Emperical Formula LH

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Investigating Substances: Synthesis with Copper and Sulfur BACKGROUND Chemists use chemical formulas to show the composition of compounds. A chemical formula shows what elements are contained in a compound, and shows how many of each element are present in one “unit” of the compound. For example, the chemical formula of water, H 2 O, tells us that water contains the elements hydrogen (H) and oxygen (O). The subscript “2” associated with the hydrogen tells us that there are two hydrogen atoms present in one unit of water. The oxygen atom has an assumed subscript of “1”, meaning that there is one oxygen atom present. The “unit” for water is called a molecule. One molecule of water always contains exactly one oxygen atom and two hydrogen atoms. We can also calculate the mass of oxygen and hydrogen in a given quantity of water. It turns out that any sample of water contains about 11.1 % hydrogen and 88.9% oxygen by mass. This can be easily calculated using the atomic masses of hydrogen and oxygen, found on the periodic table. The atomic masses of hydrogen and oxygen are, respectively, 1.0 and 16.0 amu (atomic mass units), meaning that the mass of a water molecule is 18.0 amu (this is called the formula weight or molecular weight). Therefore the mass percent of H in water is % and the mass percent of O in water is %

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Chemistry Lab Handout

Transcript of Emperical Formula LH

Page 1: Emperical Formula LH

Investigating Substances:Synthesis with Copper and Sulfur

BACKGROUND

Chemists use chemical formulas to show the composition of compounds.

A chemical formula shows what elements are contained in a compound, and shows how many of each element are present in one “unit” of the compound.

For example, the chemical formula of water, H2O, tells us that water contains the elements hydrogen (H) and oxygen (O). The subscript “2” associated with the hydrogen tells us that there are two hydrogen atoms present in one unit of water. The oxygen atom has an assumed subscript of “1”, meaning that there is one oxygen atom present. The “unit” for water is called a molecule. One molecule of water always contains exactly one oxygen atom and two hydrogen atoms.

We can also calculate the mass of oxygen and hydrogen in a given quantity of water. It turns out that any sample of water contains about 11.1 % hydrogen and 88.9% oxygen by mass. This can be easily calculated using the atomic masses of hydrogen and oxygen, found on the periodic table. The atomic masses of hydrogen and oxygen are, respectively, 1.0 and 16.0 amu (atomic mass units), meaning that the mass of a water molecule is 18.0 amu (this is called the formula weight or molecular weight). Therefore the mass percent of H in water is

% and the mass percent of O in water is %

Chemical formulas are very useful to chemists. But how are the formulas determined? One way to determine chemical formulas is by finding the mass percent of each element in the compound (which can be measured experimentally). Before we can do this, we need to find a way to count atoms and molecules (since formulas tell us numbers of atoms in a single molecule).

Any sample of a compound used in the laboratory contains an enormous number of atoms or molecules. For example, a 16 oz bottle of water contains approximately 2 x 1025

molecules of water! To avoid working with these large numbers, chemists invented a unit called the mole.

A mole of a substance is a quantity equal in grams to the atomic or molecular weight of a substance. For example, one mole of water is equal to 18.0 g.

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The molecular weight of glucose, C6H12O6, is 6(12.0) + 12(1.0) + 6(16.0) = 180.0 amu. Therefore, one mole of glucose is equal to 180.0 g of glucose.

Interestingly, a mole of a substance always contains a specific number of “units” of the substance, either atoms or molecules. The number of units in a mole of any substance is 6.02 x 1023, called Avogadro’s number.

This means that 18.0 g of water and 180.0 g of glucose both contain the same number of molecules! The difference in mass is because a molecule of glucose is 10 times heavier than a molecule of water.

The formula for a compound gives the ratio of types of atoms, but it also gives the molar ratio of the elements in the compound. For example, water contains two moles of hydrogen for every one mole of oxygen. Glucose contains six moles of carbon to every twelve moles of hydrogen to every six moles of oxygen.

There are two types of chemical formulas that we need to be concerned with: molecular formulas and empirical formulas. The formula shown above for glucose is a molecular formula: it shows the actual numbers of each type of atom in a single molecule of glucose. An empirical formula gives less information: it simply shows the ratio of the types of atoms using the smallest possible whole numbers. For glucose, the empirical formula would be CH2O, telling us that the ratio for C:H:O is 1:2:1. For water, the formula H2O is both a molecular and an empirical formula. Ionic compounds only have empirical formulas because they don’t exist as molecules.

Let’s imagine that you go into the laboratory and determine the percent composition by mass of an unknown compound. You find that the compound consists of 74.5 % chlorine and 25.5 % magnesium. Believe it or not, it is possible to determine the formula for the compound using just this information! But to do so, one needs to find a molar ratio.

The easiest way to solve this is to assume that there is a specific mass of the compound and find the number of grams of each element in that mass. To make things simple, let’s assume there is 100 g of the unknown compound. Then we can write:

74.5 g Cl25.5 g Mg

Now, if we convert each of these to moles, we can easily find the molar ratio:

To find the mole ratio, divide BOTH molar quantities by the smallest (1.05 in this case).

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This gives us a molar ratio of approximately 2 moles of Cl to one mole of Mg. Therefore, we can write our empirical formula as MgCl2. Since this is an ionic compound, only the empirical formula is needed.

OBJECTIVESBy preparing for and performing this experiment, you will:

Synthesize a compound from elements; Write the chemical formula for your product and reaction; Reinforce proper techniques for using an top-loading balance; Become familiar with the locations of lab equipment and lab waste; Use a Bunsen burner; Use a lab notebook to record your results.

MATERIALSChemicals (hood/balance bench):

Copper, Cu (s) (wire, pre-cut)Sulfur, S (s)

Equipment (dispensing room):Crucible with lid

PROCEDURES

1. Check out a crucible from the dispensing room. Clean, dry, and weigh the crucible and cover. Record this mass in your notebook.

2. Place the piece of copper wire into your crucible and weigh the crucible, cover, and copper wire.

3. Record the weight in your notebook and calculate the mass of the copper.

4. Set up the apparatus similar to the picture shown. It should be as far under the ventilation shroud as possible.

5. Weigh approximately 0.5 grams of sulfur and add this to the crucible.

6. Place the crucible with the cover in place as shown in the photo on the pipestem triangle, adjust your burner, and begin heating. (Remember to use tongs to handle the crucible once hot!)

7. Heat gently until the sulfur ceases to burn (blue flame).8. Continue heating for about 5 minutes to dull redness.9. Turn the flame off and allow the crucible to cool for 10 minutes.10. Weigh the crucible, cover, and contents and record the mass.

REPORT

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Make sure to turn your notebook pages in to your instructor before you leave lab. This experiment will require a data sheet. Details are posted on Bb.