1. Chemical Formulas & Chemical Compounds Chapter 7

2. Section 3 Using Chemical Formulas Objectives CalculateCalculate the formula mass or molar mass of any given compound. UseUse molar mass to convert between mass in grams and amount in moles of a chemical compound. CalculateCalculate the number of molecules, formula units, or ions in a given molar amount of a chemical compound. CalculateCalculate the percentage composition of a given chemical compound.

3. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chapter 7 A chemical formula indicates: the elements present in a compound the relative number of atoms or ions of each element present in a compound Chemical formulas also allow chemists to calculate a number of other characteristic values for a compound: formula mass molar mass percentage composition Section 3 Using Chemical Formulas

4. Using Chemical Formulas Formula Mass or Molecular Mass – sum of the average atomic masses of all atoms present in a compound or molecule. (Formula or Molecular Weight) H 2 OKClO 3 H – 2x1.01= 2.02K – 1x39.10=39.10 O – 1x16.00=16.00Cl – 1x35.45=35.45 18.02 amuO – 3x16.00=48.00 122.55 amu

5. Molar Mass – mass of one mole of a substance in grams. H 2 O –H 2x1.01= 2.02 –O 1x16.00=16.00 18.02 g/mol Ba(NO 3 ) 2 –Ba 1x137.33=137.33 –N 2x14.01=28.02 –O 6x16.00=96.00 261.35 g/mol

6. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chapter 7 Visual Concepts Click below to watch the Visual Concept. Visual Concept The Mole

7. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chapter 7 Visual Concepts Click below to watch the Visual Concept. Visual Concept Molar Mass as a Conversion Factor

8. Conversions using Molar Mass

9. What is the mass in grams of 2.5 mol of O 2 gas?

10. A normal tablet of ibuprofen contains 200 mg of C 13 H 18 O 2, how many moles of ibuprofen are in a single tablet? How many molecules? C – 13x12.01=156.13 H – 18x1.01= 18.18 O – 2x16.00=32.00 206.31 g/mol

11. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chapter 7 Percentage Composition of Iron Oxides Chapter 7 Section 3 Using Chemical Formulas

12. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chapter 7 Visual Concepts Click below to watch the Visual Concept. Visual Concept Percentage Composition

13. Cu 2 S Cu – 2x63.55=127.10 S – 1x32.07 = 32.07 159.17 amu

14. Chapter 7 Section 4 Determining Chemical Formulas Objectives DefineDefine empirical formula, and explain how the term applies to ionic and molecular compounds. DetermineDetermine an empirical formula from either a percentage or a mass composition. ExplainExplain the relationship between the empirical formula and the molecular formula of a given compound. DetermineDetermine a molecular formula from an empirical formula.

15. Empirical and Actual Formulas Chapter 7 Section 4 Determining Chemical Formulas

16. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chapter 7 Comparing Empirical and Molecular Formulas Chapter 7 Section 4 Determining Chemical Formulas

17. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chapter 7 Visual Concepts Click below to watch the Visual Concept. Visual Concept Comparing Molecular and Empirical Formulas

18. Steps for Determining an Empirical Formula 1.Start with the number of grams of each element, given in the problem. If percentages are given, assume that the total mass is 100 grams so that the mass of each element = the percent given. 2.Convert the mass of each element to moles using the molar mass from the periodic table. 3.Divide each mole value by the smallest number of moles calculated. 4.Round to the nearest whole number. This is the mole ratio of the elements and is represented by subscripts in the empirical formula. If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same factor to get the lowest whole number multiple. e.g. If one solution is 1.5, then multiply each solution in the problem by 2 to get 3. e.g. If one solution is 1.25, then multiply each solution in the problem by 4 to get 5.

19. Example Problem #1 A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675 g H. What is the empirical formula of the compound? Start with the number of grams of each element, given in the problem. Convert the mass of each element to moles using the molar mass from the periodic table.

20. Divide each mole value by the smallest number of moles calculated. Round to the nearest whole number. This is the mole ratio of the elements and is represented by subscripts in the empirical formula

21. Example Problem #2 NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of NutraSweet and find the molecular formula. (The molar mass of NutraSweet is 294.30 g/mol) Start with the number of grams of each element, given in the problem. If percentages are given, assume that the total mass is 100 grams so that the mass of each element = the percent given. Convert the mass of each element to moles using the molar mass from the periodic table.

22. Divide each mole value by the smallest number of moles calculated. Round to the nearest whole number.

23. This is the mole ratio of the elements and is represented by subscripts in the empirical formula. If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same factor to get the lowest whole number multiple.

24. Now, we can find the molecular formula by finding the mass of the empirical formula and setting up a ratio:

25. Homework Pages 251-254 4,6,7,10,11,14,15,16,18,21,23,24,25/Numbers 4,6,7,10,11,14,15,16,18,21,23,24,25/ 28,31,32,36,38,39,40,44,50