Chem 110 Lab
Acids & Bases
Clark College
Name: ______________________________
Partner’s Name: _______________________
LEARNING OBJECTIVES After completing this experiment, you should feel comfortable with:
•
Using burets.
•
Experimentally assessing acids and bases.
•
Examining reaction speeds
•
The concept of pH.
In Part I you will titrate acetic acid in commercially available vinegar with NaOH. You will pay
careful attention to volumes as well as speed of the reaction. In Part II you will explore the pH and
reaction speed of a different acid base titration.
GENERAL BACKGROUND
Definitions
Acid: A substance that produces H3O+ ions when placed in water, or that donates H+ to another
substance. (Although the bare proton, H+, is unlikely to be present in solution, we will use
H+ and H3O+ interchangeably as a matter of convenience or shorthand notation)
Base: A substance that produces OH- ions when dissolved in water or that accepts H+ from another
substance.
Strong/Weak Acids and Bases:
Strong acids and bases are completely or nearly 100% ionized when dissolved in water. Strong
acids include hydrogen halides, HCl, HBr, HI, HNO3, H2SO4, and HClO4. Strong bases include
group IA and IIA hydroxides, LiOH, NaOH, KOH, Mg(OH)2, Ca(OH)2, and Ba(OH)2. The group IIA
metals are not as soluble as the group IA metals but the dissolved compound does dissociate into
ions.
The majority of acids and bases are weak, meaning that most of the molecules dissolved in water
remain as molecules and only a small percentage dissociates into ions. All organic acids and bases
are weak acids and bases. Examples of weak acids include the organic acids acetic, ascorbic,
carbonic, citric, and tartaric acids, and inorganic acids such as phosphoric acid. Examples of weak
bases include ammonia and glycine.
Discussion
When a acid and a base are combined, they do so via the double replacement reaction pathway –
and a salt and water are often produced. For example, when hydrochloric acid is combined with
sodium hydroxide, we produce sodium chloride and water.
HCl + NaOH → NaCl + H2O
This reaction is a specific type of double replacement reaction termed neutralization.
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Note that the two reactants in this example combine in a 1:1 ratio. If one mole of HCl is mixed with
one mole of NaOH in aqueous solution, they will exactly combine to form a solution with neither
acid nor base in excess, hence the term neutral. If there is an excess of either the acid or the base, the
resulting solution will not be neutral.
Titration is an analytical procedure to determine the concentration of an "unknown" solution by
reacting it with a solution of a known concentration. When neutralization is carried out measuring
the exact volumes of both acid and base used to reach the equivalence point, the process is called
titration. The equivalence point is the point at which the volume of solution added will exactly
neutralize the other solution.
At the equivalence point:
moles OH- added = moles H3O+ added
This point is often determined using an indicator that turns color at a pH close to the equivalence
point. Therefore, the titration is performed to the point that the indicator changes, or to an
endpoint.
Molarity (moles of solute per liter of solution) is widely used for specifying concentration.
PART I. TITRATION OF VINEGAR
In Part I of today's lab you will titrate acetic acid (CH3COOH) in vinegar with a common household
base, sodium hydroxide (also known as lye).
You will be supplied with an alkaline (basic) solution of NaOH of a known concentration. This
concentration has been determined by our laboratory technician.
The reaction you will observe is given below.
HC2H3O2 (aq) + NaOH (aq) → NaC2H3O2 (aq) + H2O (l)
Note that while the acetic acid has 4 total hydrogens, only one ionizes and is therefore removed in
an acid-base reaction.
In general, the change from an acidic to basic solution is not a visible or otherwise obvious change.
The simplest and cheapest way to observe the change is to add a small amount of an indicator, a
substance that changes color in response to the H+ or OH- concentration of the solution. The
indicator you will be using is phenolphthalein, a substance that is clear in acidic solution and turns
pink in basic or alkaline solution.
The volume of the base solution necessary to just barely produce the color change in the indicator is
determined by adding the base solution from a buret to an erlenmeyer flask containing a specific
volume of the acid solution and several drops of the indicator. Addition is fairly rapid at first and
then is slowed to drop-wise as the end point is approached. Dropwise addition continues until the
indicator just barely changes color. The total volume of the standard base solution added from the
buret to achieve neutralization is then recorded to the nearest 0.01 mL.
EXPERIMENTAL PROCEDURE
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1. Pour about 60-70 mL (do not get extra!) of the standard sodium hydroxide solution into a clean,
dry 100-mL beaker. Pour about 20 mL vinegar into a 100 mL beaker (to be used in step 5 below).
2. Select a buret from the lab counter. Rinse it twice with distilled water to be sure it is clean. Rinse
the buret with two separate portions consisting of a few milliliters of the sodium hydroxide
solution you obtained. Drain the rinse solutions through the stopcock. Support the buret on a
stand as shown by your instructor.
3. Now fill the buret with the standard sodium hydroxide solution. Fill the tip of the buret by
opening the stopcock momentarily and allowing a small amount of the standard sodium
hydroxide solution to flow through the tip into a beaker. After the tip is filled, adjust the level of
the solution in the buret so that the bottom of the curved surface (meniscus) is opposite the scale
near the zero marking on the buret – note – it need not be at EXACTLY 0! Be sure that your eye
is level with the meniscus when you read the buret. Record the initial buret reading to the
nearest 0.01 mL. Ask your instructor to initial your first buret reading on your worksheet
before you proceed!
4. Record the molarity of the standard sodium hydroxide solution on your data sheet, using the
accuracy indicated on the label of the solution.
5. Using a graduated cylinder, pour 10.0 mL of the vinegar solution into a 125 mL erlenmeyer
flask. If you have a difficult time with the pipet, then you should get as close as you can to 10.00
mL, recording the exact amount you use to the nearest 0.01 mL on the data sheet.
6. Add 2 drops of phenolphthalein solution to the vinegar in the erlenmeyer flask.
7. Begin your titration by slowly adding into the flask 15 to 20 mL of base from your buret. Notice
that a pink color forms but disappears quickly when the erlenmeyer flask is swirled to allow
mixing of the base with the acid. Slowly drip in more of the base solution while gently swirling
the erlenmeyer flask until the pink color disappears following each addition of base.
8. When the color remains longer in the solution, add the base solution drop by drop with
swirling until one drop produces a persistent pale pink color in the solution.
9. When the end point is reached (pale pink color holds) record the final buret reading to the
nearest 0.01 mL in the data table supplied below.
10. When you clean up, please rinse the buret very well, but do not use any soap solutions. Dispose
of all waste down the drain.
11. Obtain a second set of data from another group and record in the table below. If you do not like
their data, if you do not agree with their data, then do not use it!
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DATA TABLE 1:
concentration
NaOH, M
Buret Readings, mL
Initial
Final
total mL
NaOH used
mL vinegar
solution
(tenths place)
Titration 1
Titration 2
Titration 2
names:
Calculation:
Determine the average volume of NaOH needed to neutralize the acetic acid solution. Show a
complete set-up for your calculation in the space provided below.
PART II. TITRATION OF MAGNESIUM HYDROXIDE
In Part II of today's lab you will titrate stomach acid (HCl) with antacid – Mg(OH)2 (aka one of the
components in Milk of Magnesia).
You will be supplied with solid Mg(OH)2 and a known concentration of HCl solution. This
concentration has been determined by our laboratory technician.
The reaction you will observe is given below.
2HCl (aq) + Mg(OH)2 (s) → MgCl2 (aq) + 2H2O(l)
Note that the acid solution will be reacting with a SOLID that is present in solution. You will notice
that this reaction takes longer. As you are titrating – observe what is happening to the solid as the
titration proceeds so you might comment on why this reaction takes so much more time to
complete than the vinegar/NaOH reaction.
For this reaction, we will use a different indicator – universal indicator as we watch the reaction
proceed. Placed on the instructor station is a universal indicator “key” which shows the color the
universal indicator is as the pH of the solution changes.
Neutralization of the acid and base occurs when the pH = 7. Can you hit that color?
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EXPERIMENTAL PROCEDURE
1. Pour about 60-70 mL (do not get extra!) of the standard HCl solution into a clean, dry 100-mL
beaker.
2. Measure out ~ 0.1 grams of Mg(OH)2 using the balances. Record the exact mass of Mg(OH)2
you weigh out in the table below. Dissolve the Mg(OH)2 solid into ~ 20 mL of water. Add 10
drops of universal indicator to the solution and record its starting color and initial pH.
3. Select a buret from the lab counter. Rinse it twice with distilled water to be sure it is clean. Rinse
the buret with two separate portions consisting of a few milliliters of the hydrochloric acid
solution you obtained. Drain the rinse solutions through the stopcock. Support the buret on a
stand.
4. Now fill the buret with the HCl solution. Fill the tip of the buret by opening the stopcock
momentarily and allowing a small amount of the solution to flow through the tip into a beaker.
5. After the tip is filled, adjust the level of the solution in the buret so that the bottom of the curved
surface (meniscus) is opposite the scale near the zero marking on the buret –note – it need not
be at EXACTLY 0! Be sure that your eye is level with the meniscus when you read the buret.
Record the initial buret reading to the nearest 0.01 mL. Ask your instructor to initial your first
buret reading on your worksheet before you proceed!
6. Record the molarity of the HCl solution on your data sheet, using the accuracy indicated on the
label of the solution.
7. Begin your titration by slowly adding into the flask 5 to 8 mL of acid from your buret. Notice
the color change. Observe the solid Mg(OH)2. Do not titrate more acid until the color stops
changing.
8. Continue titrating 3-5 mL of acid at a time pausing between acid additions to make note of the
color, the color change, and the appearance of the solid.
9. When the end point is reached – you are shooting for pH = 7! Record the final buret reading to
the nearest 0.01 mL in the data table supplied below.
10. When you clean up, please rinse the buret very well, but do not use any soap solutions. Dispose
of all magnesium waste (in your flask) in the special container in the hood. Excess HCl can be
put down the drain.
11. Obtain a second set of data from another group and record in the table below. If you do not like
their data, if you do not agree with their data, then do not use it!
DATA TABLE 2:
concentration
HCl, M
Buret Readings, mL
Initial
Final
total mL HCl
used
g Mg(OH)2
Titration 1
Titration 2
Titration 2
names:
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Observations: Mg(OH)2 solid in water + indicator – note color, and clarity of the solution:
Observations: Mg(OH)2 solid in water + indicator + acid and description of changes that occur during
the chemical reaction: - note – this titration takes ~ 1 hour. Pay attention to the color and clarity of the
solution.
Why does the cloudiness of the solution change as you titrate the acid into the base?
Reflection Questions
PART I: Think back to the first reaction: sodium hydroxide reacting with acetic acid
Was this titration fast? Did you have to wait a long time for the two solutions to react with one
another?
Write the ionic and net-ionic reactions for the molecular equation given below:
HC2H3O2 (aq) + NaOH (aq) → NaC2H3O2 (aq) + H2O (l)
Ionic equation: _______________________________________________________________________
Net-ionic equation: __________________________________________________________________
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PART II: Think back to the second reaction: hydrochloric acid reacting with the solid magnesium
hydroxide
Was this titration fast? Did you have to wait a long time for the two solutions to react with one
another?
Write the ionic and net-ionic reactions for the molecular equation given below:
Mg(OH)2 (s) + 2HCl (aq) → MgCl2 (aq) + 2H2O (l)
Ionic equation: _______________________________________________________________________
Net-ionic equation: __________________________________________________________________
*** - remember – what do we do with SOLIDS in ionic and net-ionic reactions
This second titration took over 1 hour to complete. Explain the following: Why was the solution
containing Mg(OH)2 and water cloudy? What happened to the solution as the HCl was added?
Why did this titration take SO long? Explain what happened chemically in the solution. You must
include the following terminology in your explanation: pH, acid, base, indicator, color changes,
dissociate, dissolve, solid, aqueous, and reaction.
The HCl used in second titration is ~ the same concentration as in your stomach. Magnesium
hydroxide is known as “milk of magnesia” and is present in almost all antacid tablets/solutions in the
amount used in lab. Given the length of time the second titration took, comment on the relationship
between taking an antacid tablet for indigestion, and how long it takes for a person to feel relief.
The length of time a reaction takes is called kinetics. Explain something interesting you have learned
with regard to reaction speed from today’s lab:
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