Lesson 2: Wrap-Up
Well, this brings us to the end of this lesson. We covered quite a bit of material that is fundamental to an understanding of chemistry.
One of the first concepts we looked at is that there are different types of knowledge. You should, of course, be able to distinguish between observations and inferences. You should also be able to describe the differences between such things as scientific laws and theories.
Another concept was that theories have to be evaluated and not just accepted as fact. With that in mind, how does Dalton's theory check out? First, did it explain established observations? Yes, it explained why elements combined in fixed ratios to form compounds (Law of Constant Composition). Did it explain new observations? Yes, with some modifications, it was able to explain the Law of Combining Volumes. Did it predict new observations? Yes, it led Dalton to formulate the Law of Simple Multiple Proportions. What it simple? Yes, but as more was discovered about atoms, it was found to be too simple.
We talked about Dalton's concept of atoms and Avogadro's concept of molecules, about elements, atomic weights, compounds, symbols and formulas. All of these are fundamental aspects of chemistry that will be used throughout this course.
We also dealt with different kinds of formulas and a host of calculations. These are not complex calculations, but you need to be proficient at them. Before long they will be taken for granted. If you have taken a chemistry course before this, you may already have been proficient at them.
As you finish studying this lesson, go back through the objectives and make sure that you can do all of them and then take the self-quiz for the lesson. Remember that you have a problem set and a lab report to be turned in for this lesson.
Lab Report Format | Additional Practice Problems & Answers | Self Quiz
Lab Report Format
(Refer to Ex. 18 in your workbook.) The report must be typed, although you may write in the calculations by hand if it is too cumbersome to figure out how to type them in an understandable format. Be sure to include your name and the names -- first and last -- of any coworkers and a title for the experiment. Then state the purpose of the experiment; give the experimental purpose here, not the educational purpose of the lesson as a whole. List and label the data you obtained in your measurements. (List all raw data. This does not include results of calculations. In this experiment, for example, the mass of the empty syringe is data, but the mass of the gas samples is not.) That list can be similar to the table in your workbook or in your own style, but be sure to label the data properly. Then show how you calculated the formula weights of the gases. This is particularly important in this experiment because the instructions do not tell you exactly how to do this. So you must make clear in your report what you have done. Show all your calculations, even the trivial ones. Be sure and label all your calculations so your instructor can easily follow what you have done.Be sure to take significant digits into account and include all the proper units. Then in the results section report the formula weights of the gases that you determined in this experiment. (The results are the answer to the question implicit in the purpose.) If you have comments to make about problems you ran into or variations that you made in the procedure, include those in a comments section. If you believe your results are too high or too low, you can make that comment and speculation on the reason for it. After that, write out your answers to the questions about the experiment that are listed in the workbook in example 19e. Your report for this experiment will be graded on the accuracy of your data and results, the correctness of your calculations, the answers to the questions, the format of your report, and even the neatness of your report.
Additional Practice Problems & Answers
Here are some additional practice problems for this lesson. (Use atomic weights that are precise to the hundredths place.)
Gram-mole Practice Problems
1. What is the weight of 0.30 mole of sulfur?
2. What is the weight of 5.5 mole of silicon?
3. How many moles are there in 45 g of Cl?
4. Change 34 g of lithium to moles.
5. What is the weight of 4.30 mole of sodium?
6. What is the weight of 1.75 mole of Ca?
7. How many moles are there in 85.3 g of P?
8. Change 0.566 g of silver to moles.
Practice Calculating Formula Weights
| 1. HCl | 5. NaCl |
| 2. SiH4 | 6. CaSO4 |
| 3. C3H6O2 | 7. C2H4O2 |
| 4. Fe(NO3)3 | 8. Sn(CO3)2 |
More Gram-mole Practice Problems
1. Change 3.4 moles of HCl to grams. 5. Change 3.4 moles of NaCl to grams. 2. Change 8.5 g of SiH4 to moles. 6. Change 8.5 g of CaSO4 to moles. 3. Change 5.20 moles of C3H6O2 to grams. 7. Change 5.20 moles of C2H4O2 to grams. 4. Change 13.2 g of Fe(NO3)3 to moles. 8. Change 13.2 g of Sn(CO3)2 to moles.
Practice Calculating Percent Composition From Formula
| 1. SO2 | 4. CH2O |
| 2. Na2CO3 | 5. CaCO3 |
| 3. FeCO3 | 6. Fe(OH)3 |
Practice Determining Empirical Formulas
1. Calculate the empirical formula of the compound that contains 1.0 g S for each 1.5 g O.
2. Calculate the empirical formula of the compound containing 75.0% C and 25.0% H.
3. Calculate the empirical formula of the compound containing 81.8% C and 18.2% H.
4. Determine the empirical formula of the compound containing 37.5% C, 12.5% H, and 50.0% O by weight.
5. Determine the empirical formula of the compound containing 26.1% C, 4.3% H, and 69.6% O by weight.
6. Determine the empirical formula of the compound containing 38.7% C, 16.1% H, and 45.2% N by weight.
7. What is the empirical formula of a compound if a 50.0 g sample of it contains 9.1 g Na, 20.6 g Cr, and 22.2 g O?
Practice Determining Molecular Formulas
1. Empirical formula C3H7, molecular weight 86 g/mole
2. Empirical formula S, molecular weight 256 g/mole
3. Empirical formula CH, molecular weight 26 g/mole
4. Empirical formula NO2, molecular weight 46 g/mole
Answers
Gram-mole Practice Problems
1. What is the weight of 0.30 mole of sulfur? 9.6 g
2. What is the weight of 5.5 mole of silicon? 150 g (only 2 sig. dig.)
3. How many moles are there in 45 g of Cl? 1.3 mol
4. Change 34 g of lithium to moles. 4.9 mol
5. What is the weight of 4.30 mole of sodium? 98.9 g
6. What is the weight of 1.75 mole of Ca? 70.1 g
7. How many moles are there in 85.3 g of P? 2.75 mol
8. Change 0.566 g of silver to moles. .00525 mol
Practice Calculating Formula Weights
1. HCl 36.46 g/mol 5. NaCl 58.44 g/mol 2. SiH4 32.13 g/mol 6. CaSO4 136.15 g/mol 3. C3H6O2 74.09 g/mol 7. C2H4O2 60.06 g/mol 4. Fe(NO3)3 241.88 g/mol 8. Sn(CO3)2 238.73 g/mol
Gram-mole Practice Problems
1. Change 3.4 moles of HCl to grams. 120 g
2. Change 8.5 g of SiH4 to moles. .26 mol
3. Change 5.20 moles of C3H6O2 to grams. 385 g
4. Change 13.2 g of Fe(NO3)3 to moles. .0546 mol
5. Change 3.4 moles of NaCl to grams. 2.0x102 g
6. Change 8.5 g of CaSO4 to moles. .062 mol
7. Change 5.20 moles of C2H4O2 to grams. 312 g
8. Change 13.2 g of Sn(CO3)2 to moles. .0553 mol
Practice Calculating Composition From Formula
1. SO2 50.05% S and 49.95% O 4. CH2O 39.99% C, 6.73% H, and 53.28% O 2. Na2CO3 43.38% Na, 11.33% C, and 45.29% O 5. CaCO3 40.04% Ca, 12.00% C, and 47.96% O 3. FeCO3 48.20% Fe, 10.37% C, and 41.43% O 6. Fe(OH)3 52.25% Fe, 44.91% O, and 2.83% H
Practice Determining Empirical Formulas
1. Calculate the empirical formula of the compound that contains 1.0 g S for each 1.5 g O. SO3
2. Calculate the empirical formula of the compound containing 75.0% C and 25.0% H. CH4
3. Calculate the empirical formula of the compound containing 81.8% C and 18.2% H. C3H8
4. Determine the empirical formula of the compound containing 37.5% C, 12.5% H, and 50.0% O by weight. CH4O
5. Determine the empirical formula of the compound containing 26.1% C, 4.3% H, and 69.6% O by weight. CH2O2
6. Determine the empirical formula of the compound containing 38.7% C, 16.1% H, and 45.2% N by weight. CH
5N 7. What is the empirical formula of a compound if a 50.0 g sample of it contains 9.1 g Na, 20.6 g Cr, and 22.2 g O? Na2Cr2O7
Practice Determining Molecular Formulas
1. Empirical formula C3H7, molecular weight 86 C6H14
2. Empirical formula S, molecular weight 256 S8
3. Empirical formula CH, molecular weight 26 C2H2
4. Empirical formula NO2, molecular weight 46 NO2
Self Quiz & Answers
List of atomic weights:
N = 14.01 g/mole O = 16.00 g/mole S = 32.07 g/mole Cl = 35.45 g/mole Cr = 52.00 g/mole
1. What is the formula weight (molecular weight) for NO2?
2. A compound has the molecular formula N2O4. What is its empirical formula?
3. Write the formula for a molecule which has two nitrogen atoms and five oxygen atoms.
4. A compound contains .328 g Cr and .672 g Cl. What is its empirical formula?
5. A compound with an empirical formula of CH2 has a molecular weight of 56. What is its molecular formula?
6. What is the relationship between a mole of an element and the atomic weight of the element?
7. How many moles are there in 8.00 g of SO2?
8. Calculate the percent composition by weight of SO2.
9. What is the weight in grams of .60 moles of S?
10. State the basic points of Dalton's theory.
Answers
1. What is the formula weight (molecular weight) for NO2? 46.0 g/mole
2. A compound has the molecular formula N2O4. What is its empirical formula? NO2
3. Write the formula for a molecule which has two nitrogen atoms and five oxygen atoms. N2O5
4. A compound contains .328 g Cr and .672 g Cl. What is its empirical formula? CrCl3
5. A compound with an empirical formula of CH2 has a molecular weight of 56. What is its molecular formula? C4H8
6. What is the relationship between a mole of an element and the atomic weight of the element?
The atomic weight of an element is the weight in grams of one mole of that element.
7. How many moles are there in 8.00 g of SO2? 0.125 moles
8. Calculate the percent composition by weight of SO2. 50.1 % sulfur and 49.9 % oxygen
9. What is the weight in grams of 0.60 moles of S? 19 g
10. State the basic points of Dalton's theory.
Elements consist of very small, indivisible particles called atoms. All atoms of a given element are identical and have a characteristic mass; atoms of different elements are different from one another. Compounds are made from atoms of different elements adhering to one another in a fixed ratio. Chemical reactions involve the rearrangement of atoms; atoms are not created or destroyed.