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
Page 1
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
Page 2
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
Page 3
(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
Page 4
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
Hydrobromic Acid
Hydrochloric Acid
Sulfuric Acid
Trichloroacetic Acid
Iodic Acid
Chloroacetic Acid
Ascorbic Acid
Acetic Acid
Propionic Acid
Water
Phenol
Sodium Hydroxide
Ka
1 x 109
1 x 104
1 x 103
1.02 x 10-2
2 x 10-1
1.7 x 10-1
1.4 x 10-3
8 x 10-5
1.8 x 10-5
1.32 x 10-5
1 x 10-7
1.3 x 10-10
1.58 x 10-14
Volume of H+
55.5 mL
57.5 mL
55.5 mL
43.5 mL
---46.0 mL
---38.0 mL
12.0 mL
----------
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
Page 5
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
1 x 103
1.02 x 10-2
0.50 M
5.00 mL
55.5 mL
0.874 M
2 x 10-1
0.50 M
10.00 mL
43.5 mL
----
O
Sulfuric Acid
HO S OH
O
O
Cl
Trichloroacetic Acid
Cl
Cl
Acid Chemistry
C
C
OH
Page 6
O
Iodic Acid
O
I
1.7 x 10-1
0.50 M
5.00 mL
----
----
1.4 x 10-3
0.50 M
10.00 mL
46.0 mL
0.362 M
8 x 10-5
0.50 M
5.00 mL
----
----
1.8 x 10-5
0.50 M
10.00 mL
38.0 mL
0.299 M
OH
O
CH2
Chloroacetic Acid
C
Cl
OH
HO
Ascorbic Acid
HO
O
HO
O
OH
O
Acetic Acid
Acid Chemistry
H3C
OH
Page 7
H3C
Propionic Acid
Water
O
CH2
C
1.32 x 10-5
0.50 M
10.00 mL
12.0 mL
0.0945 M
1 x 10-7
0.50 M
5.00 mL
----
----
1.3 x 10-10
0.50 M
5.00 mL
----
----
1.58 x 10-14
0.50 M
5.00 mL
----
----
OH
H2O
OH
Phenol
Sodium Hydroxide
Acid Chemistry
NaOH
Page 8
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
Page 9
Indicator
Methyl violet
Thymol Blue
Methyl Orange
Methyl Red
Bromocresol Green
Alizarin
Bromothymol Blue
Phenol Red
Cresol Red
Phenolphthalein
Thymolphthalein
Color Below
Lowest pH
Yellow
Red
Yellow
Red
Red
Yellow
Yellow
Red
Yellow
Yellow
Yellow
Colorless
Colorless
pH Range
0.5 to 2.0
1.2 to 2.8
8.2 to 9.1
3.1 to 4.4
4.2 to 6.3
3.8 to 5.4
5.7 to 7.1
11.0 to 12.4
6.0 to 7.6
6.4 to 8.0
7.0 to 8.8
8.0 to 9.8
9.3 to 10.5
Color Above
Highest pH
Violet
Yellow
Blue
Yellow
Yellow
Blue
Red
Purple
Blue
Red
Red
Red
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
Page 10
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
Page 11
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
Acetic Acid – HC2H3O2
Dihydrogen phosphate – HPO42-
Conjugate Base
Acetate ion – C2H3O2Phosphate – 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
Page 12
Overheads
Acid Chemistry
Page 13
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
Page 14
Lesson 2 Definitions
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
Page 15
Lesson 3 Definitions
Buffer – A material that prevents the pH of a solution from
changing drastically upon addition of an acid or a base.
Acid Chemistry
Page 16
Acid Chemistry
Page 17
Acid Chemistry
Page 18
Acid
Hydrobromic Acid
Hydrochloric Acid
Sulfuric Acid
Trichloroacetic Acid
Iodic Acid
Chloroacetic Acid
Ascorbic Acid
Acetic Acid
Propionic Acid
Water
Phenol
Sodium Hydroxide
Acid Chemistry
Ka
1 x 109
1 x 104
1 x 103
1.02 x 10-2
2 x 10-1
1.7 x 10-1
1.4 x 10-3
8 x 10-5
1.8 x 10-5
1.32 x 10-5
1 x 10-7
1.3 x 10-10
1.58 x 10-14
Volume of H+
55.5 mL
57.5 mL
55.5 mL
43.5 mL
---46.0 mL
---38.0 mL
12.0 mL
----------
Page 19
Compound
Compound Structure
Compound
Structure
H3C
O
Hydrobromic Acid
HBr
Iodic Acid
Propionic Acid
O
I
OH
O
CH2
C
OH
O
Hydrochloric Acid
HCl
CH2
Chloroacetic Acid
Water
C
Cl
H2O
OH
HO
OH
O
Sulfuric Acid
HO S OH
Ascorbic Acid
HO
O
O
Phenol
O
HO
O
Cl
Trichloroacetic
Acid
Cl
Cl
Acid Chemistry
C
O
Acetic Acid
C
OH
OH
H3C
OH
Sodium
Hydroxide
NaOH
Page 20
Acid Chemistry
Page 21
Indicator
Methyl violet
Thymol Blue
Methyl Orange
Methyl Red
Bromocresol Green
Alizarin
Bromothymol Blue
Phenol Red
Cresol Red
Phenolphthalein
Thymolphthalein
Acid Chemistry
Color Below
Lowest pH
Yellow
Red
Yellow
Red
Red
Yellow
Yellow
Red
Yellow
Yellow
Yellow
Colorless
Colorless
pH Range
0.5 to 2.0
1.2 to 2.8
8.2 to 9.1
3.1 to 4.4
4.2 to 6.3
3.8 to 5.4
5.7 to 7.1
11.0 to 12.4
6.0 to 7.6
6.4 to 8.0
7.0 to 8.8
8.0 to 9.8
9.3 to 10.5
Color Above
Highest pH
Violet
Yellow
Blue
Yellow
Yellow
Blue
Red
Purple
Blue
Red
Red
Red
Blue
Page 22
Acid
Acetic Acid – HC2H3O2
Dihydrogen phosphate – HPO42-
Acid Chemistry
Conjugate Base
Acetate ion – C2H3O2Phosphate – PO43-
Page 23
Acid Chemistry
Page 24
Laboratory
Experiments
Acid Chemistry
Page 25
Acid Chemistry
Page 26
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
100 mL graduated cylinder
600 mL beaker
Sink or trash can of water
Thermometer
Apparatus used (see next page)3
Acid Chemistry
Page 27
Acid Chemistry
Page 28
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
Acid Chemistry
25.4 torr = 1 in Hg
Page 29
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
Acid Chemistry
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
Page 30
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
Page 31
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
Page 32
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
10 mL graduated cylinder
Stirring rod
Acid Chemistry
Page 33
Procedure:
1.
2.
3.
4.
5.
6.
7.
8.
Obtain 6 test tubes and label them #1 - #6.
Place 1.0 mL of 0.10 M acetic acid in each of the six test tubes.
Add 2 drops of crystal violet to test tube #1.
Add 2 drops of thymol blue to test tube #2.
Add 2 drops of methyl orange to test tube #3.
Add 2 drops of congo red to test tube #4.
Add 2 drops of malachite green to test tube #5.
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
Page 34
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
Page 35
Acid Chemistry
Page 36
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
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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
Acid Chemistry
Color of Solution after
Cabbage Juice Addition
Approx. pH
Page 38
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
Page 39
Acid Chemistry
Page 40
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
Page 41
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
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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
Page 43
Acid Chemistry
Page 44
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
Page 45
Procedure:
1.
2.
3.
4.
5.
6.
7.
8.
Using a 25-mL graduated cylinder add 25-mL of distilled water to one of the beakers
Using a 25-mL graduated cylinder add 25-mL of buffered solution to the second beaker.
Wash the 25-mL graduated cylinder.
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.
Wash the 25-mL graduated cylinder.
Add 6 drops of universal indicator to each beaker. Record the color of the solutions.
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.
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.
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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 __________________
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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)
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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?
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