chemistry
Experiment #3a: Aluminum Content via REDOX
Reaction
Objective
To determine the aluminum content in commercial samples through stoichiometry and a standard curve of the volume of hydrogen gas produced versus the mass of aluminum consumed.
Introduction
Stoichiometry
Reaction stoichiometry is the numerical relationships between chemical amounts in a balanced chemical equation. Using
stoichiometry allows us to predict the amounts of products that will form or amounts of reactants that will be consumed
during a chemical reaction. In order to use stoichiometry to predict such amounts, the chemical equation must be balanced.
As an example, let’s look at the following chemical equation:
𝐴 + → 𝐴2𝐵3
It should be fairly obvious that the above equation is NOT balanced. In order to balance this equation, we need to look at
the relative numbers of substances (A & B) on both sides of the equation. On the left hand side (reactants side), we have
one A and one B. On the right hand side (products side), we have two A (from A2) and three B (from B3). We must have
equal relative amounts on both sides. This one is rather simple to solve. Using coefficients (NOT subscripts), we can produce:
2𝐴 + 3 → 𝐴2𝐵3
This equation is now balanced.
Balanced equations can be used to calculate the amount of reactants used or the amount of products formed in
a chemical reaction. For example, using the balanced reaction below,
𝐶2𝐻4 + 3𝑂2 → 2𝐶𝑂2 + 2𝐻2𝑂
the amount of CO2 produced can be calculated when 40.0 grams of C2H4 is reacted with excess O2 (excess means
that there is more than enough O2 for the reaction to go to completion). For such a calculation, we can use the
following general process:
The first step is to convert 40.0 grams of C2H4 into moles of C2H4 using the molar mass (28.05 g/mol). This can be
grams C2H4 moles C2H4 moles CO2 grams CO2
done using the factor-label method:
40.0 𝑔 𝐶2𝐻4 × 1 𝑚𝑜𝑙 𝐶2𝐻4
28.05 𝑔 𝐶2𝐻4 = 1.43 𝑚𝑜𝑙 𝐶2𝐻4
Next, using the balanced chemical equation, determine the number of moles of CO2 produced when
1.43 moles of C2H4 are consumed:
1.426 𝑚𝑜𝑙 𝐶2𝐻4 × 2 𝑚𝑜𝑙 𝐶𝑂2
1 𝑚𝑜𝑙 𝐶2𝐻4 = 2.86 𝑚𝑜𝑙 𝐶𝑂2
The last step is to convert 2.86 moles of CO2 into grams of CO2 using the molar mass (44.01 g/mol):
2.86 𝑚𝑜𝑙 𝐶𝑂2 × 44.01 𝑔 𝐶𝑂2
1 𝑚𝑜𝑙 𝐶𝑂2 = 125.9 𝑔 𝐶𝑂2
Alternatively, the entire process can be done at one time:
40.0 𝑔 𝐶2𝐻4 × 1 𝑚𝑜𝑙 𝐶2𝐻4
28.05 𝑔 𝐶2𝐻4 ×
2 𝑚𝑜𝑙 𝐶𝑂2
1 𝑚𝑜𝑙 𝐶2𝐻4 ×
44.01 𝑔 𝐶𝑂2
1 𝑚𝑜𝑙 𝐶𝑂2 = 125.9 𝑔 𝐶𝑂2
This process can also be used in conjunction with the ideal gas law to convert from volume of gas of one of the products into the amount of mass of the reactant needed. For example, the mass of C2H4 used to form 2.00 L of CO2 can be determined at a pressure of 1.00 atm and a temperature of 293 K. First, the number of moles of CO2 must be calculated from the ideal gas law:
𝑛 = 𝑃𝑉
𝑅𝑇
Where P is the pressure (in atm), V is the volume (in L), T is the temperature (in K), and R = 0.08206 LatmK-1mol-1. Substituting into this equation gives:
𝑛 = (1.00 𝑎𝑡𝑚)(2.00 𝐿)
(0.08206 Latm𝐾−1𝑚𝑜𝑙−1)(293 𝐾) = 0.0832 𝑚𝑜𝑙 𝐶𝑂2
From this, the mass of C2H4 used can be calculated using stoichiometry and the molar mass of 28.06 g/mol:
0.0832 𝑚𝑜𝑙 𝐶𝑂2 × 1 𝑚𝑜𝑙 𝐶2𝐻4
2 𝑚𝑜𝑙 𝐶𝑂2 ×
28.06 𝑔 𝐶2𝐻4
1 𝑚𝑜𝑙 𝐶2𝐻4 = 1.17 𝑔 𝐶2𝐻4
Experimental Procedure
This experiment will be performed in groups of 3-4. Each team will be using pure aluminum metal and two
commercial samples of aluminum foil and will set up a Hydrogen Collection Apparatus, as shown in Figures
1 and 2. Your instructor will prepare a setup ahead of time to serve as a reference.
Figure 1: Hydrogen Collection Apparatus Figure 2: Hydrogen Collection Apparatus
Part I: Measurement of Hydrogen Gas
Assemble the apparatus shown in Figure 1 at one of the lab sinks.
First, make sure that the tub is filled with water, Figures 2. Obtain and completely fill a large graduated
cylinder (250-mL) with water from the tub. Carefully invert the graduated cylinder, submerge it inside the
plastic tub in the sink and make sure that no air is trapped at the top of the graduated cylinder. Make sure
that you do not raise the graduated cylinder above the water level in the tub, or else you will need to start
over.
Connect the tubing to the filter flask, and insert the open end into the mouth of the graduated cylinder
(while under water). Obtain a stir bar and stir plate. Place the stir bar inside the filter flask. Place the
filter flask on top of the stir plate. Obtain a piece of pure aluminum wire. Make sure to clean the wire
using steel wool. Add the sample of metal to the filter flask and record the mass in your notebook.
Using the provided syringe, add 25mL of 8M hydrochloric acid to the filter flask. Turn the stir plate on
and allow the reaction to stir until no more gas is evolved (evidenced by consumption of the metal in
the filter flask). Read and record the volume of hydrogen gas produced from the graduated cylinder in a
table in your notebook. Repeat this with three additional sizes of aluminum. You may need to cut one
of the longer pieces of aluminum.
Mass of Metal Volume of Hydrogen Pure
Aluminum
1
2
3
4
Repeat the procedure using Reynolds Wrap aluminum foil and Great Value aluminum foil. Only perform
one trial with 50-110 milligrams of each type of aluminum foil. Make sure to clean both sides of the
aluminum foil using steel wool. This removes the lacquer on the foil, speeding up the dissolving process
and giving a more accurate starting weight. Read and record the volume of hydrogen gas produced in a
table in your notebook.
Mass of Metal Volume of Hydrogen
Reynolds Wrap
Great Value
Part II: Data Analysis
With the data obtained, construct a standard curve using the data you obtained for the pure
aluminum metal. Plot the volume (mL) of gas produced versus the mass (g) of pure aluminum metal
consumed. Obtain the equation of the line and the R2 value. Using the equation of the line, determine
the actual mass of aluminum metal present in both Reynolds Wrap and Great Value aluminum foil.
Assuming the pressure in the laboratory is 1.00 atm and the temperature is 293 K, calculate the mass
of aluminum in the aluminum foils using the ideal gas law and stoichiometry.