TITRATION OF HYDROCHLORIC ACID, ACETIC ACID, PHOSPHORIC
ACID AND A COLA BEVERAGE WITH SODIUM HYDROXIDE
I. INTRODUCTION
In this experiment you will use a pH meter to observe the pH changes occurring during the titration of various
acids. You will plot the data so as to obtain TITRATION CURVES of the following form.
II. OPERATION OF THE pH METER
Be extremely careful with the meter. Although it is not a delicate instrument, it is expensive and must not be
sharply jarred or splashed with water. The electrode is an extremely fragile piece of equipment. The glass
bulb at the tip is very thin soft glass.
NOTE: To take a pH measurement, immerse the probe in the solution to be tested, hold the probe against the
inner side of the beaker, swirl the solution for a few seconds to ensure complete mixing, wait 5seconds for
temperature stabilization and record the reading. After this initial reading, leave the pH probe in until the end
of the titration, taking a pH measurement after every addition from the burette. Make sure you swirl the
solution to ensure complete mixing before every reading. It isn’t necessary to wait for temperature
stabilization after the initial reading because any temperature changes are very small.
III. TITRATION OF HYDROCHLORIC ACID WITH SODIUM HYDROXIDE
1. Pipet exactly 25.00 mL of approximately 0.1 M HCl into a clean, not-necessarily dry 250 mL
beaker. Into this solution add about 5 drops of Bromothymol Blue indicator.
2. Rinse and fill a burette with 0.1 M NaOH and record the exact [NaOH] used.
3. Dip the pH probe into a beaker containing distilled water to rinse off the probe.
4. Put the rinsed pH probe into the beaker of HCl. Swirl the solution, wait for temperature stabilization, and
record the initial pH. If the pH reading fluctuates more than about 0.05 pH unit or seems erratic in any
way, or continues to drift up or down, contact your teacher. (The battery may be wearing out.)
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5. Performing the Titration: Read the rest of this section before starting the titration. The general
idea is to add a specified portion of NaOH, swirl the solution, take a reading, add another portion of NaOH
and repeat the process.
i)
The process used for every addition is
• add NaOH from the burette,
• touch the tip of the burette to the inside of the beaker to get the last drop off the tip,
• with the pH probe remaining in the solution, gently swirl the BEAKER with a rotating motion for a few
seconds until the meter reading stabilizes,
• stop the swirling and record the pH reading.
ii)
Use the following APPROXIMATE scheme to add NaOH to the HCl.
NOTE: USE YOUR COMMON SENSE ON THIS – the basic idea is to have relatively small pH
changes throughout the titration so as to end up with a smooth, easy–to–plot curve. You may have to
adjust the scheme to create a small change in pH; if one addition gives a “much bigger–than–usual”
change, reduce the amount used in subsequent additions; if the pH changes are too small, increase
the size of the next addition. Try using the scheme below. It is important to avoid large pH jumps as
much as possible to have enough points on your graph.
Tentative scheme: 1.0 mL increments from 0 mL up to 20 mL
0.5 mL increments from 20 mL up to 23.5 mL
0.1 mL increments from 23.5 mL up to 26.5 mL
0.5 mL increments from 26.5 mL up to 30 mL
5.0 mL increments from 30 mL up to 50 mL
NOTES: (1) Make a table of the TOTAL volume added and the pH at that volume. Do NOT
just give the individual volumes added each time — this latter method of recording the
volumes is highly prone to errors.
(2) If possible, try to plot the data at the same time you record it in your data table. Plotting
the data gives you a much clearer idea of how the pH changes and helps you determine
how much NaOH should be added to get to the next data point.
(3) As you approach the stoichiometric point, the Bromothymol Blue indicator changes from
YELLOW to GREENISH-YELLOW (care! getting close) to GREEN (endpoint!) to
BLUISH-GREEN (just past endpoint) to BLUE (quite basic).
IV. TITRATION OF ACETIC ACID WITH SODIUM HYDROXIDE
1. Pipette 25.00 mL of approximately 0.1 M CH3COOH into a clean, not-necessarily dry 250 mL
beaker. Add about 5 drops of Phenolphthalein indicator.
2. Fill a burette with 0.1 M NaOH and record the exact [NaOH].
3. Put the previously–rinsed pH probe into the beaker of CH3COOH, swirl, and record the initial pH.
4. Titrate the CH3COOH with the NaOH solution from the burette using the following addition scheme
and the same general procedure used for the titration of the HCl. As with the HCl titration, you may have
to adjust the scheme somewhat. The indicator changes from colourless to pink around the end point.
Tentative scheme:
0.2 mL increments from
0.5 mL increments from
1.0 mL increments from
0.5 mL increments from
0.1 mL increments from
0.5 mL increments from
5.0 mL increments from
0 up to 2 mL
2 up to 5 mL
5 up to 20 mL
20 up to 23.5 mL
23.5 up to 26.5 mL
26.5 up to 30 mL
30 up to 50 mL
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V. TITRATION OF PHOSPHORIC ACID WITH SODIUM HYDROXIDE
1. Pipette 10.00 mL of approximately 0.1 M H3PO4 solution into a clean, not–necessarily dry 250 mL
beaker. NOTE: DO NOT USE 25 mL! Do NOT add any indicator.
2. Fill a burette with approximately 0.1 M NaOH solution and record the exact [NaOH].
3. Put the previously-rinsed probe into the H3PO4 solution and keep the beaker tipped when making
the first pH readings. You will have to keep the beaker tipped for the first half of this titration in order to
have a deep enough reservoir of solution to cover the probe tip.
4. Use 0.5 mL additions for the entire titration. There are TWO stoichiometric points; that is, two
“jumps” in pH. Keep adding NaOH until you get to a total of 40.0 mL.
VI. TITRATION OF A COLA BEVERAGE WITH SODIUM HYDROXIDE
Cola beverages contain several acids, including carbonic acid and phosphoric acid. By boiling the cola for a
while, before the cola is given to you, any dissolved carbon dioxide (“carbonic acid”) is flushed out of solution.
The remaining phosphoric acid is titrated with sodium hydroxide in order to determine the [H3PO4]. (Citric acid
is also present but does not affect the titration results to any great extent.)
1. Pipette 25.00 mL of a cola beverage of your choice into a clean 250 mL beaker.
2. Clean a burette, fill it with approximately 0.01 M NaOH solution, and record the exact [NaOH].
IMPORTANT: You must use 0.01 M NaOH, NOT 0.1 M NaOH.
3. Use 0.5 mL additions for the entire titration. Keep adding NaOH until you get to 40 mL.
VII. INITIAL ANALYSIS OF THE DATA
1. Prepare graphs of pH versus volume of NaOH added for each of your four titrations as follows.
a) In order to get the required accuracy, you MUST use millimetre graph paper. Plots submitted
on any other type of paper won't be marked. All graphs must be properly labelled (title, axes, etc.) and
the scale used must be easily readable.
b) Computer-generated graphs are not acceptable because they are far too crude to allow
volumes to be read to the required ± 0.01 mL.
c) Most millimetre graph paper is 18 cm x 25 cm or 19 cm x 26 cm. Turn the paper sideways and
put the pH on the vertical axis (18 cm length) and the volume of NaOH on the horizontal axis (25cm
length).
d) When all the data points are plotted, draw a SMOOTH CURVE through the data points using a
sharp HB pencil. A graph is NOT ACCEPTABLE unless the resulting curves are EXTREMELY
SMOOTH AND CAREFULLY DONE WITH THIN LINES. Drafting instruments called “French curves”
or “flexicurves” give very good results. Again, computer–drawn graphs ARE NOT ACCEPTABLE.
2. If a titration curve is symmetric about the "inflection point", the following method is an accurate way
to find the exact stoichiometric point pH and volume. Trace the upper half of the titration curve onto
tracing paper and mark an approximate stoichiometric point pH on BOTH the graph (a1) and the tracing
paper (a2), as shown below.
a1
a2
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Now, rotate the tracing paper by 180o and lay it over the bottom half of the titration curve.
The EXACT stoichiometric point pH will be halfway between the two “approximate” values, a1 and a2.
After finding the exact stoichiometric point pH you can quickly read the exact stoichiometric point volume
off the graph.
If the titration curve is NOT symmetric about the inflection point, the following procedure should be
used.
Draw line AB perpendicular to the upper extrapolated line, and draw line CD perpendicular to the lower
extrapolated line. E is the midpoint of line AB; F is the midpoint of line CD. Draw a line from E to F. The
point at which line EF intersects the titration curve will be the stoichiometric point. The pH and volume at
the stoichiometric point can then be read off the graph in the usual way. (Incidentally, this method will also
work with symmetric curves.)
3. Find and record on each graph the exact pH and volume for each stoichiometric point using either the
symmetric or nonsymmetric method discussed in #2, above.
a) Mark every stoichiometric point on each graph with a “cross”, as shown below.
There should be two separate stoichiometric points on the H3PO4 graph (the third stoichiometric
point, corresponding to the removal of the last proton, is not observed with the low concentrations of
NaOH you used). The graph for the titration of the cola beverage may show one or two stoichiometric
points.
b) Write the pH and volume at the stoichiometric points on each graph.
c) Determine the pH and volume at the half–volume points of acetic acid and phosphoric acid,
as follows. After determining the volume of NaOH used at the stoichiometric point, divide this volume
by two and mark this “half-volume” point on the curve.
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The acetic acid graph will resemble:
The phosphoric acid graph will resemble:
V1 & V2 = stoichiometric point volumes
pH
H1 & H2 = half volumes
H1
V1
H2
V2
IMPORTANT: You must mark the volume and pH coordinates of the half-volume points and the
stoichiometric points on each graph. NOTE: H2 lies exactly halfway between V1 and V2 .
VIII. PREPARATION OF A DATA TABLE
Complete the section labelled "DATA OBTAINED FROM GRAPHS" at the end of this handout, which
requires the following information.
• the acid / solution being titrated
• the concentration of the NaOH used (0.1000 M; except for the Colas, where you used 0.0100 M)
• the volume of acid used
• the stoichiometric volume of base used (found from the graph). If more than one stoichiometric point is
observed, a separate volume must be recorded for each point.
• the pH and volume of the “half-volume” points (only for CH3COOH and H3PO4)
• the initial pH before any NaOH is added (only for HCl, CH3COOH and H3PO4)
IX. CALCULATIONS
1. Calculating the Concentrations of the Acids by Using Stoichiometric Point Data
a) Calculating the Concentration of the HCl
The reaction between HCl with NaOH is given by NaOH + HCl  NaCl + H2O .
Calculate the [HCl] using the data recorded in “DATA OBTAINED FROM GRAPHS”.
b) Calculating the Concentration of the CH3COOH
The reaction between CH3COOH and NaOH is given by
NaOH + CH3COOH  NaCH3COO + H2O .
Calculate the [CH3COOH] using the data recorded in “DATA OBTAINED FROM GRAPHS”.
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c) Calculating the Concentration of the H3PO4
There are two stoichiometric points visible on the graph.
The first stoichiometric point corresponds to the partial neutralization
NaOH + H3PO4  NaH2PO4 + H2O .
Calculate the [H3PO4] using the data recorded in “DATA OBTAINED FROM GRAPHS”.
The second stoichiometric point corresponds to the removal of a second proton:
2 NaOH + H3PO4  Na2HPO4 + 2 H2O .
Calculate the [H3PO4] using the data recorded in “DATA OBTAINED FROM GRAPHS” for the
second stoichiometric point. Note that this calculation requires the total volume of NaOH used to
get to the 2nd stoichiometric point, not just the extra volume used to go from the 1st stoichiometric
point to the 2nd stoichiometric point. (The second stoichiometric point volume IDEALLY should
be twice the first stoichiometric point volume, but actually may be slightly different.)
Average the concentration values if the [H3PO4] found for the first and second stoichiometric points do
not agree exactly (they should be close).
d) Calculating the Concentration of H3PO4 in a Cola Beverage
The calculation of the [H3PO4] in the cola beverage is based on the assumption that the first inflection
point on your graph corresponds to
H3PO4 + NaOH  NaH2PO4 + H2O .
Calculate the [H3PO4] in your cola beverage using the data recorded in “DATA OBTAINED FROM
GRAPHS”. YOU MUST ALSO NAME the brand of cola you titrated.
2. Calculating the Ka Value and the Concentration of CH3COOH from the Initial pH and the pH
Halfway through the Titration
a) At the point in the titration equal to the half volume,
pKa = pH1/2
Calculate the Ka value for acetic acid by converting the pH at the half-volume point on your acetic
acid graph to [H3O+]. This value will equal Ka for acetic acid.
b) At the start of the titration, the solution of CH3COOH has a pH value, pHINIT., which allows you to
calculate [H3O+]INIT.
[H O ]2INIT
The concentration of CH3COOH is found from
[CH3COOH] = 3
.
Ka
Calculate the [CH3COOH] using the data recorded in “DATA OBTAINED FROM GRAPHS” , the
value of Ka calculated in calculation 2 (a) and the above equation. (This value for the concentration of
CH3COOH may not agree exactly with the value found in calculation 1 (b).)
–
3. Calculating the Ka Values for H3PO4 and H2PO4
a) Calculating the Value of Ka (H3PO4) and [H3PO4]
At the point in the titration equal to the first half–volume point, half of the NaOH required to get to the
first stoichiometric point is added. At this point, half of the original H3PO4 remains unreacted and half
–
has been reacted to form H2PO4 . Hence, at this first “half volume”
–
[H3PO4] = [ H2PO4 ]
and the Ka expression
simplifies to
Ka (H3PO4) =
Ka (H3PO4) = [H3O+] .
[H 3O ][H2PO4– ]
= [H3O+]
[H3PO4 ]
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Calculate the Ka value for H3PO4 from the “DATA OBTAINED FROM GRAPHS”.
At the start of the titration, the solution is mainly H3PO4. The small amount of H3PO4 that
–
dissociates produces equal initial amounts of H3O+ and H2PO4 ; that is
–
[H3O+]INIT = [ H2PO4 ]
The Ka expression for H3PO4 is Ka =
[H 3O ]2INIT
[H3PO4 ]
Rearranging this expression gives
[H3PO4] =
[H 3O ]2INIT
Ka
Calculate the [H3PO4] from the “DATA OBTAINED FROM GRAPHS”. This value for the
concentration of H3PO4 may not agree exactly with the value found in part 1 (c).
–
b) Calculating the Value of Ka ( H2PO4 ) from the Second Half-volume Point
Assuming the first proton is removed, the removal of the second proton proceeds according to
H2PO4– + OH–  HPO2–
4 + H2O .
At the second half–volume point, half of the NaOH required to get from the first to the second
–
stoichiometric point is added. At this point, half of the H2PO4 present at the first stoichiometric point
2–
remains unreacted and half has reacted to form HPO4 . Hence, at this second half volume
–
2–
[ H2PO4 ] = [ HPO4 ]
and the Ka expression
simplifies to
Ka =
[H 3O ][HPO2–
4 ]
= [H3O+]
[H2PO–4 ]
Ka = [H3O+]
–
Calculate the Ka value for H2PO4 from the “DATA OBTAINED FROM GRAPHS”.
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DATA OBTAINED FROM GRAPHS
HCl:
[NaOH] used ........................................
= _______________
volume HCl used ....................................
= _______________
stoichiometric volume of NaOH ..................... = _______________
CH3COOH:
pH at stoichiometric point ...........................
= _______________
initial pH of HCl ......................................
= _______________
[NaOH] used ........................................
= _______________
volume CH3COOH used ...........................
= _______________
stoichiometric volume of NaOH ....................
= _______________
pH at stoichiometric point ......................
= _______________
volume at “half–volume” point .... ............ = _______________
pH at “half–volume” point ......................
H3PO4:
= _______________
initial pH of CH3COOH ..............................
= _______________
[NaOH] used ........................................
= _______________
volume H3PO4 used ................................
= _______________
first stoichiometric volume of NaOH ................ = _______________
pH at first stoichiometric point .................. = _______________
volume at first “half–volume” point ............. = _______________
pH at first “half–volume” point .................. = _______________
second stoichiometric volume of NaOH ............= _______________
pH at second stoichiometric point .............. = _______________
volume at second “half–volume” point ........ = _______________
pH at second “half–volume” point ............. = _______________
Cola:
initial pH of H3PO4 ...................................
= _______________
Brand of Cola ........................................
= _______________
[NaOH] used ........................................
= _______________
volume Cola used ...................................
= _______________
stoichiometric volume of NaOH ..................... = _______________
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