Alternate Procedure / Data Page / Pre-Lab

Procedure and Data Page for CHM 2045-L Experiment #24:
Molar Mass of an Unknown Acid
This procedure replaces the procedures shown on pages 192 & 193 in your lab manual; the data page
replaces the data page 195 & 196 in your manual. There is also an advanced study assignment at the
end of this document which replaces the advanced study assignment on page 197 in your lab manual.
Prepare for the experiment by reading the “Volume” section Appendix IV in your lab manual if
you are unfamiliar with the procedure of titration; this explains how to use and properly read a burette.
You will be expected to understand how to use a burette and how to perform a titration before
beginning this lab.
Part A: Standardization of the NaOH solution
The purpose of this part of the experiment is to determine the precise concentration of the
NaOH solution that you will use for part B; the process of determining a precise concentration for a
solution is called standardization.
Clean a burette by rinsing it with water, using first tap water and then a final rinse with distilled
or de-ionized water. Draw about 75 mL of NaOH from the container provided in the lab into a clean, dry
container – 75 mL should be enough for the entire experiment -- please do not take more than you
need! Rinse the clean burette with a few mL of the NaOH solution; remember to allow a few mL of
NaOH to run through the tip of the burette to rinse it also. Discard this rinse. Now fill the burette to
approximately the top mark and record the initial burette reading. (Remember that the correct
precision for burette reading is two places past the decimal).
Your instructor will provide a burette containing standardized HCl at the front desk of the lab.
Bring two clean 250 mL Erlenmeyer flasks to the front desk; they don’t necessarily have to be dry, but
make sure there is no tap water in either flask. Record the initial burette reading before dispensing your
HCl sample, then dispense approximately 20 mL of HCl into your first flask. Record the final burette
reading. Repeat this procedure for the second flask; each flask should now contain about 20 mL of the
HCl solution. One of these will be for trial 1 and the other will be for trial 2 – don’t get them mixed up!
Make sure you record the precise quantity of HCl used (from your final and initial burette readings); also
record the provided concentration of the HCl solution. Add 2 or 3 drops of phenolphthalein solution. At
your desk, place a white piece of paper under the first flask to aid in the detection of any color change.
Begin the titration by slowly and carefully adding NaOH from the burette at your desk to the
Erlenmeyer flask containing the HCl solution. Notice that a pink color will briefly appear, and then
rapidly disappear as you add the NaOH solution. Keep a close eye on this color change, as it will aid you
in determining when you are close to the endpoint. Swirl the flask gently and continuously as you add
the NaOH. When the pink color begins to last longer (2 or 3 seconds) before fading, you are getting
close to the endpoint, and you want to be very careful not to go past the endpoint; you should add the
NaOH drop-by-drop. The endpoint of the titration is the point at which the solution turns pink and
remains that color; you should stop immediately at the endpoint. If you go past the endpoint, even by a
few drops, you results will be affected. Swirl the flask for a few seconds to ensure that the pink color
does not fade. Once the endpoint is reached, record the final burette reading for your NaOH burette.
Note: there is no way to tell by looking at the solution how far past the endpoint you went. A
solution looks just the same if you go 2 drops past as if you go 200 drops past the endpoint! Your
instructor is not going to be able to ascertain by looking at your solution if you went too far past the
endpoint... the ONLY way to make sure you stop precisely at the endpoint is by working slowly and
carefully, and stopping immediately when the solution turns pink.
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Repeat the titration with your second HCl sample.
Using equation 3 below, calculate the molarity of your base solution, M OH - , for each trial. Use
the average value of this molarity for your calculations in Part B.
Part B: Determination of the Molar Mass of an Acid
Obtain two clean, dry 250 mL Erlenmeyer flasks and a sample of unknown acid. Into each flask,
measure about 0.25 grams of the acid. Be careful with this measurement! If you start with too much
acid, then your titration may require more than 50 mL of NaOH, and you will have to refill your burette
during the titration (this will complicate your procedure and data collection).
To the first flask, add 50 mL of distilled water and 2 or 3 drops of phenolphthalein solution.
Swirl for a few minutes to dissolve the acid. The acid may not dissolve completely, but don’t worry
because it will dissolve as you proceed with the titration.
Refill your NaOH burette before beginning; some of these unknowns may require almost a
complete burette to reach the endpoint. Record the initial volume of a NaOH before beginning. Titrate
the acid solution as you did in part A, working slowly and carefully and stopping immediately when the
solution turns pink. Record the final reading on the NaOH burette at the endpoint.
Refill your NaOH burette and repeat the titration for the second sample of the unknown acid.
Calculations:
for Part A, you need to calculate the molarity of the NaOH solution. You have been provided with the
molarity of the acid M acid , and you measured in the experiment the volumes of HCl and NaOH (V acid and
V base ).
moles of HCl = moles of NaOH added during titration
Because the moles of each substance is equal to the Molarity x Volume, this equation can also be
expressed:
M acid x V acid = M base x V base (equation 3)
(Notice that this is the same as Equation 3 in your lab manual). You know M acid , V acid , and M acid , so you
can solve for M base . Use the average value from your two trials to complete the calculations in part B.
For Part B, you must remember that the acid unknown and the NaOH react in a 1:1 molar ratio.
Therefore the number of moles of acid in your unknown are equal to the number of moles of NaOH that
you added during the titration. The moles of NaOH can be calculated from the volume of NaOH used (in
LITERS!) and the molarity of the NaOH (which you calculated in Part A):
moles of acid in unknown = moles of NaOH added during titration = M x V
Remember to convert your volumes to Liters for this calculation! Once you know how many moles of
acid were in your unknown, you can determine the molar mass of the acid by the ratio of grams to
moles:
grams of unknown
MM =
moles of acid in unknown
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Data Page
Name __________________________________________________________
Part A - Standardization of NaOH solution
Molarity of standardized HCl (provided by your instructor): ________________________________ M
Trial 1 Trial 2
Initial Reading, HCl burette ____________ mL _____________ mL
Final Reading HCl burette ____________ mL _____________ mL
Initial Reading, NaOH burette ____________ mL _____________ mL
Final Reading, NaOH burette ____________ mL _____________ mL
Part B - Determination of Molar Mass of an Unknown Acid
Unknown number
Mass of sample 1
Mass of sample 2
__________
____________ grams
____________ grams
Trial 1 Trial 2
Initial Reading, NaOH burette ____________ mL _____________ mL
Final Reading, NaOH burette ____________ mL _____________ mL
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Calculations
Name __________________________________________________________
Part A - Standardization of NaOH solution
Trial 1 Trial 2
Total volume HCl ____________ mL _____________ mL
Total volume NaOH ____________ mL _____________ mL
Trial 1 Trial 2
Molarity of NaOH ____________ M _____________ M
(use equation 3)
Average Molarity
_____________ M
Part B - Determination of Molar Mass of an Unknown Acid
Trial 1 Trial 2
Mass of sample __________ grams __________ grams
Volume of NaOH used __________ mL __________ mL
Moles of NaOH used __________ moles __________ moles
(this is also the moles of acid in the unknown)
Molar Mass of acid __________ g/mole __________ g/mole
Average Molar Mass
Unknown number
__________ g/mole
__________
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Advance Study Assignment Molar Mass of an Acid - CHM 2045
CHM 2045-L
Name _____________________________________________________________________________
note: This advance study replaces the one on page 197 in your lab manual. You don’t have to do both
of them!
1. 7.0 mL of 6.0 M NaOH solution are diluted by adding water until the total volume is 400. mL. You are
asked to find the molarity of the resulting solution.
a. First find out how many moles of NaOH there are in 7.0 mL of 6.0 M NaOH, using Equation 1
in the discussion in your lab manual. Note that the volume must be in Liters.
_____________ moles NaOH
b. Since the total number of moles of NaOH is not changed by adding water, the molarity after dilution
can also be found using Equation 1, using the final volume of the solution. Calculate that final molarity.
___________________ M NaOH
2. In an acid/base titration, 22.13 mL of NaOH solution are needed to neutralize 24.65 mL of a 0.1094 M
HCl solution. Calculate the molarity of the NaOH, using Equation 3 in the discussion section of your lab
manual:
_____________________ M NaOH
3. A 0.2678 gram sample of an unknown acid requires 27.21 mL of 0.1164 M NaOH for neutralization to
the phenolphthalein endpoint. The unknown acid reacts with the NaOH in a 1:1 molar ratio.
a. How many moles of NaOH were used in the titration?
____________________ moles NaOH
b. How many moles of acid were present in the unknown sample?
____________________ moles acid
c. What is the molar mass of the acid (use Eq. 4 in your lab manual)?
______________________ grams/mole
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