It is composed of atoms, which in turn are composed of protons, neutrons, and electrons.

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1 Basic Principles of Chemistry Lecture Notes #9 Chemical Composition So far we have learned some basic things about matter: It is composed of atoms, which in turn are composed of protons, neutrons, and electrons. Atoms make up molecules molecular elements and molecular compounds. Matter may come in the form of pure substances or mixtures. Each pure substance is characterized by its own unique set of physical and chemical properties. As chemists, our business is to study the chemical and physical properties of matter. The best way to study matter is by using the scientific method and by making measurements and other observations that we can communicate to others. We are now faced with a dilemma. How can we directly observe and measure atoms and molecules? We face two fundamental problems with trying to approach our studies: 1. Atoms and molecules are so tiny that we are not able to observe them readily without expensive and complex equipment. Recall that we can see atoms roughly with scanning electron microscopes and similar instruments. 2. There are so many atoms and molecules even in something as small as a penny that it would take us almost 100 trillion years to count them all. My office computer has a 1.6 GHz processor, and it would take my computer about 50,000 years to count that high. So how do we solve this problem? The same way we solve some more common everyday problems: Hardware stores don t count nails. Chuck-E-Cheese does not count your tickets when you redeem them. Charlie in the Chemistry stockroom does not count the printer paper he sells to students. Well, not directly anyway. They count these items by weighing them. Can you think of any other examples where this method of counting is used? What needs to be true in order to count things by weighing them? Is this true of atoms and molecules? Recall that when scientists wanted to compare the masses of protons, neutrons, and electrons, they did not have a suitable unit of mass to do this conveniently. So, they devised a new unit of mass that was equal to 1/12 the mass of an atom of 12 C. It was called the atomic mass unit (amu).

2 The amu was an example of an extremely tiny unit of mass created to report the masses of extremely tiny objects. Since atoms come in extremely large numbers, scientists also need an extremely large counting unit to report the number of atoms they have observed. What are some counting units you are familiar with from everyday life? What do these units have in common? Before we spend a lot of time learning about a new counting unit for atoms and molecules, think about these questions: Why is it so important for us to account for all of the atoms and molecules in a sample? Do you think we can account for all of the atoms and molecules in a sample? If not, what will limit our ability to do that? Our New Unit: The Mole This unit, like the amu, is based on the 12 C atom. The counting unit could have any arbitrary value under the sun, provided it was huge enough to give us reasonably manageable numbers of atoms. But remember that we are dealing with two problems here the enormity of the number of atoms that we are trying to count and their extremely tiny size. So keep in mind that we are still going to have to count these items by weighing them. If one carbon atom weighs 12 amu, wouldn t it be oh so convenient to devise a counting unit that could describe a collection of carbon atoms that weighed precisely 12 of some other mass unit particularly a unit that we can weigh on our laboratory balances? That s exactly what was done. 1 atom of 12 C weighs 12 amu. 1 MOLE of 12 C weighs 12 grams. Consider the following problem: We learned in class that the atomic mass of an atom is the total of all the protons and neutrons in the atom. The atomic mass of carbon (C) is 12.0 amu. This means that one carbon atom weighs 12.0 atomic mass units. One atomic mass unit is equal to 1.66 x kilograms. a. Fill in the numerical values in the blanks below to complete the following conversion factors: 1 C atom 1 amu 1 kg amu kg 1000 g (exact) b. Using these three conversion factors, calculate the number of carbon atoms that would be needed in order to have 12.0 grams of carbon atoms. In other words, use these factors to convert 12.0 grams into units of C atoms. Report

3 your answer to the correct number of significant figures g 1 kg 1 amu 1 C atom X X 1000 g 1.66x10-27 X = 6.02x10 23 C atoms kg 12.0 amu So, one mole of atoms (or molecules or toothpicks or marbles) = 6.02 x of those items. This extremely large number is called Avogadro s number. Let s put this number in perspective: 1 tablespoon of water contains about one mole of water (H 2 O) molecules. 22 pennies (pre-1982) contains about one mole of copper (Cu) atoms. Two 12 helium balloons contain about one mole of helium (He) atoms. This all sounds very reasonable until you begin to consider just how unfathomably large Avogadro s number is Some illustrations are: One mole of marbles would be enough to cover the surface of Earth to a depth of 50 miles. One mole of pennies would be enough to give every man, women, and child in the world $1,000,000,000,000! A glass of water contains about 10 moles of water. This number of water molecules is greater than the number of grains of sand in the Sahara Desert. An extra credit assignment: Join in a group of three to six students and devise a creative way to illustrate the immensity of Avogadro s number. The Bulletin Board is a good way to find a group and begin communication with your group. Your example must be justified by mathematical calculations. Create a "web page" to display your illustration and your computations. Use appropriate pictures and graphics to get your point across. Submit a draft of your project to me by Sunday February 26, 2006 for my approval and suggestions. Have your final page ready to turn in on Friday March 3. I will award up to 10 bonus points for your efforts. Save your "web page" in Microsoft Word or another document format, and designate one of your group members to it to me by the due date. I will post them on our course website for everyone to view. Be sure that every group member has his/her name on the project so that I can assign proper credit.

4 So, one mole of atoms (or molecules or toothpicks or marbles) = 6.02 x of those items. Now, let s put our new counting unit to work If we take any element in the periodic table, we can tell how much one mole of that element weighs by taking the mass number (or atomic mass) and changing the units from amu to grams. Atoms and Moles How many carbon (C) atoms are present in one mole of C atoms? How many C atoms are present in three moles of C atoms? How many C atoms are present in ¼ mole of C atoms? Grams and Moles How much does one mole of carbon (C) weigh? How much does one mole of helium (He) weigh? How much do two moles of helium (He) weigh? How much does ½ mole of helium weigh? The mass of one mole of an element is called that element s MOLAR MASS. Grams and Atoms How many atoms are present in 12.0 grams of carbon (C)? How many atoms are present in 2.0 grams of helium (He)? Molecules and Moles How many carbon dioxide (CO 2 ) are present in 0.35 moles of CO 2? Grams and Moles How many grams of carbon dioxide are needed to make 2 moles of carbon dioxide? How many moles of sucrose (C 12 H 22 O 11 ) are needed to make 15.0 grams of sucrose? Grams and Molecules How many molecules are present in 13.2 grams of carbon monoxide (CO)? Moles of a Compound and Moles of an Element How many moles of hydrogen (H) are found in 3.0 moles of water (H 2 O)?

5 How many moles of oxygen (O) are found in 2.38 moles of calcium carbonate (CaCO 3 )? Grams of a Compound and Grams of an Element How many grams of bromine (Br) are present in 34.0 grams of dark copper bromide (CuBr 2 )? How many grams of carbon are present in 4.56 moles of ethyl alcohol (C 2 H 6 O)? Mass Percent Composition of Compounds What is the mass percent of copper (Cu) in copper carbonate (CuCO 3 )? What is the mass percent of oxygen (O) in copper carbonate (CuCO 3 )? Determining the Empirical Formula of a Compound A sample of white powder seized during a drug raid was analyzed to determine its elemental composition. It was found to contain 80.48% carbon, 10.13% hydrogen, and 9.39% nitrogen by mass. What is the empirical formula of the compound?