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Chem 141 Titration Lab Lecture Notes Samuel A. Abrash
Q: What is the purpose of this experiment?
The purpose of this experiment is to explore two techniques by which we answer the question: How much? One technique we’ll be using to answer this is weighing. The other is titration.
Q: What do we have to weigh?
You’re going to weigh two different substances, NaOH, sodium hydroxide, on a top loading balance, and KHP, potassium hydrogen phthalate, on an analytical balance.
Q: Two different balances? What’s the difference?
The analytical balance is a much more precise instrument, allowing mass determinations to ±0.0001 g. The top loading balance typically allows masses to be determined to only ±0.01 g. Thus the analytical balance is 100 times more precise than the top loading balance.
Q: Why do we use two balances with two different precisions?
You need precise determination of the mass of KHP, because weighing is the only method you’ll use to quantify the amount of KHP in this experiment. Therefore you want to be able to record the maximum number of significant figures.
It’s ok to use the less precise top loading balances to find the approximate weight of the NaOH, since you’ll be making a finer determination of the amount of NaOH by titration later in the lab.
Q: But why can’t we just use the analytical balance to determine the exact mass of the NaOH?
NaOH is a hygroscopic compound. This means that it easily absorbs water from the air. This means that the mass of any solid “NaOH” sample that you measure will actually be a combination of NaOH and water. Since the fraction of NaOH that you’re weighing is unknown, you might as well use the less precise measuring device.
Q: You say we’ll use titration to determine the amount of NaOH. What’s titration?
Titration is a way of determining how many moles of a chemical are in a solution phase sample by completely reacting it with another chemical.
Q: What reaction do we do to determine the number of moles of NaOH?
We’re reacting KHP with NaOH. The products are water and a salt. This type of reaction is called a neutralization reaction because two caustic compounds, KHP, a mild acid, and NaOH, a strong base, are replaced by water and a weaker base.
Q: How much NaOH does it take to completely react with a sample of KHP?
This is based on the balanced reaction for neutralization of an acid by NaOH: HA NaOH NaA H O 2
Note that in this reaction there is a 1:1 mole ratio of acid to base. This equation is valid for all acids that contain only one acidic proton. All of the acids that you’ll be titrating in today’s experiment have just one acidic proton. This means we can write the reaction of KHP with NaOH as: KHP NaOH H O 2 NaKP
Q: Why is the 1:1 mole ratio significant?
It tells us that our reaction is complete when the number of moles of NaOH that we add is equal to the number of moles of KHP that are initially present.
Q: How can we tell how many moles of KHP are initially present?
Since you’ll have weighed out a sample of KHP, you simply divide by the molar mass of the KHP: mass KHP mol KHP molar mass KHP
Q: How do we determine the molar mass of the KHP?
The formula of the KHP is written in the lab manual. All you need to do is multiply the number of each type of atom by its atomic mass (which you can find on the back cover of the lab manual), and then add up the total masses for each element.
For example, percholoric acid, with formula HClO 4 , has a molar mass given by
molar mass HClO 4 (^) 1 xmH 1 xmCl 4 xmO 1.0079 35.453 4 15.999 x 100.457 g / mol
Q: Is there a special name for the point at which we’ve put equal moles of acid and base into our flask, so that the reaction is complete?
We call this the endpoint of the titration.
Q: How can you tell that you’ve reached the endpoint?
For some reactions there is a natural color change that can be used to tell when the reaction is complete. However, for our reactions today, no color changes occur.
Q: How can you tell how closely the three values agree if you’re only reporting the average?
You can tell how close the values are by also reporting the standard deviation, a statistical measure of the precision of a group of measurements. The standard deviation is calculated with the formula:
( )^2 1
s M^ M^ i N
Where M is your average molarity, Mi represents the individual measurements, and N is
the total number of molarities you measured.
What this really means is you take the absolute value of the difference between each of your three points and the average, then add the three values up, divide by two, and take the square root.
The smaller s is, the more precise your measurement is.
Q: Ok. NOW we’re done, right?
Not quite. At this point we’ve determined the concentration of the solution of NaOH we made, a process called standardization. Now we’re going to use our standardized solution to determine the concentration of acid in a commercial bottle of vinegar.
Q: What kind of acid is in vinegar?
Acetic acid, C 2 H 3 O 2 H. (Only the final H is acidic.)
Q: Why do we need to measure the concentration of the acid? It’s right there on the bottle, 5%!
One of the common tasks given to chemists is to confirm the claims made by manufacturers, whether they make cigarettes and claim that they’re low in tar, or orange juice to make sure that the vitamin C content is what is claimed. For solutions, titration is the method.
Q: Uh, by the way, what does 5% acid mean?
It means that 5% of the mass of the sample is acetic acid?
Q: How do we find the mass percent from the molarity we determine?
First you need to convert the moles of acetic acid to mass of acetic acid, but this is easy, since you just multiply the number of moles by the molar mass:
macetic acid nacetic acid M M.. acetic acid
Next you need to determine the total mass of your vinegar sample. To do this you need to know the density of your liquid, but this is given (on the final data sheet, page 12) as 1.005 g/mL. The total mass of liquid is given by the product of the volume and the density.
mvinegar Vvinegar dvinegar
The mass percent is then given by the mass of the acetic acid over the total mass multipled by 100:
%^ acetic acid 100% total
m acetic acid x m
Q: You reported the density of the vinegar to three decimal places. That means our volume measurement has to be just as precise. How can we measure volume that precisely?
We use a device called a volumetric pipette, which I’ll demonstrate for you at the end of the lab lecture.
Q: Is there anything different we’re going to do in our second titration?
You’ll have to do a 10fold dilution of your NaOH solution before using it in the titration.
Q: Why?
Your NaOH solution is much more concentrated than the vinegar. You’d use so little NaOH that your results would be very imprecise. Decreasing the concentration of the NaOH increases the volume you’ll use, and increases the precision of the measurement.
Q: How do we do the dilution.
You’ll transfer about 30 mL into a clean beaker. Transfer 10 mL from the beaker using a volumetric pipette into a 100 mL volumetric flask. Then fill with distilled (RO) water until the meniscus touches the line in the neck.
Q: How do we calculate the new concentration?
There’s a rule called the lever rule for dilution calculations:
M (^) initial Vinitial M (^) finalV final
If we solve this for the final molarity, it becomes
initial final initial final
V
M M
V
Q: How many times to we have to repeat this measurement?
Next week you’ll turn in the data report sheets on pages 11 and 12, sample calculations for molarity of the NaOH solution, for the concentration of the diluted NaOH solution and for the mass % of acetic acid (see page 7), and a brief discussion of the results including the precision, and accuracy of your results, and any sources of error in the procedure as you carried it out.
Honor Stuff: You may work together with your lab partner on all of the calculations that will be turned in on the two data report sheets. The brief discussion (part iv, on page 8) must be your individual effort. Don’t forget to sign the pledge on what you turn in.
Next week: Qualitative Analysis of Reactions – Prelab and Notebook prep due at the beginning of lab.
