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1
STANDARIZATION OF A BASE
AND TITRATION OF A VINEGAR SOLUTION
ADDITIONAL READING
The concepts in this experiment are also discussed in sections
3.6 AND 17.3 of Chemistry and Chemical Reactivity by
Kotz, Treichel, Townsend and Treichel, and in sections 4.3b,
17.3a, and 17.3b of Mindtap General Chemistry by Vining,
Young, Day, and Botch
ABSTRACT
This experiment is divided into two parts. Each student is
expected to perform the experiment individually.
In Part A, you will prepare a NaOH titrant solution, then
standardize it (determine its exact concentration) using the acid
primary standard, potassium hydrogen phthalate, KHC8H4O4,
frequently abbreviated as KHP. Note KHP is not a chemical
formula.
In Part B you will use your standardized NaOH solution to
determine the molar concentration of vinegar (an acetic acid,
CH3COOH, solution), and convert this concentration unit to a
mass percent concentration unit, and finally compare your

2. measured mass percent concentration to the value reported on
the bottle.
BACKGROUND
TITRATIONS
One of the most useful strategies in analytical chemistry is to
use a known reagent (known composition or concentration)
as a standard to analyze an unknown substance. A titration is
an analytical procedure in which a solution of known
concentration, the standard solution, is slowly reacted with a
solution of unknown concentration. The concentration of
the unknown solution can be easily calculated. Titration is
often used to measure the concentration of an acid or base,
but it can also be used for any chemical reaction if the
stoichiometry is known.
EXPERIMENTS 6 AND 7 ARE BOTH ACID BASE
TITRATION EXPERIMENTS, QUITE SIMILAR TO EACH
OTHER.
THE REASONS FOR DOING TWO TITRATION
EXPERIMENTS
A. TO GIVE STUDENTS PLENTY OF OPPORTUNITY BOTH
TO PERFECT THEIR TITRATION TECHNIQUE AND
TO LEARN TO DO THE CALCULATIONS;
B. TITRATION IS THE MOST IMPORTANT TECHNIQUE
LEARNED IN CHEM 1033 LAB.
YOU WILL DO A PRACTICAL LAB EXAM AT THE END OF
THE SEMESTER; IT WILL BE A VERY SIMILAR
TITRATION.
IT IS IMPORTANT TO REALIZE THAT TITRATION IS AN
ACQUIRED SKILL, REQUIRING PRACTICE. MOST
STUDENTS ARE NOT PROFICIENT AT FIRST, BUT IF YOU

3. WANT TO BECOME EXPERT AT IT, YOU WILL GET
THERE WITH PRACTICE.
It is critical that there be an observable change that signals that
the titration is complete. This is called the endpoint,
since it signals the end of the titration, when the equivalents of
titrant added just equal the equivalents of the analyte
unknown. When performing an acid-base titration, we
commonly use an acid-base indicator that has one color before
the
endpoint but changes sharply to a different color at the pH of
the endpoint.
Titrations are carried out using a specialized piece of glassware
called a buret, which is long tube with a dispensing valve.
The buret scale has graduated marks in units of 0.01 mL or 0.02
mL. You can apply the techniques used for reading the
graduated cylinder to reading a buret. The primary difference
between a graduated cylinder and a buret is that the zero
mark is at the top of the buret and at the bottom for a graduated
cylinder.
The titrant solution is delivered through the stopcock (the
dispensing valve) and bottom tip of the buret into a receiving
flask. Therefore, the volume in the buret will decrease and the
meniscus will drop as the solution is delivered. It is a
common mistake to read the buret volume from the bottom-up.
Simply read and record the number from the top down (to
the right number of significant figures, and with units);. You
don’t need to add or
subtract your reading from 50.
In this experiment you will place the NaOH solution into a
clean buret. The solution
that is placed into the buret is known as the titrant. When
readying the buret, over-fill

4. the buret above the 0.00 mL mark. The stopcock is then opened
to allow the solution
15
16
17
Always measure
the bottom of the
meniscus at eye
level.
←
2
to flow out of the tip until the solution level falls just below
0.00 mL. This process ensures that the tip of the buret is filled
with the solution and not with an air bubble (more on this later).
It is not critical that the initial volume of the solution be
exactly 0.00 mL. However, it is important that the value be
between 0.00 and 1.00 mL, and that the exact value be
recorded.
Once the buret is loaded with the titrant, we can then prepare
the analyte solution in the Erlenmeyer flask. It is critical
that an indicator be added to the analyte solution to signal the
end point; otherwise, you will not know when to stop the
titration and you may “over shoot” the end point.
Before the titration is started, the initial volume of titrant in the

5. buret is read to the nearest 0.01 mL, i.e., to two decimal
places; you estimate one decimal beyond the graduations. The
titrant is added slowly to the analyte solution. As the
addition continues, a fleeting color change may be noted (as a
result of the indicator), especially when the titrant first
contacts the analyte solution. It is important that the analyte
solution be swirled constantly to ensure thorough mixing.
Swirl carefully to avoid splashing of the analyte solution. The
endpoint is not reached until the entire solution has a
permanent color change (does not disappear with swirling).
With practice, it is possible to determine the endpoint within a
single drop of titrant solution. Once the color change is
permanent, the first titration trial is complete, and the final
volume
can be read from the buret. The difference between the final
and initial buret reading is the volume of titrant added to the
flask.
The volume and concentration of the standard are used to
calculate the moles of titrant added. The moles of titrant can
then be stoichiometrically compared to the moles of analyte
present, allowing us to calculate the concentration of the
analyte solution.
Since titration is an analytical technique, multiple trials are run
to ensure consistency of the results. Titrations should be
run with a minimum of two trials; however, the more trials the
better your results are likely to be. If, after performing a
titration, you find that one set of results does not agree with the
others, it should be excluded. Ideally, an additional trial
should be run to confirm the results. Titration is an accurate
and precise technique; if done correctly, it is possible to
perform experiments with less than 1% error. It will take some
practice for you to reach this level; nevertheless, you
should strive to keep your errors low. Part of your grade will
be based on the accuracy of your experimental results.

6. Consider the following example. A student performed an acid-
base titration to determine the concentration of a potassium
hydroxide solution, KOH(aq). The standard solution was 0.275
M nitric acid, HNO3. The student filled the buret above the
0.00 mark with the nitric acid solution, then drained it to 0.18
mL, ensuring the tip had no air bubbles. The student used a
volumetric pipet to dispense exactly 25.00 mL of the unknown
potassium hydroxide solution into an Erlenmeyer flask, and
added a few drops of the acid base indicator. The student
slowly added the nitric acid solution while swirling the
Erlenmeyer flask, until the color change persisted throughout
the solution, and recorded the final volume on the buret as
43.08 mL.
Volume of nitric acid used = Vfinal – Vinitial = 43.08 mL –
0.18 mL = 42.90 mL HNO3
To calculate the moles of nitric acid, use the relationship:
moles = Molarity x volume in Liters
3
2-
3
3
3 HNO mol 10 x 1.18
mL 1000
L 1 x HNO mL 42.90 x
HNO L 1
HNO mol 0.275
=
From the balanced chemical equation for the reaction between

7. potassium hydroxide and nitric acid (equation 1), we know
the stoichiometric relationship between potassium hydroxide
and nitric acid is 1-to-1.
HNO3(aq) + KOH(aq) → KNO3(aq) + H2O(l) (1)
From the stoichiometry we can determine the moles of
potassium hydroxide present.
KOH mol 10 x 1.18
HNO mol 1
KOH mol 1 x HNO mol 10 x 1.18 2-
3
3
2- =
The moles of potassium hydroxide and the volume of the
solution used (25.00 mL = 2.500 x 10-2 L) can then be used to
calculate the concentration of the potassium hydroxide solution.
M 0.472 mol/L 0.472
L 10 x 2.500
mol 10 x 1.18
KOH of L
KOH mol solution KOH ofMolarity 2-
-2
====
Each step (above) is shown individually for clarity. However, it
is possible to combine all steps together using a factor-
label sequence:

8. 3
KOH of L
KOH mol solution KOH ofMolarity ==
mol/L 0.472
KOH L 1
KOH mL 1000 x
KOH mL 00 25.
1 x
HNO mL 1000
HNO L 1 x HNO mL 42.90 x
HNO L 1
HNO mol 0.275 x
HNO mol 1
KOH mol 1
3
3
3
3
3

9. 3
=
Note that the key to a titration is the stoichiometric comparison
of the moles of one reactant (the acid) to the moles of the
other reactant (the base); without this comparison a titration is
impossible.
A simpler solution to the above titration problem consists of
using the molarity equation. The relationship is as follows.
B
BB
A
AA
n
VM
n
VM
= (2)
The quantities are as follows: MA = molarity of the Acid; VA =
volume of the Acid
MB = molarity of the Base; VB = volume of the Base
nA = coefficient for the Acid in the balanced chemical
equation
nB = coefficient for the Base in the balanced chemical
equation

10. In the acid-base reaction: HNO3(aq) + KOH(aq) →
KNO3(aq) + H2O(l), nA = 1 and nB = 1; since the mole ratio
of
acid:base is 1:1.
Using equation 2, we can compute the molarity of the KOH,
MB:
1
mL) (25.00 x M
1
mL) (42.90 x M) 275.0( B=
MB = molarity of potassium hydroxide = 0.472 M
Note that it is not necessary to convert the volumes to liters.
Leave both the volumes in mL; these units will cancel.
PRIMARY STANDARDS
As mentioned above, the standard solution must have a known
concentration. Some compounds can be used to prepare
solutions in which the concentration is accurately known.
These compounds are called primary standards. It is
important that a primary standard has the following
characteristics:
• Purity – We need to be able to accurately determine the
number of moles of primary standard present. Impurities of
any kind make this difficult, if not impossible.
• Stability – Some compounds undergo chemical reactions while
they are in the bottle or when exposed to air. Once

11. this happens, the sample is no longer pure.
• High molar mass – Compounds that have a high molar mass
are easier to weigh, thus reducing errors introduced in
mass measurement. A small error in mass results in an even
smaller error in the number of moles if the molar mass
is large.
Fortunately, there are a number of primary standards that are
not expensive or cumbersome to use. A small quantity of a
primary standard is often used to standardize a large volume of
solution. Once this solution is standardized, we call it a
secondary standard, or simply, standard. This standard can be
used in a titration (since its concentration is known) to
determine the concentration of other solutions. When you use a
secondary standard, you are performing two sets of
titrations. The first set of titrations is used to standardize the
secondary standard solution; the second set is to analyze
the unknown sample.
In this experiment, you will use the primary standard potassium
hydrogen phthalate (KHC8H4O4, abbreviated KHP) to
standardize a sodium hydroxide solution. The stoichiometry for
the reaction is 1-to-1.
KHC8H4O4(aq) + NaOH(aq) → KNaC8H4O4(aq) + H2O(l)
(3)
HC8H4O4–(aq) + OH–(aq) → C8H4O42–(aq) + H2O(l) (4)
Equation 4 is the net ionic equation. As you can see, one mole
of KHP reacts with one mole of sodium hydroxide, and the
potassium and sodium ions are spectator ions. Note that only
one of the hydrogen atoms in KHP (shown in bold) is acidic
and reacts with hydroxide ions. KHP is a monoprotic acid,
meaning the compound contains only one acidic hydrogen
atom.

12. The reason we need to standardize the sodium hydroxide
solution is that NaOH pellets are hygroscopic, which means
they readily absorb water from the air, resulting in an impure
sample as soon as the bottle is opened. Therefore, when
4
you weigh a sample of sodium hydroxide, you are actually
determining the mass of a mixture, NaOH and H2O. It is very
important that students immediately replace the cap on the
sodium hydroxide bottle to minimize the absorption
of water. Sodium hydroxide exposed to air will eventually
absorb enough water to form a solution.
To determine the concentration of the NaOH solution, we need
to know the moles of KHP used and the volume of NaOH
used to neutralize all the KHP. If we know the mass and the
molar mass of KHP, we can calculate the moles of KHP.
The stoichiometry is 1-to-1 so this is also equal to moles of
NaOH. Molarity (concentration) is moles of solute (NaOH) per
liter of solution (the volume of NaOH used). Suppose 1.2683 g
of KHP (MM = 204.22 g/mol) was titrated with 14.67 mL of
NaOH(aq). The molarity of the NaOH(aq) solution can be
determined as follows:
moles of KHP =
KHP g 204.22
KHP mol 1 x KHP g 1.2683 = 6.2105 x 10–3 mol KHP
moles KHP = moles NaOH = 6.2105 x 10–3 mol NaOH
M 0.4233 mol/L 0.4233

13. L 10 x 1.467
mol 10 x 6.2105
NaOH of L
NaOH mol solution NaOH ofMolarity 2-
-3
====
TITRATION OF VINEGAR
Once we have standardized our NaOH(aq) solution, we can use
it to determine the concentration of an unknown acid. In
this experiment you will titrate a sample of vinegar (the
analyte) with the NaOH (the titrant). Vinegar is a dilute
solution of
acetic acid, CH3COOH; a weak monoprotic acid and therefore a
weak electrolyte containing a small concentration of ions.
The reaction with NaOH is:
CH3COOH(aq) + NaOH(aq) → CH3COONa + H2O(l) (5)
CH3COOH(aq) + OH–(aq) → CH3COO–(aq) + H2O(l) (6)
You will already have determined the concentration of the
NaOH(aq) solution. You will know the volume of vinegar used
and the volume of NaOH(aq). Using the molarity equation
(equation 2), it is then quite straightforward to compute the
molarity of the vinegar solution.
MASS PERCENT CALCULATIONS
Some commercially available solutions, such as vinegar, have a
concentration unit of mass percent (m/m %) on the label.
For a solution, mass percent = 100x
solution of mass
solute of mass

14. We can convert between concentration units of mass percent
and molarity (M). For example, a solution that is labeled as
37% HCl, means there are 37 g of HCl for every 100 g of
solution. We need to know the density of the solution, which in
this case is 1.19 g/mL, and the molar mass of HCl, which is
36.46 g/mol.
HCl g 36.46
HCl mol 1 x
solution L 1
solution mL 1000 x
solution mL 1
solution g 1.19 x
solution g 100
HCl g 37
solution of L
HCl mol solution HCl ofMolarity ==
= 12 mol/L = 12 M
For the reverse calculation (molarity to mass percent), the
density, molar mass, and molarity are still needed for this
calculation:
100 x
solution g
HCl g solution HCl of percent Mass =
100 x
solution g 1.19

15. solution mL 1 x
solution mL 1000
solution L 1 x
HCl mol 1
HCl g 36.46 x
solution L 1
HCl mol 12 =
= 37 %
In this experiment you will use a density of the vinegar solution
of 1.02 g/mL, along with your experimentally-determined
value of the molarity, to calculate the mass percent, and
compare your value to that on the bottle.
5
AIR BUBBLES
An air bubble in the tip of your buret, or anywhere in the buret,
can be a serious problem. Imagine that an air bubble in
the buret occupies a volume of 1.00 mL (this is a large bubble,
but the exaggerated volume helps illustrate the problem).
Assume that a student makes an initial volume reading at 0.22
mL and reaches an end point at 28.72 mL. The student’s
calculated volume is 28.50 mL.
DV = Vfinal – Vinitial = 28.72 mL – 0.22 mL = 28.50 mL

16. However, 1.00 mL of the volume dispensed was due to an air
bubble and not the titrant solution. Therefore, only 27.50 mL
of the titrant was added to the analyte. The volume has an error
of >3%. If the air bubble had been removed before the
titration, the error could have easily been avoided. Since
titration requires precision, assuming it is carried out correctly,
every precaution should be taken to reduce errors. With
practice, it is possible to get results with <1% error.
CHECKING YOUR WORK
If you are performing the standardization of base accurately
(part A), the molarities that you calculate from multiple trials
(at least two) should agree. Similarly, the molarities of vinegar
that you measure in multiple trials in part B of the
experiment should agree closely. It is your responsibility to
make certain that you have at least two good trials (good
means you are confident you made no errors, and the results of
the two trials agree to within 1%) for each part of the
experiment.
Example: calculation whether two standardization trials agree
within 1%
[NaOH] trial
0.1333 M 1
0.1401 M 2
% agreement = 100 x |M1 – M2| / (average of M1, M2) = 100 x
|.1333 – .1401| / ((.1333 + .1401)/2) = 5.0%
thus, these two values agree within 5% (remember you should
aim for two good trials that agree 1% or batter)

17. SAFETY/HYGIENE/WASTE DISPOSAL
1. Wear your goggles!
2. Caution: many chemicals, in particular sodium hydroxide
pellets and acids, are corrosive. You should wear gloves
and be sure your skin is not exposed. Discard the gloves after
use, and wash your hands. Acid or base exposure
may result in an itchy sensation. If you have any sensation of
itching, burning, or tingling, start to thoroughly flush the
area with water and inform your lab instructor. Continue
flushing until several minutes after the sensation has
subsided. Inform your lab instructor.
3. Never raise containers of solution, especially corrosive
solutions, above shoulder level.
4. Never weigh chemicals directly on a balance pan. Use
weighing paper, glassware, or other secondary container.
5. Waste solutions from this experiment should be disposed of
in the waste container.
6
PROCEDURE
THIS EXPERIMENT WILL BE DONE INDIVIDUALLY
Part A: Standardization of NaOH

18. 1. Students are to work individually in this experiment.
2. Prepare 250 mL of 0.15 M NaOH. Weight the NaOH on a top
loading balance (confirm the mass needed with your
instructor before you begin), and dissolve the NaOH in ~ 200
mL of DI-water in a 250 ML volumetric flask. Be sure
that all of the NaOH has dissolved, then dilute to the mark with
DI water, and MIX THOROUGHLY.
Note: Sodium Hydroxide (NaOH) is hygroscopic, which means
it absorbs water from the air. Therefore it is very
important that you recap the sodium hydroxide reagent bottle
immediately after use. Even if someone is waiting for
the bottle, you should recap it before handing the bottle over.
This will also reduce the likelihood that the cap will be
lost. Your grade and the grade of the class will in part be based
on the condition of the lab, hoods, and balances.
DID YOU RECAP THE REAGENT BOTTLE?!
3. Check that the tip of your buret is secure. Many of you will
have 2 piece (Teflon tip body and glass bore) burets.
Look at the image below. Make sure
a. that the Teflon body is a snug fit in the glass bore
b. that the glass tip fits snugly in the Teflon body
c. that the valve turns freely
If the tip comes off during the experiment, you have wasted a
great deal of time and solution.
4. Set up the buret in the buret stand as indicated in the figure
on the next page. Place a waste beaker under the buret.

19. Since we cannot assume that the last student using the buret
adequately cleaned it or that they used NaOH, we want
to be very sure that the only substance in the buret is NaOH,
and of course water. With the stopcock closed, fill your
buret with DI-water. Allow the buret to drain into a 100-mL
waste beaker. Besides rinsing out the buret, this step lets
you
(a) check that your buret tip does not leak.
(b) Check that no droplets are left behind on the buret’s wall as
the buret drains (if droplets are left, the buret
should be cleaned)
(c) Discard the rinsings from your buret in the sink.
7

20. 5. Next add ~ 20 – 25 mL of a NaOH rinsing solution to your
buret. DO NOT RAISE NaOH SOLUTION ABOVE
SHOULDER LEVEL WHILE ADDING SOLUTION TO THE
BURET. BRING THE TOP OF THE BURET DOWN TO
SHOULDER LEVEL. Make sure that the entire inner surface
(including the stopcock) of the buret is rinsed. Allow the
rinsing solution in the buret to drain into a 100-mL waste
beaker. Discard the rinsings from your buret in the waste
carboy.
6. Transfer about 50 mL of your own NaOH solution into a
clean, dry beaker. Very carefully, use a funnel to add your
NaOH solution from the beaker to the buret. DO NOT RAISE
NaOH SOLUTION ABOVE SHOULDER LEVEL. Take
care that you do not over fill the buret. Students often watch
the volume in the buret, as an indicator to stop pouring,
but forget that the volume still in the funnel will also be added
to the buret. The NaOH solution should be slightly
above the 0.00 mL mark.
7. Carefully open the stopcock and allow the solution level to
drop to, or just below the 0.00 mL mark. Note the tip of the
buret, which originally was filled with air, should now be filled
with solution. MAKE SURE THERE ARE NO AIR
BUBBLES.
8. Weigh out 0.5-0.6 g of KHP on the top-loader balance, then
re-weigh it accurately on the analytical balance TO 4
DECIMAL PLACES. Do not add or remove KHP while the
sample is inside the chamber of the analytical balance.
Record the exact mass (4 decimal places), and quantitatively
transfer (transfer every last speck of it) the KHP to a
clean Erlenmeyer flask. Add ~20 mL of DI-water to the flask to

21. dissolve the KHP. Make sure ALL the KHP is
dissolved.
KEEP THE ANALYTICAL BALANCES CLEAN!
YOUR GRADE MAY BE AFFECTED IF THEY ARE NOT!
9. Add 3 drops of phenolphthalein indicator to the KHP solution
in the Erlenmeyer flask. Place the flask under the tip of
the buret. Put a piece of white paper underneath the flask to
help you detect the color change.
10. Check to see that sure there are no air bubbles in the tip of
the buret. If there are, you will need to drain a small
amount into a waste beaker. Record the initial volume of NaOH
solution in your lab notebook. This initial volume
should be between 0.00 and 1.00 mL.
11. For the first trial, you will approximate the volume needed
to reach the endpoint. With constant swirling of the
Erlenmeyer flask, open the stopcock to allow the NaOH to drip
into the Erlenmeyer flask at a rate of several drops per
second. Make sure that the NaOH solution does not adhere to
the sides of the flask (if it does, rinse it down with a
small squirt of DI water). You should slow the rate of addition
as the pink color of the phenolphthalein appears, then
disappears with swirling. Note: the pink color may appear and
disappear but the end point is reached when the entire
8
solution remains pale pink. If it’s hot pink, you have overshot
(gone past) the endpoint. It is unlikely that you will
achieve an accurate endpoint during your first trial. At the

22. endpoint, the solution should be a very, very faint pink and
the color change should occur from just one last drop of NaOH.
12. Record the final volume of NaOH in your lab notebook.
The difference between the final and initial volume represents
the total volume of NaOH added. For this first trial, the volume
is an estimate of where the endpoint occurs. Save the
titrated sample for comparison to the next trial.
13. Repeat steps 8-12. Make sure you have sufficient NaOH in
your buret before starting the second trial. Some burets
will not deliver titrant when the meniscus reaches 45 mL.
Calculate the molarity of NaOH for this second trial.
14. Refill your buret and check for air bubbles in the tip before
proceeding. Repeat steps 8-12.
15. You will have a total of three trials, one estimate (#1), and
two data trials (#2 and #3). The molarity of NaOH
calculated from trials 2 and 3 should agree closely. If the
difference is greater than 1%, you should do a fourth trial.
Use the molarities from your two best trials to calculate the
average molarity of the NaOH solution.
Part B: Determining the Concentration of Acetic Acid in a
Vinegar
Solution
16. Obtain about 25 mL of vinegar solution in a clean, dry,
labeled beaker. Record the mass percent of acetic acid in the

23. vinegar solution which appears on the label of the bottle.
17. Using a pipet pump and a 5 mL pipet, pull a small volume of
vinegar solution into the pipet and roll it around to rinse it.
Discard the rinse into a waste beaker.
18. Draw more vinegar solution into the pipet, but this time go a
little above the ring in the neck of the pipet. Do not draw
any solution into the pipet pump. Wipe the tip of the pipet with
a clean tissue (Kimwipe).
19. Holding the tip against the inside edge of the beaker, slowly
release the solution, using the pump, until the bottom of
the meniscus coincides with the ring.
20. Slowly remove the pipet from the beaker and place it at a
45o angle over the neck of a clean Erlenmeyer flask. Press
the release lever of the pump (do not use the wheel) so as to
allow the solution to drain into the flask. Then carefully
move the pipet into the flask so as to allow the tip of the pipet
to touch the glass surface for about 10 seconds.
Withdraw the pipet from the flask. You have now dispensed
5.00 mL of vinegar solution into the flask.

24. 21. Add approx 20 mL of water, and 3 drops of phenolphthalein
to the vinegar solution.
22. Repeat steps 10-12 until you have a total of two consistent
data trials for the vinegar.
23. Use the information from your best two trials to determine
the average molar concentration of the vinegar. From the
molar concentration you should be able to calculate a mass
percent concentration of acetic acid in the vinegar.
24. Unused NaOH (still in volumetric flask or in buret) should
be recycled in the NaOH recycle container. NO WASTE IS
TO BE DISPOSED OF IN THIS CONTAINER.
25. All waste solutions should be disposed of in the waste
carboy.
26. Hand in the yellow copy of your data sheet to your
instructor before you leave the lab.

25. 9
SUGGESTED SETUP FOR LAB NOTEBOOK (All lab data goes
in the notebook; do not record lab data here)
Part A: Standardization of NaOH
TRIAL 1 TRIAL 2 TRIAL 3 TRIAL 4 TRIAL 5
Mass KHP _________ _________ _________ _________
_________
Initial Buret Vol. (Vi) _________ _________ _________
_________ _________
Final Buret Vol. (Vf) _________ _________ _________
_________ _________
Vol. of NaOH used =
DV = Vf – Vi _________ _________ _________ _________

26. _________
Molarity of NaOH _________ _________ _________
_________ _________
Show a sample calculation:

27. Average Molarity of NaOH (using best two trials) _________
Part B: Determining the Concentration of Acetic Acid in a
Vinegar