 10.5 Blue

Acid Chemistry

Acid Conjugate Base

Acetic Acid – HC2H3O2 Acetate ion – C2H3O2-

Dihydrogen phosphate – HPO42- Phosphate – PO4

3-

Acid Chemistry

Acid Chemistry

Laboratory

Experiments

Acid Chemistry

Acid Chemistry

Laboratory #1

Preparation (for nine groups):

0.5 M HCl – Slowly add 8.3 mL of 6 M hydrochloric acid to approximately 50 mL of distilled

water in a 100 mL volumetric flask. Cap and shake. Then add distilled water to the line, cap and

shake.

0.5 M Acetic Acid – Slowly add 2.90 mL of glacial acetic acid to approximately 50 mL of

distilled water in a 100 mL volumetric flask. Cap and shake. Then add distilled water to the

line, cap and shake.

Or

0.5 M Acetic Acid – Add 59.41 mL of vinegar to a 100 mL volumetric flask. Add some distilled

water, cap and shake. Then add distilled water to the line, cap and shake.

Magnesium strips cut into small pieces.

Determine the barometric pressure

Stop Watch

Materials Needed Per Group

10 mL graduated cylinder

Large test tube

Small test tube

Large Erlenmeyer flask

Rubber tubing with stopper and glass tubing


600 mL beaker

Sink or trash can of water

Thermometer

Apparatus used (see next page)3

Acid Chemistry

Acid Chemistry

Background:

An acid, when reacted with magnesium metal, dissociates to form hydrogen gas. This gas can be

collected and evaluated to determine the relative strength of an acid. This lab will allow you to

determine the relative acidities of two different acids; one will be a strong acid and the other a

weak acid.

The method used in this laboratory involves collecting the gas generated by the magnesium /

acid reaction above water in a closed container. In order to determine the amount of hydrogen

collected, and hence the strength of the acid, you will be using the ideal gas law to solve for the

number of moles of hydrogen generated.

The ideal gas law is: PV = nRT, where P stands for the pressure of the gas, V the volume of the

gas, n the number of moles of gas, R the ideal gas constant, and T the temperature of the gas in

Kelvin. The ideal gas constant comes in many forms, but the one that you will be using for this

lab is R = 0.0821 atm L / mol K. The ideal gas constant dictates the units that are acceptable for

pressure (atm), volume (L), and temperature (K).

The volume and temperature, associated with the gas, will be simple to determine. The volume

of the gas will be measured directly from the 100-mL graduated cylinder. This method of

measurement is acceptable since gases spread out to fill the space available. The temperature of

the gas will be the same as the temperature of the water that the gas bubbled through. The

pressure is the most complicated of the variables to determine.

The surface molecules of a liquid absorb energy to break free of the remaining liquid molecules

and enter into the gas phase. This energy is provided by heat. In this experiment, the amount of

water that will enter into the gas phase is determined by the temperature of the water. The higher

the temperature of the water, the greater the amount of water molecules that will break free from

the liquid phase to enter into the gas phase. So the gas collected in the 100-mL graduated

cylinder will be a combination of water (water vapor) and hydrogen. According to Dalton’s Law

of Partial Pressure, the overall pressure associated with a gas is equal to the sum of the pressures

of every gas involved. In this experiment, therefore, the total pressure of the gas, which is

barometric pressure, will be the sum of the pressure of the water molecules (water vapor) plus

the pressure of the hydrogen gas.

Ptotal = Pwater + Phydrogen

To solve for the pressure of the hydrogen gas, the equation can be rearranged to give:

Ptotal – Pwater = Phydrogen

Once the pressure, volume, and temperature of the hydrogen has been determined, the ideal gas

law can be used to calculate the number of moles of hydrogen generated by the acid.

Important pressure conversion units:

1 atm = 760 torr 1 atm = 760 mm Hg 25.4 torr = 1 in Hg

Acid Chemistry

Procedure

1. Lower a 600-mL beaker into a filled sink of water.

2. Lower a 100-mL graduated cylinder into a filled sink of water. Allow the cylinder to

completely fill with water so that no air bubbles remain in the cylinder.

3. While the 600-mL beaker and the 100-mL graduated cylinder are under water, place the

cylinder into the beaker.

4. Measure the temperature of the water in the beaker.

5. Obtain approximately 0.200 grams of magnesium and record the weight.

6. Place the magnesium in a small test tube.

7. Obtain a clean, dry 10-mL graduated cylinder. Place 10-mL of acid into the cylinder and

record the volume to the nearest 0.1 mL.

8. Pour the acid into the large test tube.

9. Gently slide the small test tube into the large test tube. Make sure that none of the acid

splashes into the tube with the magnesium.

10. Stand the large test tube in an Erlenmeyer flask.

11. Put the stopper end of the tubing into the large test tube so that the seal is tight.

12. Place the other end of the tubing with the curved glass gently under the lip of the 100-mL

graduated cylinder. Be careful not to let any air into the cylinder during this step.

13. Carefully mix the acid and magnesium by slowly tilting the large test tube. Bubbling will

occur.

14. Continue mixing for 4 minutes.

15. Read the volume of the gas in the graduated cylinder.

16. Carefully pour the resulting solution down the drain capturing any magnesium that

remains.

17. Dry the magnesium and weigh it. Record the new weight.

18. Clean your glassware and repeat for the next acid. (You may use the magnesium left from

step 15 for your next acid.)

Water Vapor Table

Temperature in oC Water Vapor Pressure in Torr

18 15.5

19 16.5

20 17.5

21 18.7

22 19.8

23 21.1

24 22.4

25 23.8

26 25.2

27 26.7

28 28.3

Acid Chemistry

Data and Results

Acid 1 Acid 2

Temperature in oC

Temperature in K

Initial Mass of Magnesium in g

Final Mass of Magnesium in g

Mass of Magnesium used

Atmospheric Pressure

Atmospheric Pressure in atm

Water vapor press at this temperature

Water vapor in atm

Pressure of H2 (atmospheric pressure

in atm – water vapor pressure in atm)

Volume of gas collected in mL

Volume of gas in L

Ideal Gas Constant

Moles of Hydrogen (H2)

Acid Chemistry

Conclusions and Post – Lab Questions

1. What is the concentration of the H

+ in solution for Acid 1?

a. You used 10 mL of acid solution in this experiment. What is the volume in liters?

b. How many moles of H+ is present in the solution if 2 moles of H

+ gives 1 mole of

H2?

c. What is the molarity of the H+ in solution if Molarity (M) equals moles of solute

divided volume of solution?

2. What is the concentration of the H+ in solution for Acid 2?

a. You used 10 mL of acid solution in this experiment. What is the volume in liters?

b. How many moles of H+ is present in the solution if 2 moles of H

+ gives 1 mole of

H2?

c. What is the molarity of the H+ in solution if Molarity (M) equals moles of solute

divided volume of solution?

3. Which acid, Acid 1 or Acid 2, was the strongest? Explain.

Acid Chemistry

Laboratory #2

Preparation (for nine groups):

0.1 M Acetic Acid – Slowly add 0.57 mL of glacial acetic acid to approximately 50 mL of

distilled water in a 100 mL volumetric flask. Cap and shake. Then add distilled water to the

line, cap and shake.

0.01% Malachite Green indicator – Dissolve 0.01 g Malachite Green in enough water to make

100 mL of solution.

0.10% Congo Red indicator – Dissolve 0.1 g Congo Red in enough water to make 100 mL of

solution.

1.0 % Crystal Violet indicator – Dissolve 1 g Crystal Violet in enough water to make 100 mL of

solution.

0.1% Methyl Orange indicator – Dissolve 0.1 g Methly Orange in enough water to make 100 mL

or solution.

0.04% Thymol Blue indicator – Dissolve 0.04 g Thymol Blue in enough water to make 100 mL

of solution.

0.02% Methyl Red indicator – Dissolve 0.02 g Methyl Red in enough water to make 100 mL of

solution.

Materials Needed Per Group

6 test tubes


Stirring rod

Acid Chemistry

Procedure:

1. Obtain 6 test tubes and label them #1 - #6.

2. Place 1.0 mL of 0.10 M acetic acid in each of the six test tubes.

3. Add 2 drops of crystal violet to test tube #1.

4. Add 2 drops of thymol blue to test tube #2.

5. Add 2 drops of methyl orange to test tube #3.

6. Add 2 drops of congo red to test tube #4.

7. Add 2 drops of malachite green to test tube #5.

8. Add 2 drops of methyl red to test tube #6.

Data

For each indicator, color in the box on that line that most accurately describes the color you

observed in that test tube.

Acid Chemistry

Conclusions

1. According to Crystal Violet the pH was between _____________________.

2. According to Thymol Blue the pH was between _____________________.

3. According to Methyl Orange the pH was between _____________________.

4. According to Congo Red the pH was between _____________________.

5. According to Malachite Green the pH was between _____________________.

6. According to Methyl Red the pH was between _____________________.

7. According to all indicators the pH of this solution was between ___________________.

8. Explain how indicators can be used to discern the pH of a solution.

Acid Chemistry

Acid Chemistry

Laboratory #3

Preparation:

Pull and tear the cabbage leaves into small pieces until you fill a large beaker or another

container and add water to submerge the leaves. Then add heat to boil the water for

approximately 5 to 7 minutes or till the color of the water is purple like the cabbage. Pour up this

solution into dropper bottles.

Make a laminated sheet of the photograph showing the color of cabbage juice in solution

depending on ph for each station.

Have your students bring in a sample of a household product that is clear and relatively colorless.

Materials Needed per Station:

5-8 test tubes per group

Droppers for each sample they will be testing

Dropper bottle of cabbage juice solution

Photograph showing the color of cabbage juice in solution depending on pH

Acid Chemistry

Procedure

1. Add 1 dropper full of the solution to be analyzed to a test tube.

2. Add 5 drops of cabbage juice extract to the test tube.

3. Compare to the photograph to determine the pH of the household product.

Data

Household Product Color of Solution after

Cabbage Juice Addition Approx. pH

Acid Chemistry

Conclusions and Post-Lab Questions

1. If the pH of your household product is less than 7, what does it tell you about the amount

of H+ in your household product compared to those with a pH greater than 7?

2. Place the household chemicals that you tested in order of increasing acidity.

3. Which of your household products has the greatest amount of H+ in solution?

Acid Chemistry

Acid Chemistry

Laboratory #4

Preparation: (per 20 groups)

0.1 M acetic acid - Dilute 60.24 mL of store bought vinegar solution to make 500 mL of

solution.

0.1 M acetate – dissolve 6.56 g of sodium acetate in enough water to make 800 mL of solution.

Materials:

100-mL graduated cylinder

150-mL beaker or a glass

Straw

0.1 M acetic acid

0.1 M acetate solution

pH probe

Discussion:

They should observe that the pH drops. This is due to the fact that carbon dioxide, exhaled by

the body, when combined with water makes carbonic acid.

CO2 + H2O → H2CO3

The addition of carbonic acid to water that includes a buffer does not affect the pH to appreciable

amounts.

Acid Chemistry

Procedure:

1. Fill a glass with 60 mL of distilled water

2. Check the pH with a pH probe and record the pH.

3. Using a straw, blow bubbles into the water for 1 minute

4. Retest the pH and record.

5. Clean out your glass.

6. Add 15 mL of 0.1 M acetic acid to the glass.

7. Add 15 mL of 0.1 M acetate to the glass.

8. Add 30 mL of distilled water to the mixture.

9. Check the pH with a pH probe and record the pH.

10. Using a straw, blow bubbles into the water for 1 minute.

11. Retest the pH and record.

12. Clean out your glass.

Data

pH of distilled water ___________________________

pH of water after 1 minute of bubbles ___________________________

pH of the 15 mL/15 mLbuffer solution ___________________________

pH of buffer solution after 1 minute of bubbles ___________________________

Acid Chemistry

Post-Laboratory Questions

1. Give the name and formula for the acid that you used to make the buffer solution.

2. Give the name and formula of the conjugate base used to make the buffer.

3. How do the acid and the conjugate base differ from one another?

4. Given your observations, explain how a buffer works.

5. What do you think will happen to the pH of your solution when the buffer is used up?

6. Given your observations from the first experiment, explain why it is essential that all living

systems contain buffers. Give an example.

Acid Chemistry

Acid Chemistry

Laboratory #5

Preparation: (Enough for 14 groups)

Buffer solution – add 0.75 g baking soda to 350-mL of distilled water.

Acid Rain solution – add 4 mL of 1 M H2SO4 to 2 L of distilled water.

Universal Indicator - Pull and tear red cabbage leaves into small pieces until you fill a large

beaker or another container and add water to submerge the leaves. Then heat to boil the water for

approximately 5 to 7 minutes or till the color of the water is purple like the cabbage. Pour up this

solution into dropper bottles.

Materials:

Distilled water

Buffer solution

Acid Rain solution

Universal indicator solution

3 – 250 mL beakers

1 – 25 mL graduated cylinder

1 – 10 mL graduated cylinder

Dropper

Safety goggles

1 M hydrochloric acid

Acid Chemistry

Procedure:

1. Using a 25-mL graduated cylinder add 25-mL of distilled water to one of the beakers

2. Using a 25-mL graduated cylinder add 25-mL of buffered solution to the second beaker.

3. Wash the 25-mL graduated cylinder.

4. Using a 25-mL graduated cylinder add 25-mL of 1 M hydrochloric acid to the third

beaker for defining what is meant by the color pink.

5. Wash the 25-mL graduated cylinder.

6. Add 6 drops of universal indicator to each beaker. Record the color of the solutions.

7. Using a dropper add the acid rain solution drop-by-drop to the beaker containing distilled

water. Swirl after each addition and record the color. Continue until the color turns pink

and remains stable.

8. Using a 10-mL graduated cylinder, add 10-mL of acid rain solution to the beaker with the

buffered solution. Swirl after each addition and record the color of the solution. Continue

adding acid rain until the solution turns pink and remains stable.

Acid Chemistry

Data:

Color of HCl solution _______________________

Distilled Water Buffer Solution

Initial Color

Color After:

________________________

1 drops __________________

2 drops __________________

3 drops __________________

4 drops __________________

5 drops __________________

6 drops __________________

7 drops __________________

8 drops __________________

9 drops __________________

10 drops __________________

11 drops __________________

12 drops __________________

13 drops __________________

14 drops __________________

15 drops __________________

__________________________

10 mL __________________

20 mL __________________

30 mL __________________

40 mL __________________

Acid Chemistry

50 mL __________________

60 mL __________________

70 mL __________________

80 mL __________________

90 mL __________________

100 mL __________________

110 mL __________________

120 mL __________________

130 mL __________________

140 mL __________________

150 mL __________________

Volume of Acid Used:

(assume 20 drops equals 1 mL)

Acid Chemistry

Post Laboratory Questions

1. Why was there a difference in the amount of acid rain needed to change the pH of these

two solutions?

2. What is a buffer?

3. Acidosis in the human body occurs when the pH of the blood drops below 7.35.

Alkalosis in the human body occurs when the pH goes above 7.45. Both are dangerous

and can cause death. During a regular day an individual is likely to have a carbonated

beverage with a pH around 3 or coffee with a pH of 5. What does the blood contain so

that the things that we ingest do not harm us?

by: michael wild, matt huber, jasmine gilbert and dr...

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