Chapter 4

1. Chapter 4 Molar Relationships

2. Part 1 Counting Atoms • Objective: Define the term mole and describe how it is used in chemistry. • A mole is the amount of a substance that contains 6.02 x 1023 particles. • A pair is ….. • A dozen is … • A gross is … • A mole is …

3. A mole • You can have a mole of… • Atoms, molecules, an element, a compound. • The 6.02 x 1023 is called Avogadro’s number and answers the question “how many.” • The average animal mole is 15 cm long. • 6.02 x 1023moles head to tail would stretch from the earth to Alpha Centauri and back. 4.37 light-years away.

4. The Mole • An average animal mole weighs 150 g. • 6.02 x 1023moles would weigh as much as 1% of the earth. Weigh 1.3 times the mass of the moon. Weigh 60 times the combined mass the Earth’s oceans!! • An average animal mole is 5 cm tall. • 6.02 x 1023moles would cover the earth at a depth of 8 million moles !!

5. Mass of a mole of atoms • The atomic mass is found on the periodic chart. We also call this the atomic mass unit or amu. • The amu for aluminum is 26.98. So, you can say that one atom of Al has a mass of 26.98 amu. • Even better is the molar mass. • The molar massis the mass of one mole of an element. If you have 26.98 gmm of Al you must have 6.02 x 10 23 atoms of Al.

6. Mass of a mole of atoms • The mass of one atom of Al is 26.98 amu • The mass of a mole of Al atoms is 26.98 gmm • Our book drops the mm. 26.98 g of Al = a mole Practice: How many moles of carbon are in 26 g of carbon? 26 g of carbon x 1 mole of carbon = 2.2 moles 12 g of carbon

7. A chart may help • You can have a mole of • atoms, molecules and compounds • Practice: Calculate the gmm of CuSO4 • Cu = 63.5 + S = 32.1 + O = 4 x 16 = 160.6 g • So, 160.6 g is made up of 6.02 x 10 23 units of CuSO4

8. 4.3 Molar Conversions for Elements • How many moles of nickel are in 25 g of nickel? • How many grams of copper are in 2.50 moles of copper?

9. 4.4 Molar Conversions for Compounds • How many molecules of CO2 are in 0.75 moles of CO2? • How many O2 atoms are in 1.75 moles of CO2? • You practice pg 129 #’s 3,4, pg 131 #’s 5,6 and pg 132 #’s 7,8.

10. Part 1 Review • Page 135 #’s 11, 12, 13, 14, 16, 17

11. Part 3 Formula Calculations • Objective: To determine the percent composition of elements that compose a particular compound. • Part over whole. What percent of carbon is in a 50 g sample of carbon dioxide? CO2 C = 12 g, O = 2 x 16 so CO2 = 44g Remember part/whole= 12/44 x 100 = 27.3 % Practice page 144 #’s 32, 33

12. 4.10 Empirical Formulas • Objective: understand how to find empirical and molecular formulas from percent compositions data of a compound. • An empirical formula is the smallest whole-number ratio of the various types of atoms in a compound. Huh? • H20 and H202 and H30 and OH

13. 4.10 Empirical Formulas • Like this. What is the empirical formula for an iron oxide compound with the following composition? 36 gram sample is 78% iron. 78g of iron x 1mole of iron = 1.4 moles 55. 8 g of iron 22g of oxygen x 1 mole of oxygen = 1.4 moles 16.0 g of oxygen

14. 4.10 Empirical Formulas • So the empirical formula is FeO • Practice: Page 146 #’s 34 and 35 • Table 4-3 is helpful!!

15. 4.11 Molecular Formulas • This is easy. • Take the molecular formula mass • empirical formula mass • This results in a multiplier used to multiply the subscripts of the empirical formula. • Try this. Calculate the molecular formula for ethylene glycol(CH3O). Molar mass is 62g/mol • 62/31 = 2 so C2H6O2 • Try #38 page 150.

16. Significant Digits 4.12 • Measurements are precise when they are clustered. Page 151 figure 4-16 • Measurements are accurate when they are in agreement with the true value. • Describe a basketball players shots that are… • Precise but not accurate. • Precise and accurate.

17. Percent Error • Error is the difference between the accepted value minus the experimental value. • Percent error is the absolute value of the calculated error divided by the accepted value times one hundred. • Percent error = I error I x 100 % • accepted value

18. Sig figs • Significant figures in a measurement include all of the digits that are known, plus a last digit that is estimated. • The rules: • 1) Every nonzero digit in a reported measurement is assumed to be significant. • 24.7 0.743 714

19. Sig figs • 2) Zeros between nonzero digits are significant. • 7003 meters 1.505 meters 40.89 km • 3) Leftmost zeros appearing in front of nonzero digits are not significant. • 0.0005 mm, 0.43 mL, 0.000000009 m • Changing these to scientific notation cures this problem.

20. 0.0005 mm, 0.43 mL, 0.000000009 m • Changing these to scientific notation cures this problem. • 5 x 10 -4 mm, 4.3 x 10 -1, 9 x 10 -9 • 4) Zeros at the end of a number and to the right of a decimal point are always significant. • 43.00 meters, 1.010 km, 9.000 mL

21. 5) Zeros at the rightmost end of a measurement that lie to the left of an understood decimal point are not significant. • 7000 meters, 27210 km, • 6) Unlimited significant figures occur when you count. There are 21 students in class. • 21.00000000000000000000000000000000000000000000000000000000000000000000000

22. 6) Unlimited significant figures occur when you count. There are 21 students in class. • 21.00000000000000000000000000000000000000000000000000000000000000000000000 • Unlimited significant figures occur when you have a known defined quantity like • 60 minutes = 1 hour • 1000 mL = 1 liter

23. Rounding • Five and above = round up, • four or less is dropped • Addition and subtraction. Line up the decimal points and round to the least number of significant figures. • 6.12 m + 9.35 m + 61.2 m • 34.61 m - 17.3 m

24. Rounding and sig figs • Multiplication and division • Round to the same number of significant figures as the measurement with the least number of significant figures. • 7.55 m x 0.34 m • 2.4526 meters / 0.0200 meters

25. Part 3 Review • Page 153 # s 40, 41, 42, 43, 45, 46, 49, 50, 51 • Test Prep Questions • Page 157 #’s 5 do one., 13, 14, 20a,e, 31, 37, 40 a-f, 41 a,b,d.