UP Chemistry Acid Base Titration Lab Report

i need a lab report about Acid-Base Titration

Laboratory Report

Reports consist of the following sections:-Report Sheets – the completed Data and Report Sheets cut out of your lab manual.Calculations – If there are calculations, you must clearly and logically show all of the calculations to getfull credit.Discussion – Explain your results, referring to the chemistry theory where appropriate. Did the dataconform to what was expected? What does it show in relation to the experiment? If the results are notwhat was expected explain why. Possible sources of error (estimate experimental error in numericalresults). Conclusion – briefly state any final result(s), or conclusions that can be drawn from the experiment.References – any sources consulted for the writing of the report. These should be referenced (cited) inthe body of the text. 

The discussion should be at least one page.

Put the calculation separately.

The conclusion should NOT be short

here all information including titration (volumetric) Experiment instruction including worksheet, and calculation links. Also attached separately table 2 and table 3 data screen shot. Table 1 can be calculated based on instruction in the lab manual. In Table 1 KHP concentration 0.001572 M to standardize NaOH in table 2. on FIELS.

GENERAL CHEMISTRY LAB SERIES
UNIVERSITY OF BRIDGEPORT-CHEMISTRY DEPARTMENT
Acid-Base Titration (Volumetric Glassware uses)
OBJECTIVES
• To standardize a NaOH(aq) solution by titration
with potassium hydrogen phthalate.
• To determine the acid concentration of an
unknown solution by titration with a standard
NaOH(aq) solution.
• To master using volumetric flasks, volumetric
pipets, and burets.
• To gain understanding of titration calculations.
BACKGROUND INFORMATION
General Overview
A titration is an analysis in which a solution with
a known reagent concentration (the standard) is
used to determine the unknown concentration of a
different reagent (the analyte). The standard and
analyte chemically interact in a reaction with known
stoichiometry. In the most common type of titration,
in an acid/base neutralization reaction occurs.
Titration is a form of volumetric analysis that
requires careful volume measurements.
The
experimental setup for a typical titration is shown in
Figure 1. A buret is a long glass tube with fine
graduated markings that allow for relatively high
precision measurement of delivered volume. A
valve (stopcock) at the bottom of the buret controls
the delivery of solution from the buret to the
receiving flask.
The reagent which is being
dispensed by the buret is known as the titrant.
Depending upon the experimental circumstances,
the titrant may be either standard or analyte.
In the more common configuration, an exactly
measured volume of analyte is transferred to the
receiving flask with a volumetric pipette.
The
standard titrant is then delivered from the buret to
the receiving flask in a controlled manner where it
reacts with and depletes the analyte. The titration is
stopped at the exact point when all of the analyte is
consumed – this is the equivalence point. At the
equivalence point the buret valve is closed, and the
volume of standard titrant delivered is read from the
buret. The known concentration of the standard
titrant, the volume of standard titrant, the volume of
analyte, and the reaction stoichiometry are then
used to calculate the concentration of the analyte.
Dr. H.
Figure 1: A typical titration apparatus
Indicators and Endpoints
So how do you know when you reach the
equivalence point? For most titrations, the analyte,
titrant, and product solutions are all colorless and
thus indistinguishable. To make the equivalence
point observable, an indicator is added to the
sample solution. The indicator (In) is a substance
that changes color when it reacts with the titrant (T).
In (Color 1) + T → In (Color 2)
Prior to the equivalence point, the analyte is in
excess and quickly consumes any added titrant and
the indicator retains its original color. However, at
the equivalence point, all of the analyte has been
consumed. Thus, the next drop of titrant will react
with the indicator and convert it into its product color.
In this experiment, we will be using phenolphthalein
as our indicator. For our titrations, phenolphthalein
will initially be colorless and turn pink at the
endpoint. Another indicator, bromothymol blue is
also available in the lab (it will turn from blue to
yellow at the equivalence point).
The point at which the analyst first notices a color
change is known as the endpoint. In an ideal
titration, the endpoint occurs at the exact same
volume as the equivalence point. In practice, it is
possible for the analyst to go past the equivalence
point before a color change is observed. When this
occurs, the difference between the equivalence point
and the endpoint is known as the titration error.
PAGE 2 OF 8
Standards and Standardization
In order to accurately determine the analyte
content of your sample, the titrant concentration
must be accurately known. A primary standard is a
pure substance from which very accurate solution
concentration can be made. Thus, the primary
standard functions as the chemical measuring stick
used to quantify the analyte. However, sometimes
the analyte cannot be directly titrated with the
primary standard. For example, in this experiment,
both the primary standard and the analyte are acids
and cannot react with each other.
In these
circumstances, a secondary standard must be used
as an intermediate measuring stick between the
primary standard and the analyte.
A secondary standard is a substance capable
of reacting with both the primary standard and the
analyte. When a secondary standard is used, the
titration
analysis
occurs
in
two
stages:
standardization and sample analysis.
In the
standardization titrations, the secondary standard
is titrated with the primary standard. From these
titrations, the concentration of the secondary
standard can be accurately determined. Then, the
sample is titrated with the secondary standard to find
the analyte concentration.
In this experiment, our primary standard will be
potassium hydrogen phthalate and our secondary
standard will be sodium hydroxide. Potassium
hydrogen phthalate is an acid salt with the formula
KHC8H4O4, MM = 204.2 g/mol. For shorthand,
chemists frequently refer to potassium hydrogen
phthalate as “KHP” where P2– = C8H4O42–. Because
KHP is a weak acid, an acid-base neutralization
reaction occurs when it is added to sodium
hydroxide:
KHP(aq) + NaOH(aq) → H2O(l) + NaKP(aq)
Or in net ionic form:
–
–
2–
HP + OH → H2O(l) + P
Our experimental procedure will consist of three
parts. First, we will be preparing a primary standard
KHP solution with an accurately known
concentration. Second, we will use this standard
KHP solution to establish the conentration of a
sodium hydroxide solution. Finally, we will use this
secondary standard NaOH(aq) to determine the acid
concentration of a unkown solution.
At this point, you may be wondering why we can’t
eliminate the entire standardization step by simply
preparing an NaOH(aq) solution with an accurate
concentration. Two problems prevent us from doing
this.
First, solid NaOH is very hygroscopic—
meaning it quickly absorbs moisture from the air. If
Dr. H.
a NaOH pellet is placed on an analytical balance in
humid air, you can actually see the NaOH pick up
mass. Because NaOH(s) pellet begins absorbing
water as soon as it is exposed to the air, it is
impossible to know how much its mass is pure
NaOH and how much is absorbed water. This
uncertainty in true NaOH(s) mass will compromise
the accuracy of the NaOH(aq) concentration.
The second problem is that NaOH also reacts
with the carbon dioxide in air:
NaOH(s) + CO2(g) → NaHCO3(s)
Although this reaction is slower, it produces the
same effect as water absorption. In both cases,
impurities will make up a significant proportion of the
mass of a “NaOH(s)” pellet and the resulting
concentrations will be inaccurate. For this reason,
NaOH cannot be used as a primary standard.
Even after your NaOH(aq) solution is prepared,
the absorption of CO2 is a cause for concern. If your
NaOH(aq) solution is left exposed to the atmosphere
for long periods of time, its concentration will
change. For this reason the NaOH(aq) reagent
bottle should be kept tightly capped while not in use.
Titration Calculations
Just as at the heart of every titration is a reaction,
at the heart of every titration calculation is a reaction
equation. Consider the general reaction.
aA + sS → pP+qQ
where A is the analyte, S is the standard, and P and
Q are products. The symbols a, s, p, and q are the
stoichiometric coefficients of the equation. These
coeffiecients give the mole ratio in which reactants
combine and products are generated.
The
relationship between the number of analyte moles
(nA) and the number of titrant moles (nT) is:
n A = nS ´
a
s
(1)
For every possible type of titration calculation,
Equation (1) is always used. Similarly, because the
titrant is always in solution form, the number of
moles of titrant is always calculated as
n = M·V
(2)
where M is the concentration and V is the volume.
The part of a titration calculation which is different
from case to case is the method of calculating the
number of moles of analyte. If the analyte is a
solution then the number of analyte moles is
calculated from equation (2). If the analyte is a
solid, then the number of analyte moles is calculated
from its mass, m, and its molar mass, MM.
PAGE 3 OF 8
n=
m
MM
(3)
While there are many different varieties of
titration calculations, all of them are based upon
some combination of Equations (1), (2), and (3).
A systematic stepwise solution using Equations (1)
and (2) is shown in the sample calculation in Box 1
below.
Streamlining Calculations 1: Combined Equation
The most common types of titration calculations
are those involving two solutions. A standard
titration calculation of this type will consist of three
steps in which equation (2) is used twice and
equation (1) is used once (see sample calculation
below). The calculation can be streamlined by
combining all these steps into one grand equation.
This is achieved through substitution of nA = MA·VA
and nS = MS·VS into equation (1):
M AV A = M SVS ´
a
s
(4)
This equation is valid for any titration involving two
solutions. It is very important to notice the a/s factor
at the end of equation (4). Many students forget this
factor in their calculations and obtain erroneous
answers when a ≠ s. The use of this equation is
shown in the streamlined titration calculation in Box
1 below.
Streamlining Calculations 2: Using Millimole Units
For typical laboratory experiments, titration
volumes are in mL rather than in L. Since the units
of molarity are mol/L, both our analyte and titrant
volumes must be divided by 1000 to calculate the
number of moles. However, a close examination of
the sample calculation below reveals that the mL→L
conversions are unnecessary. Notice that we are
dividing by 1000 in step 2 and multiplying by 1000 in
step 4. Thus, in the final answer, the two mL→L
conversion steps cancel each other out.
So can we just ignore units in our calculations?
No, no, no! Don’t succumb to this temptation which
each year leads students to commit violence to their
Chem120 grades. The proper approach is to do our
multiplication and division of 1000 in our units rather
than in our calculation. Notice what happens if we
divide both the numerator and denominator of
molarity by 1000 and substitute the milli prefix:
M =
1 mol (1/1000) mol 1 mmol
=
=
1L
(1/1000) L
1 mL
Box 1: Sample Titration Calculation
The concentration of a citric acid (H3C6H5O7) solution is being determined by titration with a standard
0.1067 mol/L NaOH solution. When 25.00 mL of the citric acid are titrated, it requires 18.34 mL of the NaOH
standard solution to reach the endpoint. What is the citric acid concentration?
Stepwise Solution
1. Start by writing a balanced equation for the reaction:
H3C6H5O7 (aq) + 3 NaOH(aq) → 3 H2O(l) + Na3C6H5O7 (aq)
2. Next, calculate the number of moles of the NaOH standard using equation (2).
nS = M SVS = 18.34 mL ´
1 L
´ 0.1067 mol L NaOH = 1.9569 ´ 10-3 mol NaOH
1000 mL
3. Now, we use equation (1) to find the moles of the citric acid analyte.
n A = nS ´
a
1 mol H 3C6 H 5O7
= 1.9569 ´ 10-3 mol NaOH ´
= 6.5229 ´ 10 -4 mol H 3C6 H 5O 7
s
3 mol NaOH
4. And finally, we use a rearranged equation (2) to find the citric acid analyte concentration.
n A 6.5229 ´ 10-4 mol H 3C6 H 5O7 1000 mL
MA =
=
´
= 0.02609 mol H 3C6 H 5O7
L
VA
25.00 mL
1 L
Streamlined Solution – Using Rearrangement of Equation (4) and Millimole Units
MA =
M SVS a 0.1067 mmol mL ´ 18.34 mL 1 mol H 3C6 H 5O7
´ =
´
= 0.02609 mol L H 3C6 H 5O7
VA
s
25.00 mL
3 mmol NaOH
PAGE 4 OF 8
Thus, another valid definition of molarity is mmol/mL.
Using this definition allows us to conveniently keep
our titration volumes in milliliters and eliminate two
conversion calculations. The use of millimole units
is shown in the streamlined in Box 1.
Concentration of an Acid Sample
Once we have established the concentration of
our sodium hydroxide solution (the secondary
standard) we will use it to determine the
concentration of an acid solution with an unknown
concentration. It is often the case that the acid
concentration is significantly higher than the
concentration of our standard solution and, if we
titrated directly, it would require an inordinate
amount of titrant. This is not a significant problem as
we can accurately dilute the acid to bring its
concentration down to a more reasonable level. The
following scenario gives an example of the
manipulations and calculations involved.
Acid Sample Concentration Analysis
A chemist needs to accurately determine the
concentration of an industrial waste sample that
contains oxalic acid (a diprotic acid) with a
concentration of about 0.7 mol/L. She has available
the KOH solution standardized in the previous
example (0.1566 mol/L or 0.1566 mmol/mL). Using
a volumetric pipet she transfers a 25.00 mL portion
of the original oxalic acid solution to a 250.0 mL
volumetric flask. After diluting and mixing she uses
a volumetric pipet to transfer 20.00 mL of this
solution into an Erlenmeyer flask. She then titrates
this sample to the endpoint with the standard KOH
solution. It takes 17.52 mL of the KOH solution to
reach the endpoint. What is the concentration of the
original unknown?
Titration Calculation
We will use the streamlined solution shown at the
end of the previous example to calculate MA, the
concentration of the diluted oxalic acid solution (the
analyte). To calculate this concentration we need to
extract the values for the following variables from the
data given in the problem.
MS
VS
VA
a
s
concentration of standard titrant
volume of standard titrant required
volume of analyte titrated
stoichiometric coefficient of analyte
stoichiometric coefficient of standard titrant
We find a and s from the balanced chemical
equation for the reaction. The formula for oxalic acid
is H2C2O4 and it reacts with potassium hydroxide
solution according to the following.
Dr. H.
H2C2O4(aq) + 2KOH(aq) → K2C2O4(aq) + 2H2O(l)
The KOH solution is the titrant so s = 2 mol KOH
and the oxalic acid solution is the analyte, so a = 1
mol H2C2O4.
The concentration of the standard KOH solution
is given as 0.1566 mol/L and 17.52 mL of this
solution was needed to reach the endpoint, thus we
have MS and VS.
Determining VA can be trickier because, with a
dilution step and a titration step, we see several
volumes associated with the oxalic acid solution.
Since we are dealing with the titration step in this
calculation we must look for the volume of solution
that is titrated.
Careful reading reveals that
20.00 mL of the dilute oxalic acid is placed in the
Erlenmeyer for titration, so VA is equal to 20.00 mL.
We now have values for all the variables so it is
just a matter of running the calculation.
MA =
(0.1566 mol KOH L )(17.52 mL ) æ 1 mol H 2C 2O 4 ö
ç
÷
è 2 mol KOH ø
20.00 mL
M A = 0.06859 mol/L H 2C 2O 4
This is not the final answer though – this calculation
gives the concentration of the dilute oxalic acid
sample. But we want to know the concentration of
the original oxalic acid solution, so we have one
more calculation to perform – a dilution calculation.
Dilution Calculation
Dilution calculations are usually presented as
follows:
M iVi = M f V f
(5)
where M represents concentration and V represents
volume (just as with the titration calculation) and the
subscripts indicate initial (i) conditions and final (f)
conditions, that is before and after the dilution.
Sometimes you may see this relation written with 1’s
and 2’s in place of the I’s and f’s, but it works the
same either way.
The original oxalic acid solution concentration is
represented by Mi, thus we need to rearrange
Equation 5 to solve for Mi.
Mi =
M fVf
Vi
(6)
The final concentration, Mf, is the concentration of
the dilute oxalic acid determined by the titration and
calculated above. As part of the dilution process the
chemist withdrew a 25.00 mL sample and placed it
in a 250.0 mL volumetric flask, thus Vi = 25.00 mL
PAGE 5 OF 8
and Vf = 250.0 mL. Substituting into equation (6),
the original solution is found to have an oxalic acid
concentration of 0.6859 mol/L.
Mi =
(0.06859 mol L )(250.0 mL ) = 0.6859 mol
25.00 mL
L
REFERENCES
For additional information on the topics covered
in the Background Information, consult the following
sections from Chemistry (3rd ed.), Nivaldo J. Tro,
Pearson, 2011.
• Solution Concentrations & Dilutions, Section 4.4
• Acid-Base Titrations, Section 4.8
PRACTICAL ASPECTS
Throughout this experiment you will need to
exercise careful laboratory technique in order to
obtain successful results (and thus a good accuracy
and precision score). Review the “Lab Equipment
Instructions” document for the proper use burets and
pipets used in this experiment.
beaker and the funnel with three small portions of
deionized water into the volumetric flask. Fill the
volumetric flask half-full of deionized water, stopper
it, and swirl to dissolve the crystals. When all solid is
dissolved, fill the volumetric flask to the mark with
deionized water and mix thoroughly by inverting 20
times or more.
2. Standardization of the Sodium Hydroxide
Solution.
Buret Preparation
Review the “Lab Equipment Instructions” section
on buret usage before proceeding further. Rinse
your burets thoroughly with tap water and deionized
water. A clean buret will drain freely without forming
drops on the inside surface.
Insure that the
stopcock functions properly; it must turn freely and
must not leak when turned off.
EXPERIMENTAL PROCEDURE
Rinse the buret with three 5-mL portions of the
NaOH solution running a portion of each rinse
solution through the stopcock and tip of the buret.
Discard the rinse solution. Fill the buret with the
NaOH solution. Place the buret cap on the buret to
reduce the rate of reaction between sodium
hydroxide and atmospheric carbon dioxide.
Chemicals:
Standardization Titrations
0.1 M sodium hydroxide (NaOH)
Potassium hydrogen phthalate (KHP),
KHC8H4O4, anhydrous
Dilute acid solutions of unknown concentration
Phenolphthalein indicator solution, or
bromothymol blue solution (an alternate
indicator)
Special Equipment (To be checked out):
50-mL buret;
buret funnel
buret cap
100 mL volumetric flask
One 10 mL volumetric pipet
One pipet bulb
1. Preparation of a Standard KHP Solution:
Refer to the “Lab Equipment Instructions”
document on volumetric flask technique before
reading further. Carefully weigh a clean, dry 50 mL
beaker as accurately as possible. Place between 3
g and 3.5 g of anhydrous potassium phthalate, KHP,
in the beaker and weigh accurately. Transfer the
weighed crystals to a clean 100 mL volumetric flask,
being careful to prevent loss of the crystals. To
ensure every bit of KHP is transferred, rinse the
Review the “Lab Equipment Instructions”
document covering the use of volumetric pipets and
burets before proceeding.
Using a 10 mL volumetric pipet, transfer 10 mL of
the KHP solution into a clean, rinsed 250 mL
Erlenmeyer flask.
Touch off the pipet tip if
necessary and rinse the inside walls of the flask with
a minimum amount of deionized water. Add enough
DI water such that the solution covers the bottom of
the flask. Add 2 drops of phenolphthalein indicator.
(Alternatively, if you have difficulty seeing pink, you
may use thymol blue as the indicator. The color
change for thymol blue is from yellow to blue.)
NOTE: be consistent from one run to the next for
good precision and accuracy.
Your first titration (Trial 0) will be a rough titration
to get an approximate volume of titrant needed to
reach the endpoint. This approximate volume,
although not used to calculate the average
concentration of your unknown, will give you practice
using the buret and will allow you to more quickly
complete subsequent trials.
Record the initial buret volume to the nearest
0.01 mL. Position the Erlenmeyer flask under the
NaOH buret. Place a sheet of white paper or a
paper towel under the flask to help you see the
indicator more easily. Begin titrating the KHP with
PAGE 6 OF 8
NaOH by opening the stopcock of the NaOH buret
almost all the way. Immediately begin stirring the
solution by swirling the flask. As the end point is
approached, the faint pink color will appear in parts
of the unmixed NaOH solution will persist for several
seconds.
For Trial 0 slow the flow somewhat (to a rapid
rate of drops instead of a continuous flow) and stop
when the pink color of the indicator persists. Record
the final buret volume and calculate the volume of
titrant used (final volume − initial volume).
For Trials 1, 2, and 3 follow the same procedure,
except slow the rate of addition when you reach
approximately 90% of the volume of Trial 0. The
addition rate should be such that you are adding
single drops and thoroughly mixing the solution after
each drop is added. The end point is reached when
a single drop (or fraction of a drop) produces a faint
pink color which lasts for at least 30 seconds after
vigorous swirling. The pink color will fade after the
solution has been standing awhile because carbon
dioxide will be absorbed from the air, and will make
the solution slightly acidic again. (You can verify the
effect of carbon dioxide by blowing on your finished
titration solutions while swirling and watching the
indicator change back to its original color.) Record
the final buret volume and calculate the volume of
titrant used (final volume − initial volume).
When the titration is finished, the neutralized
mixture in the titration flask may be disposed of by
pouring down the sink with running water.
Refill your buret with the sodium hydroxide
solution as necessary.
Complete at least three standardization titrations.
Enter your titration data into spreadsheet. Run an
additional titration (Trial 4) if the volume range
between the highest and lowest value of Trials 1, 2,
and 3 is greater than 0.2 mL.
Contact your
instructor if, after your fourth trial, this difference is
still greater than 0.2 mL.
Calculations
From your titration data, calculate the
concentration of your NaOH solution. Because of
the volumetric accuracy of your measurements, you
should be able to report your answers to four
significant digits (unless your titration volumes are
less than 10 mL). Calculate the average, standard
deviation, and relative standard deviation of your
concentration.
We will need to recycle the volumetric glassware
from this part of the experiment for use in the
second part. However, we do not want you to
dispose of your KHP solution unless we know you
Dr. H.
have (reasonably) good data. Have your instructor
or TA check your data, and, once given the OK,
discard the leftover KHP solution in the sink with the
tap water running. Rinse the volumetric flask, flask
cap, and pipet three times with regular tap water and
then do a final rinse with a small amount of DI water.
If you accidentally drew KHP solution into your pipet
bulb rinse it as well.
3. Concentration of an Acid Solution
Each student will be assigned one of the
available acid solutions of unknown concentration
for analysis. Record the letter, chemical name, and
chemical formula of your acid unknown in your
report and in the spreadsheet.
As discussed above, the acid unknowns are too
concentrated to analyze directly so we must dilute
them before analysis. There are common use pipets
associated with each of the acid unknowns. Use the
appropriate common pipet to transfer a sample of
the assigned unknown to your rinsed 100 mL
volumetric flask. NOTE: do not worry if there is
residual DI water in the volumetric flask – it will not
affect your results. Fill the flask about two-thirds full
with DI water, insert the cap, and carefully mix the
solution by inverting the flask at least ten times.
Remove the cap and CAREFULLY fill to the line with
DI water. It is a good idea to use a dropper instead
of your squirt bottle as the dropper gives you more
control – if you go over the line you will need to
dump the solution and start over. Reinsert the cap
and mix again.
Your own 10 mL volumetric pipet should be
rinsed at this point, but it probably contains residual
DI water from the rinsing. This water must be
removed otherwise it will dilute your acid and cause
erroneous results. To rinse the pipet, transfer about
20 mL of the acid from your flask into a clean, dry
50 mL beaker. Draw a small portion of this acid into
your pipet and let it rinse through the pipet into the
sink. Repeat this step with a second portion of acid.
Thought Question: why do we need to remove
residual water from the pipet but not from the
volumetric flask?
Use the rinsed pipet to transfer 10 mL of your
unknown acid from the volumetric flask into a rinsed
250 mL Erlenmeyer flask. Rinse the walls with DI
water and add enough DI water so that the solution
just covers the bottom of the flask. Add your
preferred indicator and titrate the acid with the
standardized NaOH solution to the endpoint. As
with the standardization titrations in Part 2 above,
use the first titration (Trial 0) to get an approximate
volume and then carefully titrate three more samples
(and a fourth if necessary, see below).
PAGE 7 OF 8
When the titration is finished, the neutralized
mixture in the titration flask may be disposed of by
pouring down the sink with running water.
Complete at least three unknown acid titrations.
Enter your titration data into spreadsheet. Run an
additional titration (Trial 4) if the volume range
between the highest and lowest value of Trials 1, 2,
and 3 is greater than 0.2 mL.
Calculations
Using your NaOH titration volumes and the
standardized concentration determined in the first
part of this experiment, calculate the concentration
of the acid. This is the concentration of the acid in
your volumetric flask but we want you to report the
concentration of the acid as it is in the main
container. Remember that you titrated the diluted
acid solution so you will need an additional dilution
calculation to determine the concentration of the
original acid.
Waste Disposal
Pour any unused KHP, NaOH, and unknown acid
solutions together into your largest beaker and mix
by stirring. Flush this mixture down the sink with
copious amounts of running water. Allow the tap
water for run for about a minute after pouring
unused solution in the sink.
Clean your area, replace any equipment you may
have taken from the common equipment drawer,
and make sure all reagent bottles at your station are
closed.
Before leaving the lab…
• Be sure you have cleaned and returned all of the
equipment that you checked out from the
storeroom.
• Be sure you have properly entered and saved all
of your data in the Lab06.xls spreadsheet.
• Police and clean up your area – you will not be
“signed out” until your station is clean.
• Have your laboratory report “signed out” by your
instructor or teaching assistant.
NAME:
INSTRUCTOR NAME:
1. KHP SOLUTION PREPARATION
Data (MUST be completed in ink)
Standard KHP Solution
Mass of KHP + beaker
Mass of beaker (0 if tared)
Mass of KHP
Final Dilution Volume
Calculations (should be completed in pencil)
1.
Concentration of Standard KHP Concentration
Calculate the concentration of your KHP solution in mol/L. Show your work. In your calculations, be sure to
show units, and report your answer to the correct number of significant figures.
2. NAOH STANDARDIZATION
Buret Volume Data (in ink)
* Trial 0
Trial 1
Trial 2
Trial 3
†
Trial 4
Volume KHP Delivered
NaOH buret: Final reading
Initial reading
Volume Titrant Delivered
‡
Instructor Initials for Excluded Data →
* Trial 0 is a practice run to get a rough idea of the volume of titrant needed – do not transfer this value
to the Excel spreadsheet.
†
Run a fourth trial if the volume range between the highest and lowest value of Trials 1, 2, and 3 is
greater than 0.2 mL.
‡
Instructor must confirm and initial appropriate box for excluded data, otherwise five points will be
deducted from accuracy and precision score.
Use the space below to explain: (i) if a known error was made in any of your trials; and/or (ii) your reasoning if
you decide to exclude one of your trials from your results.
Dr. H.
Calculations (should be completed in pencil)
Calculate the exact molarity of your NaOH solution to the proper number of significant digits. You must show a
complete set of calculations for Trial 1. It is not necessary to show your work for successive trials, but you must
record the final result for all trials in a table of your own construction. From your trial results, calculate the
average NaOH molarity, the standard deviation, and the percent relative standard deviation – include the results
of these calculations in your table.
3. UNKNOWN ACID SOLUTION CONCENTRATION
Unknown #: __________
Acid Name: ______________________________
Formula: ______________
Write a balanced chemical equation for your titration reaction.
__________________________________________________________________________________________
Dilution Data
Volume Initial (common pipet volume)
Titration Data
Volume Final (volumetric flask volume)
* Trial 0
Trial 1
Trial 2
Trial 3
†
Trial 4
Volume Analyte, Diluted Acid
NaOH buret: Final reading
Initial reading
Volume Titrant Delivered
‡
Instructor Initials for Excluded Data →
* Trial 0 is a practice run to get a rough idea of the volume of titrant needed – do not transfer this value
to the Excel spreadsheet.
†
Run a fourth trial if the volume range of Trials 1, 2, and 3 is greater than 0.2 mL.
‡
Instructor must confirm and initial appropriate box for excluded data, otherwise five points will be
deducted from accuracy and precision score.
Use the space below to explain: (i) if a known error was made in any of your trials; and/or (ii) your reasoning if
you decide to exclude one of your trials from your results.
Titration Calculations (should be completed in pencil)
Calculate the concentration of your diluted acid sample to the proper number of significant digits. You must show
a complete set of calculations for Trial 1. It is not necessary to show your work for successive trials, but you
must record the final result for all trials in a table of your own construction.
Dilution Calculations (should be completed in pencil)
Calculate the concentration of your original acid sample to the proper number of significant digits. You must show
a complete set of calculations for Trial 1. It is not necessary to show your work for successive trials, but you
must record the final result for all trials in a table. From your trial results, calculate the average concentration of
your unknown, the standard deviation, and the percent relative standard deviation – include the results of these
calculations in your table. Record your final results in the summary section below.
Results Summary:
Unknown Number: __________Acid Name: ______________________________
Formula: ______________
Original Solution Concentration (avg.): ___________________ +/– Standard Deviation: ___________________
Relative Standard Deviation: ____________________
Dr. H.