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1 STOICHIOMETRY LEARNING OUTCOMES At the end of this unit students will be expected to: UNIT 1 THE MOLE AND MOLAR MASS define molar mass and perform mole-mass inter-conversions for pure substances explain how a major scientific milestone, the mole, changed chemistry CALCULATIONS AND CHEMICAL EQUATIONS identify mole ratios of reactants and products from balanced chemical equations identify practical problems that involve technology where equations were used state a prediction and a hypothesis based on available evidence and background information perform stoichiometric calculations related to chemical equations STOICHIOMETRIC EXPERIMENTATION design stoichiometric experiments identifying and controlling major variables use instruments effectively and accurately for collecting data identify and explain sources of error and uncertainty in measurement using precision and accuracy communicate questions, ideas, and intentions, and receive, interpret, understand, support, and respond to the ideas of others identify various constraints that result in trade-offs during the development and improvement of technologies APPLICATIONS OF STOICHIOMETRY identify various stoichiometric applications predict how the yield of a particular chemical process can be maximized explain how data support or refute the hypotheses or prediction of chemical reactions compare processes used in science with those used in technology analyse society's influence on science and technology

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3 UNIT 1 AN INTRODUCTION TO STOICHIOMETRY LESSON 1 How does a chemist count out atoms of the element sodium, Na In industry, when making extremely large amounts of chemicals through chemical reactions, how do chemists know how much of the reactants to react or how much product will be formed This unit will answer these questions and other questions related to amount of matter. The word stoichiometry comes from the Greek words, stoicheion (meaning any first thing or principle) and metron (meaning measure). Stoichiometry deals with the mass-mass or molemole relationship among reactants and products in a balanced chemical equation. It answers questions like how much of one reactant will react with a given amount of another reactant and how much product will be formed. Whether a compound is being prepared by the chemical industry or in a high school laboratory, it deals with extremely large numbers of atoms, molecules, or ions. For example, if we were to prepare the flavouring agent, ethyl butyrate, we might use 10.0 g of butyric acid. This relatively small mass of butyric acid contains molecules, a huge number. If we had an imaginary device that could count molecules at the rate of one million per second, it would take over two billion years for the device to count that many molecules. To overcome the problem of dealing with extremely large numbers of particles, chemists use the mole concept. A mole is an amount (quantity) of matter that contains particles of that matter. Depending on the make-up of matter, the particles could be atoms, molecules, ions, or formula units. The symbol for the mole unit is mol. The mole is one of the seven base units of the standard SI system of measurement. German chemist, Wilhelm Ostwald, introduced the mole unit around The mole is used by chemists to communicate the amount of matter they are using. The number is called Avogadro's number, after the Italian scientist, Amedeo Avogadro, ( ). Another definition of the mole is: a mole of a substance is the quantity of that substance that contains the same number of chemical units (atoms, molecules, formula units, or ions) as there are atoms in exactly 12 g of the isotope carbon-12. The mole is thus a unit of measurement. A certain number of moles of a substance contain a certain number of grams and a certain number of particles. National Mole Day is celebrated yearly across North American high schools on October 23 rd, starting at 6:02 a.m. The mole was a major milestone in chemistry; chemists could quantify their observations and use it to make quantitative predictions. The mole concept makes it extremely simple to convert from numbers of atoms or molecules or ions (in moles) to mass in grams, and vice versa. Avogadro's number provides a way of converting between the number of particles on the micro scale (which we cannot see) and the number of moles on the macro scale (which we can see). CORRESPONDENCE STUDY PROGRAM PAGE 17

4 CHEMISTRY 11 TEXT READING 4.3 Balancing Chemical Equations and Translating Balanced Chemical Equations on page 138. LESSON EXERCISES 1. Explain what is meant by stoichiometry in your own words. 2. Explain what is meant by the mole unit of measurement in your own words. 3. Why is it important to be able to use the mole unit of measurement 4. What is the value of Avogadro's number 5. How many molecules are there in two moles of water JOURNAL ENTRY PAGE 18 CORRESPONDENCE STUDY PROGRAM

5 UNIT 1 CALCULATING AVERAGE ATOMIC MASSES LESSON 2 The mass of an atom is the sum of all of the electrons, protons, and neutrons it contains. The standard for atomic mass is the isotope carbon-12. It has a mass of exactly u. Isotopes are atoms of the same element that have different masses because of different numbers of neutrons. One atomic mass unit, 1 u, is 1/12 the mass of a carbon- 12 atom. One atomic mass unit, 1 u, is equivalent to g. It is easier to use atomic mass units because the mass of an atom in grams, g, is extremely small. The mass of all element isotopes is compared to the carbon-12 atom. SAMPLE PROBLEM 1 What is the mass of an atom with a mass 7/6 times greater than the carbon-12 atom The mass of an atom 7/6 times bigger than a carbon-12 atom is 7/ u = u. SAMPLE PROBLEM 2 What is the mass of an atom 4/13 times the mass of a chromium atom, if a chromium atom is 13/3 times the mass of a carbon-12 atom. The mass of the atom is: 4/13 13/ u = u. CORRESPONDENCE STUDY PROGRAM If you looked up the atomic mass of an element (the mass of a single atom), you would find that atomic masses are not whole number integers. The atomic masses used in this course are located in the Chemistry Data Booklet. Where do these masses come from Examine hydrogen; this element has three different isotopes, 1 1H, 2 1H, and 3 1H with masses of approximately 1.0 u, 2.0 u, and 3.0 u, respectively. If we took an average of these three isotopes, we would get an average of 2.0 u for the hydrogen atom. Looking up the atomic mass of hydrogen, we see that it is 1.01 u. Why the big difference The atomic mass of an element is the weighted average of all of the naturally occurring isotopes of an element. The weighted average takes into consideration just how much there is of each isotope in nature. In the case of hydrogen, per cent of all of the hydrogen atoms found in nature are the lightest isotope of hydrogen, 1 1H, with an approximate mass of 1.0 u. So the weighted average is close to 1.0 u. SAMPLE PROBLEM 3 Natural neon consists of per cent 20 10Ne, actual mass u; per cent 21 10Ne, actual mass u; and 8.82 per cent 22 10Ne, actual mass u. Calculate the average atomic mass of natural neon. Assume that we have a sample of neon atoms. Calculate the average by finding the sum of the masses of all of the atoms and then dividing by If we wanted to find the average weight of a student in a classroom, we PAGE 19

6 CHEMISTRY 11 would add up the weights of all of the students in the room and then divide by the number of students in the classroom. The average atomic mass of natural neon is found in a similar manner: First calculate the amount of each isotope in atoms % = Ne % = Ne % = Ne Then calculate the total atomic mass for each isotope, add the masses together, and divide by ( u u u) = 20.2 u. SAMPLE PROBLEM 4 Natural nitrogen consists of per cent 14 7 N, actual mass u; and 0.37 per cent 15 7 N, actual mass u. Calculate the average atomic mass of natural nitrogen for a sample size of LESSON EXERCISES 1. Natural magnesium consists of per cent 24 Mg, actual mass u; per cent 25 Mg, actual mass u; 12 and per cent Mg, actual mass u. Calculate the average atomic mass of natural magnesium. 2. Natural silicon consists of per cent 28Si, actual mass u; 4.70 per cent 14 29Si, actual mass u; and 3.09 per 14 Si, actual mass u. cent Calculate the average atomic mass of natural silicon. 3. Natural argon consists of per cent 36Ar, actual mass u; per 18 Ar, actual mass u; and cent per cent Ar, actual mass u. Calculate the average atomic mass of natural argon. JOURNAL ENTRY First calculate the amount of each isotope in atoms % = N % = N Then calculate the total atomic mass for each isotope, add the masses together, and divide by The average atomic mass of natural nitrogen is: u u = u PAGE 20 CORRESPONDENCE STUDY PROGRAM

7 UNIT 1 DETERMINING ATOMIC MASS, MOLECULAR MASS, FORMULA MASS, AND MOLAR MASS LESSON 3 Atomic mass is the mass of an atom of an element in atomic mass units, u. This value is given in the periodic table, or international table of atomic masses, or in the Chemistry Data Booklet. For example, the atomic mass of an atom of the element sodium, Na, is u. The atomic mass of an atom of the element xenon, Xe, is u. Molecular mass is the mass of a molecule of a compound or an element, in atomic mass units, u. It is the sum of all of the atomic masses of all of the atoms in the molecule. For example, the molecular mass of a molecule of water, H 2 O, is the sum of: 2H = u = 2.02 u and 1O = u = u. The molecular mass of water is u. The molecular mass of F 2 is the sum of the atomic masses of F and F: u u = u. Not all compounds are made up of molecules; which result from covalent bonding. Some compounds form by ionic bonding and are made up of ions. There are no molecules present so the compounds cannot be defined using molecular mass. These compounds consist of formula units instead of molecules. A formula unit is the smallest unit of an ionic compound that represents the properties of the ionic compound. It is the CORRESPONDENCE STUDY PROGRAM simplest whole number ratio of positive and negative ions in the ionic compound. A formula unit will have the same formula as the formula for the ionic compound. Formula mass is the mass of a formula unit. A formula unit of table salt, NaCl (sodium chloride) is NaCl. The formula mass of the ionic compound NaCl is the sum of the atomic masses of the ions in the formula unit. This is: 1Na + = u + 1Cl - = u = u. Molar mass is the mass of one mole of atoms, molecules, ions, or formula units in grams. It is the atomic mass, molecular mass, or formula mass expressed in grams instead of atomic mass units. The molar mass is numerically equal to the atomic mass, molecular mass, or formula mass, but in grams. Looking at the chemicals we have referred to in this lesson, we know the molar mass of Na is g, and the molar mass of Xe is g, the molar mass of H 2 O is g, the molar mass of F 2 is g and the molar mass of NaCl is g. To determine the molar mass of a substance, look up the atomic mass, or molecular mass, or formula mass, and drop the u unit and replace it with the g unit. Does this mean that the atomic mass of Na, u is the same as the molar mass of Na, g No, only the numbers are the same, the amounts are completely different because of different units. For example, the atomic mass of Na is the mass of a single Na atom. One atomic PAGE 21

8 CHEMISTRY 11 mass unit = g. The amount in grams is: g = g. The molar mass of Na is the mass of a mole of Na atoms, Na atoms. This would be g. As you can see, g is much smaller than g. The table below summarizes the relationship between molar mass and atomic mass, molecular mass, and formula mass. substance Na Xe H 2 O F 2 NaCl particles atoms atoms molecules molecules formula units atomic mass molecular mass formula mass molar mass u g u g 3H = u = 3.03 u u u = u The molecular mass of CH 3 OH is: 1C = u = u 4H = U = 4.04 u 1O = u = u u u u = u u g u g u g SAMPLE PROBLEM 3 What are the formula masses of K 2 SO 4 and (NH 4 ) 3 PO 4 The formula mass of K 2 SO 4 is: 2K = u = u 1S = u = u 4O = u = u SAMPLE PROBLEM 1 What are the atomic masses of Mn and K The atomic mass of Mn is u and the atomic mass of K is u. SAMPLE PROBLEM 2 What are the molecular masses of NH 3 and CH 3 OH The molecular mass of NH 3 is: 1N = u = u Formula mass = u The formula mass of (NH 4 ) 3 PO 4 is: 3N = u = u 12H = u = u 1P = u = u 4O = u = u Formula mass = u SAMPLE PROBLEM 4 What is the molar mass of Al, C 2 H 6, and Mg(OH) 2 PAGE 22 CORRESPONDENCE STUDY PROGRAM

9 UNIT 1 The atomic mass of Al is u. The molar mass of Al is therefore g. The molecular mass of C 2 H 6 is u. The molar mass of C 2 H 6 is therefore g. The formula mass of Mg(OH) 2 is u. The molar mass of Mg(OH) 2 is therefore g. TEXT READING Molar Mass on page 143. LESSON EXERCISES 1. What is the atomic mass of Li, Ag, and Kr 2. What is the molecular mass of CCl 2 F 2, C 2 H 5 OH, and NBr 3 3. What is the formula mass of Na 2 C, (NH 4 ) 2 SO 4, and Ca 3 (PO 4 ) 2 4. What is the molar mass of Ne, SiCl 4, K 3 PO 4, Mg, Sr(N ) 2, and CH 3 Br JOURNAL ENTRY CORRESPONDENCE STUDY PROGRAM PAGE 23

10 CHEMISTRY 11 MOLE CONVERSIONS LESSON 4 A mole is an amount or quantity of matter that contains particles of that matter. Depending upon the make-up of matter, the particles could be atoms, molecules, ions, or formula units. There are different ways of expressing the amount of matter you have, such as grams or moles. By knowing the number of moles or grams you have, you can calculate how many particles you have. Otherwise, even if you could count individual particles, you could not count the extremely large number of particles in a sample of matter. For a given sample of matter, moles, grams, and number of particles are all interrelated. If you know one of the three (moles, grams, or number of particles) you can calculate how much you have of the other two. These conversions are known as mole conversions. The three quantities are found in the mole triangle. moles grams number of particles Use the factor label method or dimensional analysis when doing mole conversions. In this method, you write down and keep track of the units or dimensions. You multiply a given quantity by a ratio (conversion factor), which converts the given quantity into the quantity you require. There is a relationship between the two quantities in the ratio. You treat the same units as factors that you would cancel in the same way that numbers are 3 cancelled in algebra. For example, 2 2 = 3. The 2s are identical numbers that we can cancel out. Likewise, we can cancel similar units by treating them as similar numbers. For example, g mol/g = mol. SAMPLE PROBLEM 1 How many mol and atoms are there in a 22.3 g sample of copper, Cu Multiply 22.3 g of Cu by a conversion factor (ratio) that will change it into mol g mol Cu g Cu What is the relationship between # mol of Cu and # g of Cu Choose a relationship that is simple to figure out. Assume there is 1 mol of Cu. Now determine how many g of Cu there are in 1 mol of Cu. The molar mass of Cu is g. Therefore, the ratio is: 1 mol Cu g Cu We can now complete the calculation 22.3 g Cu 1 mol Cu = mol Cu g Cu The number of atoms in 22.3 g of Cu is 22.3 g Cu atoms Cu/g Cu. Now determine the relationship between # atoms of Cu and # g of Cu. Use Avogadro's number, , for the number of Cu atoms. This is also the number of Cu atoms in 1 mol of Cu. Put the molar mass of Cu on the bottom of the ratio to give: PAGE 24 CORRESPONDENCE STUDY PROGRAM

11 UNIT atoms Cu g Cu Now complete the calculation: 22.3 g Cu atoms Cu g Cu = atoms Cu. SAMPLE PROBLEM 2 How many grams and atoms are there in mol of potassium, K units. The method of doing these mole conversions is similar to the method used when working with elements and atoms. SAMPLE PROBLEM 4 How many mol and molecules are there in 25.0 g of water, H 2 O 25.0 g H 2 O 1 mol H 2 O g H 2 O = 1.39 mol H 2 O mol K g K = 29.3 g K 1 mol K mol K atoms K 1 mol K = atoms K SAMPLE PROBLEM 3 How many grams and mol are there in atoms Mg atoms Mg g Mg atoms Mg = 10.3 g Mg g H 2 O molecules H 2 O g H 2 O = molecules H 2 O SAMPLE PROBLEM 5 How many g and formula units are there in 1.15 mol of AgN 1.15 mol AgN g AgN 1 mol AgN = 195 g AgN 1.15 mol AgN formula units AgN 1 mol AgN atoms Mg 1 mol Mg atoms Mg = mol Mg So far, we have restricted our calculations to those involving elements. Now look at mole conversions involving compounds with molecules and formula = formula units AgN SAMPLE PROBLEM 6 How many mol and g are there in molecules of NH 3 CORRESPONDENCE STUDY PROGRAM PAGE 25

12 CHEMISTRY molecules NH 3 1 mol NH molecules NH 3 = mol NH molecules NH g NH molecules NH 3 = 1.27 g NH 3 SAMPLE PROBLEM 7 How many atoms are there in 42.5 g of CF g CF molecules CF g CF 4 = molecules CF 4 5 atoms/1molecule = atoms TEXT READING Mass-Amount Conversions on pages LESSON EXERCISES 1. How many g and atoms are present in mol of barium, Ba 2. How many mol and atoms are there in g of iron, Fe 3. How many mol and g are there in atoms of rubidium, Rb 4. How many g and molecules are present in 2.76 mol of propane, C 3 H 8 5. How many mol and formula units are present in g of CaSO 4 6. How many mol and g are present in formula units of KCl 7. How many atoms are present in g of C 2 H 6 8. How many formula units are there in g of LiF 9. Which is the largest mass in g, 25.2 g of CaBr 2, 1.80 mol of CH 4, or molecules of NCl Why isn't the atomic mass of potassium, K, the same as the molar mass of potassium JOURNAL ENTRY PAGE 26 CORRESPONDENCE STUDY PROGRAM

13 UNIT 1 CALCULATING THE PERCENTAGE COMPOSITION OF A COMPOUND LESSON 5 Quite often if a compound is being identified, its percentage composition is calculated. This involves calculating the percentage by mass of each element in the compound. To calculate the percentage composition for an element in a compound we need to know the total mass of the element in the compound and the total mass of the compound. The percentage of the element in the compound is the total mass of the element in the compound divided by the total mass of the compound. This ratio is then multiplied by 100. SAMPLE PROBLEM 1 What is the percentage composition of propane gas, C 3 H 8, used in barbecues The molecular mass of C 3 H 8 is u. %C = total mass of C in compound 100 mass of compound There are three C atoms in C 3 H 8 so the total mass of C is u = u. %C = u 100 = 81.68% u %H = total mass of H in compound 100 mass of compound There are eight H atoms in C 3 H 8 so the total mass of H is u = 8.08 u. %H = 8.08 u 100 = 18.3% u SAMPLE PROBLEM 2 What is the percentage composition of Ca 3 (PO 4 ) 2 The formula mass of Ca 3 (PO 4 ) 2 is u. %Ca = total mass of Ca in compound 100 mass of compound There are three Ca atoms in Ca 3 (PO 4 ) 2 so the total mass of Ca is u = u. %Ca = u 100 = % u %P = total mass of P in compound 100 mass of compound There are two P atoms in Ca 3 (PO 4 ) 2 so the total mass of P is u = u. %P = u 100 = 19.97% u %O = total mass of O in compound 100 mass of compound There are eight O atoms in Ca 3 (PO 4 ) 2 so the total mass of O is u = u. %O = u 100 = % u SAMPLE PROBLEM 3 Calculate the percentage composition of CH 3 COOH. The molecular mass of CH 3 COOH is u. %C = total mass of C in compound 100 mass of compound CORRESPONDENCE STUDY PROGRAM PAGE 27

14 CHEMISTRY 11 There are two C atoms in CH 3 COOH so the total mass of C is u = u. %C = u 100 = 40.01% u %H = total mass of H in compound 100 mass of compound There are four H atoms in CH 3 COOH so the total mass of H is u = 4.04 u. %H = 4.04 u 100 = 6.73% u %O = total mass of O in compound 100 mass of compound There are two O atoms in CH 3 COOH so the total mass of O is u = u. %O = u 100 = 53.26% u SAMPLE PROBLEM 4 Calculate the percentage of O in CaC. In this question we are only calculating the percentage by mass of a single element in the compound, oxygen. The formula mass of CaC is u. %O = total mass of O in compound 100 mass of compound There are three O atoms in CaC so the total mass of O is u = u. is not, then you know there is a mistake in the solution to the problem. Also, percentage composition problems ask you to calculate the percentage by mass of each element in the compound. So, in sample problem 2, we did not calculate the percentage of PO 4 because PO 4 is not an element. We needed to calculate the percentage of P and O. LESSON EXERCISES 1. Calculate the percentage composition of one of the components of gasoline, octane, C 8 H Calculate the percentage composition of the antacid, milk of magnesia, Mg(OH) Calculate the percentage composition of baking soda, sodium hydrogen carbonate, NaHC. 4. Calculate the percentage composition of the acid found in car batteries, sulphuric acid, H 2 SO Calculate the percentage composition of the explosive, TNT, trinitrotoluene, C 7 H 5 (NO 2 ) What is the percentage of N in smelling salts, (NH 4 ) 2 C JOURNAL ENTRY %O = u 100 = 47.94% u In percentage composition problems, the sum of all of the percentages of all of the elements should be 100 per cent or very close to 100 per cent. If it PAGE 28 CORRESPONDENCE STUDY PROGRAM

15 EMPIRICAL FORMULA LESSON 6 Empirical formula is the simplest formula for writing a compound. It is the simplest whole number ratio of atoms or ions in the formula for the compound. It does not tell us the actual number of atoms of an element in a molecule of the compound. Determining the empirical formula for a compound can be used to identify the compound. To calculate the empirical formula of a compound we must know either the masses of all of the elements in the compound or the percentage composition by mass of the compound. SAMPLE PROBLEM 1 What is the empirical formula for a compound which on analysis shows 2.2% H, 26.7% C, and 71.1% O UNIT 1 numbers represent the subscripts of the elements in the empirical formula and we have calculated the empirical formula. If a mole ratio is not close to a whole number, then multiply the ratio by the smallest whole number to make the ratio a whole number. Then multiply all of the mole ratios by this whole number to give the subscripts for the empirical formula. These steps are summarized in the table below. H C O 2.2 g 26.7 g 71.1 g 2.2 g H 1 mol H 1.01 g H = 2.2 mol H 2.2 mol 2.2 mol = g C 1 mol C g C = 2.22 mol C 2.22 mol 2.2 mol = g O 1 mol O g O = 4.45 mol O 4.45 mol 2.2 mol = The empirical formula is H 1 C 1 O 2 or HCO 2. We are trying to find the subscripts X, Y, and Z in the empirical formula for the compound, H x C y O z. Use the following steps: Assume you have 100 g of the compound. Since per cent means out of 100, we will know how many g we have of each element. Change % to g Change g to mol Divide the smallest number of moles into itself and all of the other number of moles. If the results of the above step are whole numbers for each element or very close to whole numbers for each element, then these CORRESPONDENCE STUDY PROGRAM SAMPLE PROBLEM 2 What is the empirical formula of a compound which on analysis shows 69.9% Fe and 30.1% O Fe O 69.9 g 30.1 g 69.9 g Fe 1 mol Fe 55.8 g Fe = 1.25 mol Fe 1.25 mol 1.25 mol = g O 1 mol O g O = 1.88 mol O 1.88 mol 1.25 mol = = = 3.00 PAGE 29

16 CHEMISTRY 11 Since we did not have all whole numbers or numbers very close to whole numbers, multiply 1.50 by the smallest whole number that will make 1.50 a whole number. In this case, multiply by two. Then multiply all other ratios by two. Therefore, 1.00 becomes The empirical formula of the compound is Fe 2. SAMPLE PROBLEM 3 When g of antimony, Sb, is heated with an excessive amount of sulphur, S, a chemical reaction takes place. The antimony joins with some of the sulphur and the excess sulphur escapes, leaving only the compound. If g of the compound was made, what is the empirical formula of the compound This problem is worded differently than the two previous problems. We are not given percentages of each element in the compound. However, if we know the mass of each element in the compound, we can determine the empirical formula by changing g to mol and following the rest of the steps. We know the mass of the antimony, Sb. It is g. The compound contains only Sb and S. We know the mass of the compound is g. Sb g g g Sb 1 mol Sb g Sb = mol Sb mol mol = S g S 1 mol S g S = mol S mol mol = The mass of S present in the compound is g g = g S. The empirical formula is Sb 2 S 3. TEXT READING Molecular Compounds and Empirical Molecular Formulas on pages LESSON EXERCISES 1. Determine the empirical formula of a compound which on analysis shows: 38.67% C; 16.23% H; and 45.10% N. 2. Determine the empirical formula of a compound which on analysis shows: 71.65% Cl; 24.27% C; and 4.07% H. 3. A g sample of phosphorous was reacted with enough oxygen to form a white compound made up of only P and O and having a mass after reaction of g. Determine the empirical formula of this compound. 4. Determine the empirical formula of a compound with the following percentages: Fe = 63.53% and S = 36.47%. 5. Determine the empirical formula of a compound which on analysis shows: 26.52% Cr; 24.52% S; and 48.96% O. 6. Determine the empirical formula of a compound with the following percentages: 63.1% C; 11.92% H; and 24.97% F. JOURNAL ENTRY = = PAGE 30 CORRESPONDENCE STUDY PROGRAM

17 UNIT 1 DETERMINING MOLECULAR FORMULA LESSON 7 The empirical formula gives us the minimum information about a compound. It just shows the simplest whole-number ratio of atoms of the elements in a compound. A formula that provides more information about a compound is the molecular formula. The molecular formula shows the actual number of atoms of each element joined together to form a single molecule of the compound. More information is provided by a structural formula for a compound. A structural formula shows exactly which atoms are joined together in a molecule. The diagram below summarizes information provided by an empirical formula, a molecular formula, and a structural formula. Hydrogen Peroxide empirical formula HO molecular formula H 2 O 2 structural formula CORRESPONDENCE STUDY PROGRAM O H H O To determine the molecular formula for a compound we need to know its empirical formula and its molecular mass. The molecular mass is some whole number multiple of the mass of the empirical formula. We can determine the molecular formula by using the following equation: (mass of empirical formula)x = molecular mass (the total mass of the empirical formula that was given or determined)(some whole number) = (total mass of the molecule) We then multiply each subscript in the empirical formula by the whole number x to give the molecular formula. SAMPLE PROBLEM 1 A certain compound has an empirical formula of CH 2 O and a molecular mass of 180 u. What is the molecular formula of the compound (mass of empirical formula)x = molecular mass ( u u u)x = 180 u (30.02 u)x = 180 u x = 6.00 Since the empirical formula is CH 2 O, multiply each subscript in the empirical formula by 6.00 to give the molecular formula C 6 H 12 O 6. SAMPLE PROBLEM 2 The empirical formula for a compound is P 2 O 5 and the molecular mass of the compound is u. What is the molecular formula of the compound (mass of empirical formula)x = molecular mass ( u u)x = u ( u)x = u x = PAGE 31

18 CHEMISTRY 11 Since the empirical formula is P 2 O 5, we multiply each subscript in the empirical formula by 2 to give the molecular formula P 4 O 10. SAMPLE PROBLEM 3 A certain compound was found to have 21.9% Na; 45.7% C; 1.9% H; and 30.5% O. What is the molecular formula for this compound if its molecular mass is 210 u First determine the empirical formula. Na C H O 21.9 g 45.7 g 1.9 g 30.5 g mol 3.80 mol 1.9 mol 1.91 mol mol mol = mol mol = mol mol = mol mol = The empirical formula is NaC 4 H 2 O 2. (mass of empirical formula)x = molecular mass = u ( u)x = 210 u x = 2.00 The molecular formula is Na 2 C 8 H 4 O 4. LESSON EXERCISES 1. The empirical formula of a certain compound is CH. It has a molecular mass of 52 u. What is its molecular formula 2. A certain compound was analysed and found to have 32.0% C; 6.7% H; and 18.7% N; with the rest of the compound made up of oxygen, O. The molecular mass is 75.0 u. What is the molecular formula of the compound 3. What is the molecular formula for a compound whose molecular mass is 116 u and whose empirical formula is CHO 4. What is the molecular formula for a compound whose molecular mass is 90 u and whose empirical formula is CH 2 O 5. A certain compound was found to contain 48.48% C; 5.05% H; 14.14% N; and 32.32% O. The molecular mass is 198 u. What is its molecular formula 6. A certain compound was found to have the following percentage composition: 49.38% C; 3.55% H; 9.40% O; and 37.67% S. The molecular mass is u. What is its molecular formula JOURNAL ENTRY PAGE 32 CORRESPONDENCE STUDY PROGRAM

19 STOICHIOMETRY CALCULATIONS LESSON 8 A balanced chemical equation provides a lot of information. It tells what the reactants are in a chemical reaction and what the products are. A balanced chemical equation obeys the law of conservation of mass. The total mass of all reactants before a chemical reaction equals the total mass of products after the chemical reaction takes place. This concept leads us to the mole-mole or mass-mass relationship in a balanced chemical equation, known as stoichiometry. Examine this relationship through the following balanced chemical equation: C (s) + O 2(g) CO 2(g) From this equation we can infer that one atom of C reacts with one molecule of O 2 to produce one molecule of CO 2. From a mass relationship, 12 u of C would react with 32 u of O 2 to produce 44 u of CO 2. The law of conservation of mass is preserved. We could also say that 10 atoms of C reacts with 10 molecules of O 2 to make 10 molecules of CO 2. Again, from a mass relationship, 120 u of C reacts with 320 u of O 2 to produce 440 u of CO 2. We could also say that atoms of C reacts with molecules of O 2 to produce molecules of CO 2. This next relationship is very important. Based on the last statement and since there are particles in one mol of a substance, we can say that one mol of C reacts with one mol of O 2 to make CORRESPONDENCE STUDY PROGRAM UNIT 1 one mol of CO 2. That is, 12 g of C reacts with 32 g of O 2 to produce 44 g of CO 2. This is the molemole and mass-mass relationship in a balanced chemical equation. It is the stoichiometry of a chemical reaction. This information is summarized in the table below: C (s) + O 2(g) CO 2(g) 1 atom 1 molecule 1 molecule 12 u 32 u 44 u 10 atoms 10 molecules 10 molecules 120 u 320 u 440 u atom 2 s molecules molecules 1 mol 1 mol 1 mol 12 g 32 g 44 g Examine some calculations involving balanced chemical equations using the factor-label method or dimensional analysis to solve the problems. This is the same method used for mole conversions. SAMPLE PROBLEM 1 Given the reaction: C (s) + O 2(g) CO 2(g) a) How many g of O 2 reacts with 9.8 g of C b) How many g of CO 2 are produced when 19.6 g of C is completely reacted with enough O 2 c) How many mol of CO 2 are produced from reacting 0.85 mol of O 2 d) How many g of C are needed to react with 1.45 mol of O 2 e) How many molecules of O 2 are needed to react with 5.26 g of C PAGE 33

20 CHEMISTRY 11 f) How many molecules of CO 2 are made by reacting atoms of C a) C (s) + O 2(g) CO 2(g) 1 mol 1 mol g g 9.8 g g 9.80 g C g O 2 = 26.1 g O g C b) C (s) + O 2(g) CO 2(g) 1 mol 1 mol g g 19.6 g g 19.6 g C g CO 2 = 71.8 g CO g C c) C (s) + O 2(g) CO 2(g) 1 mol 1 mol 0.85 mol mol 0.85 mol O 2 1mol CO 2 = 0.85 mol CO2 1mol O 2 d) C (s) + O 2(g) CO 2(g) 1 mol 1 mol g g g 1.45 mol 1.45 mol O g C = 17.4 g C 1 mol O 2 e) C (s) + O 2(g) CO 2(g) 1 mol 1 mol g molecules 5.26 g molecules 5.26 g C molecules O 2 = molecules O g C f) C (s) + O 2(g) CO 2(g) 1 mol 1 mol atoms molecules atoms molecules atoms C molecules CO atoms C = molecules CO 2 The sample problem we just solved involved a stoichiometry in which one mol reacted with one mol to produce one mol. Not all chemical reactions have this simple one to one stoichiometry. Examine some problems where the stoichiometry is not as simple. These problems are solved using a similar method to sample 1 problems. SAMPLE PROBLEM 2 Given the reaction: Fe 2 + 3CO 2Fe + 3CO 2 a) How many g of Fe 2 reacts with g of CO b) How many g of CO 2 are produced from reacting 45.7 g of Fe 2 with sufficient CO c) How many mol of Fe are produced from reacting 1.80 mol of CO d) How many g of Fe 2 are needed to react with 6.50 mol of CO e) How many molecules of CO are needed to react with g of Fe 2 f) How many atoms of Fe can be made by reacting formula units of Fe 2 PAGE 34 CORRESPONDENCE STUDY PROGRAM

21 UNIT 1 a) Fe 2 + 3CO 2Fe + 3CO 2 1 mol 3 mol g g g g e) Fe 2 + 3CO 2Fe + 3CO 2 1 mol 3 mol g molecules g molecules g Fe molecules CO = molecules CO g Fe g CO g Fe 2 = g Fe g CO b) Fe 2 + 3CO 2Fe + 3CO 2 1 mol 3 mol g g 45.7 g g f) Fe 2 + 3CO 2Fe + 3CO 2 1 mol 2 mol formula units atoms formula units atoms formula units Fe 2 = atoms Fe atoms Fe formula units Fe g Fe g CO 2 = 37.8 g CO g Fe 2 2 c) Fe 2 + 3CO 2Fe + 3CO mol CO 3 mol 2 mol 1.80 mol mol 2 mol Fe = 1.20 mol Fe 3 mol CO d) Fe 2 + 3CO 2Fe + 3CO 2 1 mol 3 mol g g 6.50 mol 6.50 mol CO g Fe 2 = 346 g Fe2 3 mol CO LESSON EXERCISES 1. Given the following reaction: N 2 O 4 + 2N 2 H 4 3N 2 + 4H 2 O a) How many g of N 2 H 4 reacts with g of N 2 O 4 b) How many g of N 2 are formed from reacting g of N 2 H 4 c) How many mol of H 2 O are produced by reacting 7.5 mol of N 2 H 4 d) How many g of N 2 O 4 are needed to react with 3.65 mol of N 2 H 4 e) How many molecules of N 2 O 4 are needed to react with g of N 2 H 4 f) How many molecules of N 2 can be made by reacting molecules of N 2 O 4 CORRESPONDENCE STUDY PROGRAM PAGE 35

22 CHEMISTRY Given the following reaction: MnO 2 + 4HCl MnCl 2 + Cl 2 + 2H 2 O a) How many g of MnO 2 reacts with g of HCl b) How many g of H 2 O are formed from reacting g of MnO 2 c) How many mol of Cl 2 are made by reacting 14.5 mol of HCl d) How many g of HCl are needed to react with 4.25 mol of MnO 2 e) How many molecules of HCl are needed to react with g of MnO 2 f) How many molecules of H 2 O can be made by reacting molecules of HCl 3. a) How many g of Al 2 (SO 4 ) 3 can be made by reacting g of Al with enough H 2 SO 4 according to the following equation: 5. How many g of CaCl 2 can be made from g of HCl according to the following equation CaO (s) + 2HCl (aq) CaCl 2(aq) + H 2 O (l) 6. Given the following balanced chemical equation: 2VO (s) + 3Fe 2 (s) 6FeO (s) + V 2 O 5(s) a) State the mol-mol relationship in this reaction. b) State the mass-mass relationship in this reaction. c) Show that this balanced chemical equation obeys the law of conservation of mass. JOURNAL ENTRY 2Al (s) + 3H 2 SO 4(aq) Al 2 (SO 4 ) 3(aq) + 3H 2(g) b) How many g of Al reacts with g of H 2 SO 4 c) How many mol of Al 2 (SO 4 ) 3 can be made from 6.00 mol of H 2 SO 4 4. Given the following reaction: 4FeS (s) + 7O 2(g) 2Fe 2 (s) + 4SO 2(g) a) How many g of O 2 reacts with g of FeS b) How many g of Fe 2 can be made from g of O 2 c) How many mol of Fe 2 can be made from mol of FeS PAGE 36 CORRESPONDENCE STUDY PROGRAM

23 UNIT 1 LIMITING REAGENT LESSON 9 In the last lesson, we calculated how much of one reactant was needed to react with a certain amount of another reactant. However, sometimes reactants are just mixed together and we try to predict how much product will be made. Under these conditions there is usually too much of one reactant present (a reactant in excess). In this case one of the reactants is used up completely. Some of the other reactant is used with some left over (in excess). The reactant that is completely used up is called the limiting reagent (reagent is a chemical). It is called the limiting reagent because it is completely used up and you cannot make any more product. It limits the amount of product that can be made. Examine the following reactions and diagram to help understand the concept of limiting reagent. C (s) + O 2(g) CO 2(g) C + O O O C O C C C C C + O O O O O O O O O C O C O C O We can see that one atom of carbon, C, reacts with one molecule of oxygen, OO, to make one molecule of carbon dioxide, OCO. If we reacted five carbon atoms, C, with three oxygen molecules, OO, it is only possible for three molecules of carbon dioxide to form. + C C Oxygen, O 2, OO, is the limiting reagent because it is completely used up and limits the amount of carbon dioxide product made. We cannot make any more CO 2 because there is no OO to react with the two carbon atoms in excess. It is easier to understand the concept of limiting reagent when we relate it to numbers of particles reacting. It is more difficult when talking about grams of reactants that are reacting. However, we have learned in this unit that grams of a chemical and numbers of particles are directly related to each other. In the same way so many particles of one reactant reacts with so many particles of another reactant, we know that a certain mass of one reactant reacts with a certain mass of another reactant. Examine some limiting reagent problems. SAMPLE PROBLEM 1 Given the reaction C (s) + O 2(g) CO 2(g) How many g of CO 2 can be made by reacting g of C with g of O 2 C (s) + O 2(g) CO 2(g) 1 mol 1 mol 1 mol g g g g g g One approach to solving a limiting reagent problem is to assume that one of the reactants is the limiting reagent. It doesn't matter which reactant you choose. Then calculate to see if there is enough of the other reactant to completely use up the reactant assumed to be the limiting reagent. In this question, assume that carbon is the limiting CORRESPONDENCE STUDY PROGRAM PAGE 37

24 CHEMISTRY 11 reagent. Next, calculate to see if the assumption is correct g C g O 2 = g O g C 2 Our calculation shows that if all of the C is to be used up, making C the limiting reagent, it will need to react with g O 2. However, we only have g of O 2 available. Therefore, C cannot be the limiting reagent. The other reactant, O 2, is the limiting reagent. If you can prove that one of the reactants is not the limiting reagent, the other reactant is the limiting reagent. We do not have to show that O 2 is the limiting reagent now that we have shown that C is not the limiting reagent. However, for demonstration purposes, look at the calculations as if we had assumed in the beginning that O 2 was the limiting reagent. Calculate how much C we would need to use up the O 2 completely g O g C = g C g O 2 We need g of C to use up all of the O 2 and make O 2 the limiting reagent. We have g of C to react, more than enough. This calculation proves that O 2 is the limiting reagent. Now that we know which reactant is used up completely, we can calculate how much product can be made. If we used g of C (the reactant in excess) to determine how much product is made, the amount calculated would be incorrect and more than the actual amount we can make in this reaction g O g CO 2 = g CO g O 2 Therefore, g CO 2 will be made in this reaction. SAMPLE PROBLEM 2 What is the maximum amount of Fe 2 that can be made by reacting g of FeS with g of O 2 according to the following reaction 4FeS (s) + 7O 2(g) 2Fe 2 (s) + 4SO 2(s) 4 mol 7 mol 2 mol g g g g g g Assume that FeS is the limiting reagent, (LR) g FeS g O 2 = g O g FeS If FeS is the limiting reagent we would need g of O 2 to use up the FeS completely. There is g of O 2 available to react, more than enough. Therefore, FeS is the limiting reagent g FeS g Fe 2 = g Fe2 O g FeS 3 The maximum amount of Fe 2 that can be made by reacting g of FeS with g of O 2 is g of Fe 2. SAMPLE PROBLEM 3 Given the reaction: Li 3 N (s) + 3H 2 O (l) 3LiOH (s) + NH 3(g) How much LiOH can be made by reacting g of Li 3 N with g H 2 O PAGE 38 CORRESPONDENCE STUDY PROGRAM

25 UNIT 1 Li 3 N (s) + 3H 2 O (l) 3LiOH (s) + NH 3(g) 1 mol 3 mol 3 mol g g g g g g Assume that Li 3 N is the limiting reagent g Li 3 N g H 2O = g H2 O g Li 3 N If Li 3 N is the limiting reagent, we need g of H 2 O to use up the Li 3 N completely. However, we only have g of H 2 O available to react. Therefore, Li 3 N is not the limiting reagent. The H 2 O is the limiting reagent g H 2 O g LiOH = g LiOH g H 2 O Therefore, g of LiOH can be made from reacting g of Li 3 N with g H 2 O. SAMPLE PROBLEM 4 a) Assume that CH 4 is the limiting reagent mol CH g O 2 = 107 g O2 1mol CH 4 We need 107 g of O 2 to use up all of the CH 4 if CH 4 is the limiting reagent. However, we have only g O 2 available to react. The CH 4 cannot be the limiting reagent. The O 2 is the limiting reagent g O g CO 2 = g CO g O 2 Therefore, g CO 2 can be made by reacting 1.68 mol of CH 4 with g O 2 b) g O g CH 4 = g CH g O g CH 4 needed 1.68 mol CH g CH 4 = 27.0 g CH4 1 mol CH g CH 4 available 27.0 g CH g CH 4 = 15.6 g CH 4 left over, unreacted. Given the reaction: CH 4(g) + 2O 2(g) CO 2(g) + 2H 2 O (g) a) How much CO 2 can be made by reacting 1.68 mol of CH 4 with g of O 2 b) How many g of the reactant in excess will be left over unreacted CH 4(g) + 2O 2(g) CO 2(g) + 2H 2 O (g) 1 mol 2 mol 1 mol g g g 1.67 mol g g LESSON EXERCISES 1. Given the reaction: Mg (s) + 2HCl (aq) MgCl 2(aq) + H 2(g) How many g of MgCl 2 can be made by reacting g of Mg with g of HCl 2. Given the following reaction: (NH 4 ) 2 SO 4 + Pb(N ) 2 2NH 4 N + PbSO 4 How much NH 4 N can be made by reacting g of (NH 4 ) 2 SO 4 with g of Pb(N ) 2 CORRESPONDENCE STUDY PROGRAM PAGE 39

26 CHEMISTRY Given the reaction: 2Na 3 PO 4 + 3BaCl 2 Ba 3 (PO 4 ) 2 + 6NaCl What is the maximum amount of NaCl that can be produced by reacting g of Na 3 PO 4 with g of BaCl 2 4. Given the reaction: P 4 + 5O 2 P 4 O 10 a) How much product can be made by reacting g of P 4 with g of O 2 b) How many g of the reactant that is in excess will be left over unreacted 5. Given the reaction: SF 4 + 2H 2 O SO 2 + 4HF How much HF can be produced by reacting g of SF 4 with 0.75 mol of H 2 O JOURNAL ENTRY PAGE 40 CORRESPONDENCE STUDY PROGRAM

27 UNIT 1 PERCENT YIELD OF A CHEMICAL REACTION LESSON 10 In this unit we calculated how much product could be made if we reacted a certain amount of reactants, or if the limiting reagent was completely consumed. This yield of products we calculated is the predicted or theoretical yield under ideal conditions. Theoretical yield is the amount of products predicted in a chemical reaction based on stoichiometry calculations and assuming a complete reaction of all of the limiting reagent. The theoretical yield, however, is not usually equal to the actual yield for a chemical reaction. Actual yield is the measured amount of products made. What would make the actual yield less than the theoretical yield for a reaction There are a number of possible reasons. 1. One of the major reasons is the "competing reaction" or "side reaction" that is taking place. Examine the following reaction: main reaction 2X + 3Y 2 2XY 3 We are trying to produce XY 3. We calculate the theoretical yield of XY 3 based on how much of the reactants react. If there is a competing or side reaction competing reaction XY 3 + Y 2 XY 5 in which some of the XY 3 produced is further reacted to make XY 5, then the actual yield will be lowered. There will be less XY 3 to be measured in the container after the chemical reactions. 2. Another possible reason for the actual yield to be less than the theoretical yield is that one of the reactants or products may be very volatile and evaporate. 3. Sometimes some of a product might stick to the inside walls of a glass container used. 4. If an insoluble product (precipitate) is produced and filtered to separate it from the solvent, a small amount of this product actually dissolves in the solvent and represents lost product in our measurements. When chemicals are made in huge quantities in the chemical industry it is useful to know the per cent yield for a chemical reaction. actual yield per cent yield = 100 theoretical yield This is important because it shows exactly how much chemicals have to be reacted to make a certain amount of products. The per cent yield tells us how close the actual amount of products made is to the predicted maximum theoretical amount. The higher the per cent yield, the better. SAMPLE PROBLEM 1 Given the following reaction: Fe (s) + S (s) FeS (s) Calculate the per cent yield for this reaction if g of FeS was made when g of Fe was completely used up by an excess of S. Fe (s) + S (s) FeS (s) 1 mol 1 mol 1 mol g g g CORRESPONDENCE STUDY PROGRAM PAGE 41

28 CHEMISTRY 11 actual yield percent yield = 100 theoretical yield We know the actual yield, but we must calculate the theoretical yield g Fe g FeS = g FeS g Fe The theoretical yield is g FeS. This is the maximum amount of product we could hope to make. percent yield = g FeS 100 = 89.52% g FeS SAMPLE PROBLEM 2 Given the following reaction 3Fe + 4H 2 O Fe 3 O 4 + 4H 2 Calculate the percent yield if g of Fe 3 O 4 was produced when g of H 2 O was reacted with an excess of Fe. 3Fe + 4H 2 O Fe 3 O 4 + 4H 2 4 mol 1 mol g g The theoretical yield is: g H 2 O g Fe 3O 4 = g Fe3 O g H 2 O percent yield = g Fe 3O = 82.16% g Fe 3 O 4 SAMPLE PROBLEM 3 Given the following reaction: H 2 S (g) + 2NaOH (aq) Na 2 S (aq) + 2H 2 O (l) A mass of 6.75 g of NaOH was reacted with 4.50 g of H 2 S. One of the products, Na 2 S, had a mass of 5.95 g. Calculate the percent yield. H 2 S (g) + 2NaOH (aq) Na 2 S (aq) + 2H 2 O (l) 1 mol 2 mol 1 mol g g g 6.75 g 4.50 g g Before we can calculate the percent yield we must determine which reactant is the limiting reagent and then calculate the theoretical yield. Assume that NaOH is the limiting reagent g NaOH g H 2S = 2.88 g H 2 S g NaOH If NaOH is the limiting reagent we need 2.88 g of H 2 S to use it up completely. We have 4.50 g of H 2 S available, more than enough. Therefore, NaOH is the limiting reagent. The theoretical yield is: 6.75 g NaOH g Na 2S = 6.59 g Na 2 S g NaOH percent yield = 5.95 g Na 2S 100 = 90.3% 6.59 g Na 2 S LESSON EXERCISES 1. Given the following reaction AgN(aq) + KBr (aq) AgBr (s) + KN(aq) Calculate the per cent yield if 7.15 g of AgBr was produced when 7.00 g of AgN was reacted with an excess of KBr. 2. Aspirin, C 9 H 8 O 4, is produced by reacting salicylic acid, C 7 H 6, with acetic PAGE 42 CORRESPONDENCE STUDY PROGRAM

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