Chem 1312
Handout Experiment
ONE
GETTING THE END POINT TO APPROXIMATE
THE EQUIVALENCE POINT
Laboratory Time
Required
Special Equipment and
Supplies
Two hours
Objective
Safety
First Aid
Balance
Burets
Buret clamp
pH Electrode
pH Meter
Potassium hydrogen phthalate, KHP (s)
0.1 M Sodium hydroxide, NaOH (aq)
Vinegar
Phenolphthalein
Universal Indicator
pH = 7 Buffer
pH = 4 Buffer
10-mL Pipet
Pipet bulb
Stirring bar
Magnetic stirrer
In performing this experiment, the student will consider the
factors that affect an indicator’s suitability for use in an
acid/base titration.
Bases, such as sodium hydroxide, can cause skin burns and
are especially hazardous to the eyes. Although vinegar is a
dilute solution of a weak acid (acetic acid), it is,
nevertheless, advisable to avoid splashing it in the eyes..
Following skin contact with either sodium hydroxide,
wash the area thoroughly with water. Should sodium
hydroxide (or even vinegar) get in the eyes, rinse them
with thoroughly with water. At least 20 minutes of
flushing with water is recommended. Then seek medical
attention.
There are a number of experiments in your lab manual that involve acid/base indicators.
In Experiment 8, phenolphthalein is used to mark the end point in the titration of vinegar
with sodium hydroxide. Methyl orange serves the same purpose in the titration of
hydrochloric acid with ammonia (Experiment 30). In Experiment 13, bromothymol blue
changes color when carbon dioxide dissolves in water, creating an acidic solution. In this
experiment we explore the question of why different indicators change color at different
pH values and illustrate that the end point of a titration (signaled by an indicator’s color
change) does not necessarily correspond to the equivalence point in the titration.
PRINCIPLES
The equivalence point in a titration is the point at which the acid and base have been
mixed in “stoichiometric proportion”, meaning that neither acid nor base is in excess. For
example, if 50.00 mL of 0.1000 M acetic acid is titrated with 0.1000 M sodium
hydroxide, the equivalence point will be reached when 50.00 mL of base have been
added to the acid. At that point in the titration, the acid and base have exactly
“neutralized” each other. The question is “How can one determine that this point has
been reached?”
In the early portion of the titration, the chemical amount of base that has been
added to the titration mixture is less than the chemical amount of acid present and the pH
of the titration is less than 7 (at 25°C). At the equivalence point, the titration mixture is
essentially a 0.05000 M solution of sodium acetate and should have a pH of 8.7 (at
25°C), because of the hydrolysis of the acetate ion (see Experiment 22). Further addition
of base raises the pH well above 7 (25°C) because there is no longer any acid present to
react with the extra sodium hydroxide. These changes in pH take place without any
apparent changes in the appearance of the reaction mixture. The equivalence point is
revealed either by monitoring the changing pH of the solution (by the use of a pH meter)
or by finding an indicator that will change color in the vicinity of the equivalence point.
However, not all indicators will change color at the desired. point in the titration.
Indicators
As discussed previously (see Experiment 21), indicators are weak acids that have
different colors at low pH values (when the indicator is predominantly in its HIn form)
and at high pH values (when the indicator is predominantly in its In– form). A good rule
of thumb is that the indicator will display its low pH color when the ratio of [HIn] to [In–]
has a value of 10 or more; conversely, the indicator will display its high pH color when
the ratio of [HIn] to [In–] has a value of 0.1 or less. Consider an indicator which is blue in
the HIn form and is red in the In– form. If the indicator has a Ka value of 1.0 × 10–4, it will
appear blue at pH’s below 3 and red at pH’s above 5 (see EquationsONE.1 through
ONE.3).
Ka
[H3O+] [In–]
[HIn]
=
(ONE.1)
1.0 × 10–4
=
[H3O+](1); [H3O+] = 1.0 × 10–3 when [HIn] = 10[In–]
(10)
(ONE.2)
1.0 × 10–4
=
[H3O+](10); [H3O+] = 1.0 × 10–5 when [In–] = 10[HIn]
(1)
(ONE.3)
Calculations similar to those shown in Equations ONE.1 through ONE.3 reveal
that, if the indicator had a Ka value of 1.0 × 10–8, it would appear blue at pH’s below 7
and red at pH’s above 9. Thus, an indicator with Ka = 1.0 × 10–4 would change color well
before the equivalence point was reached in the titration of acetic acid by NaOH while an
indicator with Ka = 1.0 × 10–8 would change color just after the equivalence point had
been reached in the same titration.
Universal Indicator
Universal Indicator is actually a mixture of various indicators, chosen so that the mixture
will undergo several color changes as the pH of the solution being titrated varies from a
value of 4 to a value of 10. In this experiment, you will perform a pH titration of acetic
acid by sodium hydroxide in the presence of Universal Indicator. This will permit you to
determine whether a given color change occurs near the equivalence point or not.
PROCEDURE
Procedure in a Nutshell
Standardize 0.1 M NaOH via titration with KHP, using phenolphthalein as an
indicator. Dilute 10.00 mL of vinegar to 100.00 mL. Standardize the resulting 0.1 M
HC2H3O2 via titration with the NaOH, again using phenolphthalein as an indicator.
Add 3 drops of Universal Indicator to 25 mL of the dilute acetic acid and titrate it
with NaOH, monitoring the course of the titration with a pH meter.
Standardizations
Clean a buret and prepare it for use in the standardization of sodium hydroxide according
to the directions provided in the Introduction. Accurately weigh, to the nearest 0.1 mg,
0.4 g of KHP. Place the KHP in an Erlenmyer flask, dissolve it in 25 mL of distilled
water, and add 2 to 3 drops of phenolphthalein. Read and record the initial volume of
NaOH in the buret to the nearest 0.01 mL. Titrate the KHP solution until a faint pink
color, which does not disappear when the solution is swirled, is obtained. Read and
record the final volume of NaOH in the buret to the nearest 0.01 mL. Refill the buret with
sodium hydroxide.
Rinse a pipet with two small portions of vinegar. Then, use the pipet to deliver a
10-mL sample of vinegar to a clean (but not necessarily dry) 100-mL volumetric flask.
Add 50 mL of distilled water to the vinegar and swirl the flask carefully to mix its
contents. Dilute the solution to the 100-mL mark with distilled water; invert the capped
flask several times to promote further mixing of the diluted vinegar.
Clean a second buret and prepare it for use in the standardization of the dilute
acetic acid according to the directions provided in the Introduction. Deliver a 25 mL
HC2H3O2 sample to an Erlenmyer flask (read and record the initial and final buret
readings to the nearest 0.01 mL). Add 2 to 3 drops of phenolphthalein to the acid. Read
and record the initial volume of NaOH in the buret to the nearest 0.01 mL. Titrate the
HC2H3O2 solution to the phenolphthalein end point. Read and record the final volume of
NaOH in the buret to the nearest 0.01 mL.
Calculate the molarities of the NaOH and the HC2H3O2 solutions. Determine the
volume of NaOH that will have to be added to 25 mL of HC2H3O2 to reach the
equivalence point.
pH Titration
Prepare the pH meter for titration (set the Function Knob to “Standby”, the Slope Knob
to “100%”, and attach the electrode). Immerse the electrode in the pH = 7 buffer solution.
Set the pH meter’s “Temperature” Knob to the room temperature (probably 20°C to
24°C). Move the Function Knob to “pH” and adjust the Calibration Knob until the pH
meter reading matches the pH of the buffer exactly. Return the Function Knob to
“Standby”.
Rinse the electrode with distilled water and shake it to dry the bulb. Immerse the
electrode in the pH = 4 buffer solution. Move the Function Knob to “pH” and adjust the
Slope Knob until the pH meter reading matches the pH of the buffer exactly. Return the
Function Knob to “Standby”.
Refill the base buret with sodium hydroxide and the acid buret with acetic acid.
Deliver a 25 mL HC2H3O2 sample to an Erlenmyer flask (read and record the initial and
final buret readings to the nearest 0.01 mL). Add 2 to 3 drops of Universal Indicator to
the acid. Read and record the initial volume of NaOH in the buret to the nearest 0.01 mL.
Rinse the electrode with distilled water and shake it to dry the bulb. Immerse the
electrode in the HC2H3O2 solution. Put the stirring bar in the solution. Adjust the
magnetic stirrer so that the movement of the stirring bar during the course of the titration
does not endanger the pH electrode. Move the pH meter’s Function Knob to “pH”.
Record the initial pH of the HC2H3O2.
Begin the titration. Start by adding 1-2 mL increments of NaOH to the diluted
vinegar. Record the pH meter reading after each addition of NaOH. Note the pH of the
titration mixture at which each color change occurs.
As the volume of base added approaches the calculated equivalence point,
decrease the size of the NaOH increment. (Add the NaOH dropwise for 2 mL before and
2 mL after the anticipated equivalence point). Continue adding NaOH and recording the
pH (and any color change) until you have gone 10 mL beyond the equivalence point.
Prepare a plot of pH (y-axis) versus volume of NaOH added (x-axis). The
equivalence point is marked by a large increase in the pH upon the addition of a small
volume of base.
Calculations
Titrations are usually performed to find the molarity of an acid or base. Assume that the
molarity of your HC2H3O2 solution is correct and calculate the molarity of NaOH (using
M aVa = M bVb ) using the volume of base added that corresponds to (1) each of the
observed color changes and (2) the equivalence point as determined on the pH plot.
Calculate the % difference between the molarity based on the pH titration and the
molarity based on each color change. State the values of Ka required for an indicator to be
suitable for use in a titration of HC2H3O2 by NaOH.
Disposal of Reagents
Excess KHP should be placed in the containers used for solid waste. Solutions should be
neutralized and diluted. They may then be flushed down the drain.
Date_____Name_______________________________Section_____Desk Number_____
Summary Report on Handout Experiment ONE (take your data in your notebook
as usual; Summary Report Sheets are shown only as a guide for the preparation of
the data table in your notebook)
Standardization of NaOH Solution
Mass of KHP and container
_____________
Mass of container
_____________
Mass of KHP
_____________
Final buret reading, NaOH
_____________
Initial buret reading, NaOH
_____________
Volume used, NaOH
_____________
Molarity of NaOH solution
_____________
Standardization of HC2H3O2 Solution
Final buret reading, HC2H3O2
_____________
Initial buret reading, HC2H3O2
_____________
Volume used, HC2H3O2
_____________
Final buret reading, NaOH
_____________
Initial buret reading, NaOH
_____________
Volume used, NaOH
_____________
Molarity of HC2H3O2 solution
____________
pH Titration
Buret reading, mL
Volume of base
pH
color
added, mL
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Buret reading, mL
Volume of base
pH
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Buret reading, mL
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pH
color
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Date_____Name_______________________________Section_____Desk Number_____
Pre-Laboratory Exercises for Handout Experiment ONE
These exercises are to be completed after you have read the experiment but before you come to lab to
perform it.
Consider the titration of 40.00 mL of 0.2000 M acetic acid with 0.2000 M sodium
hydroxide. Calculate the pH of the titration mixture at the following points in the
titration.
1) start (no NaOH added)
2) after the addition of 10.00 mL of NaOH
3) after the addition of 20.00 mL of NaOH
4) after the addition of 30.00 mL of NaOH
5) after the addition of 39.95 mL of NaOH
6) after the addition of 40.00 mL of NaOH
7) after the addition of 40.05 mL of NaOH
8) after the addition of 42.00 mL of NaOH