West Virginia University Acid-Base Classification of Salts Chemistry Lab Report

Creating a solution of known molarity and converting units
In this assignment, you will weigh out a sample of baking soda (NaHCO3) and create a solution of
known molarity.
1. Start Virtual ChemLab, select Reactions and Stoichiometry, and then select Creating a Solution of
Known Molarity from the list of assignments. The lab will open in the Titration laboratory.
2. In the laboratory, a bottle of baking soda (sodium bicarbonate, NaHCO3) will be next to the
balance, and an empty beaker will be on the stir plate. Drag the empty beaker to the spotlight
next to the balance, click in the Balance area to zoom in, place a weigh paper on the balance, and
tare the balance.
3. Open the bottle by clicking on the lid (Remove Lid). Pick up the Scoop and scoop out some
sample by first dragging the scoop to the mouth of the bottle until it snaps in place. Then pull the
scoop down the face of the bottle. As the scoop is dragged down the face of the bottle it will
pickup different quantities of solid. Select the largest sample possible and drag the scoop to the
weighing paper on the balance until it snaps in place and then let go. This will put approximately
1 g of sample on the balance. Repeat with a second Scoop. Record the mass of the NaHCO3 in
the data table.
4. Drag the weigh paper to the beaker and add the NaHCO3 sample to the beaker. Click on the green
Zoom Out arrow to return to the laboratory.
5. Drag the beaker to the 50 mL graduated cylinder (the largest one) by the sink and empty the
sample into the cylinder. Hold the cylinder under the tap until it fills with water to make an
aqueous solution of NaHCO3. (When the graduated cylinder is full it will automatically snap back
into place.) Note that normally the solid is added and dissolved before the volume is measured
when making a molar solution. Also, chemists normally use a volumetric flask for making molar
solutions, but this is not available in the simulation.
6. Determine the liters of solution in the cylinder and record the data in the data table.
7. Determine the moles of NaHCO3 in the sample and record the data in the data table.
Data Table
mass NaHCO3
moles NaHCO3
liters solution
8. The molarity is calculated using the formula:
mass solute⁄
moles solute
molar mass solute
molarity =
=
L solution
L solution
Calculate the molarity of the NaHCO3 solution (show working).(give your answer to 3 sig figs)
The virtual lab that you just carried out made a few critical mistakes in the preparation of a solution
A better approach is outlined in this video.

Can you identify three differences between the virtual lab approach and the approach used in the
video
1.
2.
3.
Converting solution units (show working)
Molality
Molality is used when dealing with colligative properties such as freezing point depression
(i.e. the freezing point of a solution vs. that of the pure solvent). It is defined as the moles of
solute divided by the mass (in kg) of solute, so its units are mol/kg.
1 mol/kg = 1molal = 1 m
Molality can be determined with the help of these two formulae
Mass solution = mass solute + mass solvent
molality =
mass solute⁄
moles solute
molar mass solute
=
kg solvent
kg solvent
If the density of the solution from the virtual lab is measured to be 1.047 g/mL, calculate the molality
of the NaHCO3 solution in units of mol/kg (Hint: use the density to work out the mass of the solution
and then use the first formula to work out the mass of the solvent).(give your answer to 3 sig
figs)(show working)
Mass percent
This is another common unit for measuring concentration
mass % solute=
mass solute
x 100
mass solution
Calculate the mass percent of sodium bicarbonate in the solution (show working – answer to 3 sig
figs).
Graphs
There are no graphs to include
Acid-Base Classification of Salts – Part I
In this assignment you will be asked to classify aqueous solutions of salts as to whether they are
acidic, basic, or neutral. This is most easily done by first identifying how both the cation and anion
affect the pH of the solution and then by combining the effects. After predicting the acid-base
properties of these salts, you will then test your predictions in the laboratory.
1. State whether 0.1 M solutions of each of the following salts are acidic, basic, or neutral. Explain
your reasoning for each by writing ionic equations to describe the behavior of each salt in water:
NaCN, KNO3, NH4Cl, NaHCO3, and Na3PO4.
NaCN:
KNO3:
NH4Cl:
NaHCO3:
Na3PO4:
Once you have predicted the nature of each salt solution, you will use Virtual ChemLab to confirm
your prediction. Each solution must be approximately 0.1 M for your comparisons to be valid. Most
of the solutions in the Stockroom are approximately 0.1 M already. Three solutions must be prepared
from solid salts. One of these salt solutions is already prepared and on the lab bench ready for you to
measure the pH.
2. Start Virtual ChemLab, select Acid-Base Chemistry, and then select Acid-Base Classification of
Salts from the list of assignments. The lab will open in the Titrations laboratory.
3. On the stir plate, there will be a beaker of 0.10 M ammonium chloride (NH4Cl) that has already
been prepared. The pH meter has been calibrated and is in the beaker. Record the pH of the
NH4Cl solution in the data table on the following page. When finished, drag the beaker to the red
disposal bucket, and drag the bottle of NH4Cl to the stockroom counter.
4. Click in the Stockroom to enter. Double-click on the NH4Cl bottle to return it to the shelf and then
double-click on the NaHCO3 and KNO3 bottles to move them to the Stockroom counter. Return to
the laboratory.
5. Open the beaker drawer (click on it) and drag a beaker to the spotlight next to the Balance. Click
and drag the bottle of NaHCO3 and place it on the spot light near the balance. Click in the Balance
area to zoom in. Place a weigh paper on the balance and tare the balance. Open the bottle by
clicking on the lid (Remove Lid). Pick up the Scoop and scoop up some salt by dragging the Scoop
to the mouth of the bottle and then down the face of the bottle. Each scoop position on the face of
the bottle represents a different size scoop. Pull the scoop down from the top to the first position
(approximately 0.20 g) and drag it to the weigh paper in the balance until it snaps into place.
Releasing the scoop places the sample on the weigh paper. Record the mass in the table below.
Now drag the weigh paper from the balance to the beaker until it snaps into place and then empty
the salt into the beaker. Return to the laboratory and drag the beaker to the stir plate.
6. Drag the 25 mL graduated cylinder to the sink under the tap until it fills. When filled, it will
return to the lab bench and will indicate that it is full when you place the cursor over the cylinder.
Drag the 25 mL cylinder to the beaker on the stir plate and empty it into the beaker. Place the pH
probe in the beaker and record the pH in the data table. Drag the beaker to the red disposal bucket.
Double-click the bottle of NaHCO3 to move it to the Stockroom counter. Repeat steps 5 and 6 for
KNO3.
7. Click in the Stockroom. The stock solutions of NaCN and Na3PO4 are already approximately 0.1
M. Double-click each bottle to move them to the counter and return to the laboratory. With these
solutions you can pour a small amount into a beaker that you have placed on the stir plate and
place the pH probe in the solution. Record each pH in the data table. Drag each beaker to the red
disposal bucket when you have finished. Were your predictions correct?
Data Table
solution
NH4Cl
mass used
(g)
–
concentration
(M)
pH
acidic, basic or
neutral
NaHCO3
KNO3
NaCN
–
Na3PO4
–
8. For some of the salts, the concentrations were not quite 0.1 M. If you had used exactly 0.1 M
solution would this change whether they are acidic basic or neutral? Discuss
9. Use Le Chatelier’s Principle to qualitatively determine what the pH of a 0.2 M NaCN would be
compared to the pH of a 0.1 M solution of NaCN. Explain your reasoning.
10. . Use Le Chatelier’s Principle to qualitatively determine what the pH of a 0.2 M NH4Cl would be
compared to the pH of a 0.1 M solution of NH4Cl. Explain your reasoning