Learning Goals

1. Build confidence in yourself as an independent experimentalist;
2. Perform chemical degradation analysis, then use the results to calculate an empirical formula;
3. Use a gravity filtration apparatus;
4. Use a drying oven;
5. Use a Bunsen burner to heat a solid;
6. Observe a redox reaction and a precipitation reaction;
7. Identify sources of uncertainty associated with the procedure and the analytical balance.

Introduction

John Dalton

John Dalton was a chemist, physicist, meterologist, and teacher who lived 1766 – 1844 in England. (Can you believe that he was 34 years old when Middlebury College was founded?!) He was born into a poor Quaker family, so he started supporting himself by working as a teacher at age 12 and continued throughout his life. You can thank Dalton for much of the atomic theory that you are learning in Chem 103.

Over the next two weeks, your laboratory experiments will use Dalton’s atomic theory. In today’s lab, you will determine the molecular formula of a copper compound.

Background

Copper is normally in the 0, +1, or +2 oxidation state. This means copper can combine with 0, 1, or 2 chloride ions. Sometimes water molecules cling to cations and form hydrates. It is difficult to predict the number of water molecules that cling to a salt, but it must always be an integer value.

Today, you will determine the empirical formula of a compound that contains copper, chlorine, and water (Cu_xCl_y zH₂O). You will decompose the compound in order to solve for x, y, and z. The decomposition is in two steps. First, you will heat the compound to dehydrate the compound. Then, you will reduce the copper cations to copper metal to drive off the chlorine. By analyzing how the mass changes after each step, you will be able to solve for the mole ratios in the original compound. A tutorial of this lab is found here.

Procedure

Note: do this lab individually, but share a hood with a “hood mate” and consult with them as needed.

Dry the hydrate

1. 1. Set up a ring stand with a clay triangle on the ring in a fume hood. Place a Bunsen burner underneath with a really small blue flame.
   2. Record the mass of a clean, dry crucible in Table 2.
   3. Place about 1 g of the unknown salt in the crucible. Record the mass of the combined crucible and salt in Table 2. Break up any big chunks with a spatula.
   4. Place your salt and crucible on the clay triangle. Heat super gently. Like giving a baby a warm bath. Stir carefully to prevent burning. Observe the color change from blue-green to brown. If your salt begins to spatter or turn black, lower the heat either by turning down the gas or by raising up your sample from the burner. Don’t hesitate to start over if your salt spatters.
   5. Once the salt is completely brown with no signs of yellow or green, remove the heat and let it cool for about 5 minutes.
   6. Record the mass of the cool crucible containing the dehydrated salt in Table 2. If the mass is not stable, let it cool longer.

Reduce the copper salt to elemental copper

1. Transfer the dehydrated salt to a 50 mL beaker. Rinse the crucible twice with approximately 8 mL of DI H2O each time and pour the rinsate into the same beaker as the dehydrated salt to ensure all of the salt is transferred. Swirl to completely dissolve the salt. Note: the salt turns green when it is dissolved in just a little water, and it turns blue when it is fully hydrated.
2. Coil an approximately 10 cm piece of aluminum wire so that it fits into the beaker. Place it in the beaker and ensure it is completely immersed in the solution. Wait about 30 minutes for the copper to deposit on the surface of the wire. It is finished when the solution is colorless and the bubbling slows. While it is reacting, wash and dry the crucible (but keep yours), set up the next part of the experiment, get a drink of water, work on homework, do your pre-lab for next week (you might want to read through the procedure – it’s a tough one, for sure), go ask Prof. Larrabee questions, call someone in your family, run a 5k, take a nap, etc.
3. Use a spatula to scrape off as much copper as you can from the aluminum wire and put it into the solution.
4. Using that filter paper, set up a funnel filtration apparatus on top of a small Erlenmeyer flask similar to the one shown on the right.
5. Slowly pour the copper and the solution onto the filter paper. Use small amounts of water to rinse all of the copper onto the paper.
6. Open the filter paper, label it, and place it on the watch glass. Put it in the drying oven for about 10 minutes until the copper is dry and crumbly and the paper is dry and begins to brown.
7. When your filter paper starts to brown, carefully transfer all of your copper to your crucible. Put it back in the drying oven for 5 minutes.
8. When you check on your copper, break up the large chunks with a spatula to ensure it is dry throughout. Don’t let it turn black.
9. The drying process is complete when the copper is red/brown and the consistency of Grape Nuts cereal. Let the crucible and copper cool, then weigh it. Record the mass in Table 3. Consult your instructor or TA if the mass of copper is greater than 0.40 g.
10. Pour the filtrate down the drain. Rinse, dry, and return the aluminum wire to where you found it. Save your solid copper!

Oxidize elemental copper to copper oxide

1. Return the copper and crucible to the Bunsen heating set up.

2. Heat the copper and crucible over a hot flame and this time: Burn, baby, BURN. Heat until the copper turns completely black.

3. Let it cool until it is cool enough to handle. Weigh the product. Does the mass increase or decrease? What could cause the mass to increase? What could cause the mass to decrease?

Data analysis

To determine the formula, you will first convert the masses to moles, then divide by the smallest number to get the mole ratios. For example, if you were using a hydrocarbon (composed only of hydrogen and carbon) and you found that it contained 0.25 moles of carbon and 0.99 moles of hydrogen, then you would divide both by 0.25:

0.25 moles carbon / 0.25 = 1 mole carbon

0.99 moles hydrogen / 0.25 = 3.96 moles hydrogen, so approximately 4 moles of hydrogen (remember it must be an integer value!)

Therefore, the empirical formula is CH₄.

Report

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