1. Chemical Quantities Chapter 10

2. The Mole  a mole is an amount of matter  mass is also an amount of matter, however the mole is much more useful to chemists, in some cases  a mole is based an a certain amount of particles, sometimes called representative particles, these include:  atoms when dealing with an element  molecules when dealing with molecular compounds and,  formula units when dealing with ionic compounds

3. Counting particles  When we study chemical reactions, we will find that we often want to tract the changes of particles, not grams or ounces  To discuss reactions in terms of one particle, or 10 particles or even 1000 particles is silly, these are ridiculously small amounts of matter  To make a mole practical, we need to consider a very large number of particles. The number we use is 6.02 x10 23 particles = 1 mole. This is know as Avagadro’s number.

4. The mole is just another unit  Remember all the way back to chapter 3, the book told us that the mole was the SI Unit for the amount of substance  That means that we can use the relationship 6.02x10 23 particles = 1 mole is a conversion factor we can use to convert from particles to moles and moles to particles  Example:  How many moles is 3.5 x 10 24 atoms of neon?

5. More examples  How many moles is 7.1 x 10 22 formula units of sodium chloride?  How many molecules are there in 2.9 moles of carbon dioxide?  How many formula units are there in 0.45 moles of silver nitrate?

6. The Mole and Mass  When you have one mole (6.02 x 10 23 atoms) of any element, the mass just happens to equal the element’s atomic mass in grams (not amu)  For example, if you have one mole of carbon, it would have mass of 12.01 grams, if you have one mole of silver, it would have a mass of 107.9 grams  If you want to try to understand why this is, see Table 10.2 on p. 293

7. Some more examples  How many grams is one mole of helium?  How many moles is 22.99 grams of sodium?  How many grams is 2.5 moles of helium?

8. Molar Mass  We can create mole-mass conversion factors for any compound by adding up the atomic masses of the atoms in the compound. This is called the molar mass.  For example, find the molar mass of SO 3 :  How many grams are in 1 mole of SO 3 ? 2 moles of SO 3 ?

9. Practice finding Molar Mass  What is the molar mass of carbon dioxide?  What is the molar mass of nitrogen gas?  What is the molar mass of calcium hydroxide?

10. Mole – Mass Relationships  Molar mass is a conversion factor you can use to convert moles to grams or grams to moles  How many moles is 25.0 g of water?  If you need 13.7 moles of water, how many grams do need to weigh out?

11. More Mole - Mass Examples  How many grams of silver nitrate are needed if you need 0.64 moles to run the reaction?  How many moles are there in 50.0 grams of aluminum sulfide?

12. Here’s a tricky one: How many grams is 8.45 x 10 22 molecules of ethanol (C 2 H 5 OH)?

13. Convert 4.24 x 10 22 molecules of PCl 5 to moles Calculate the molar mass of Fe 2 (SO 4 ) 3

14. Mole – Volume Relationship  At standard temperature and pressure (STP), one mole of any gas occupies 22.4 L of volume  standard temperature is 0oC and standard pressure is 1 atm or 101.3 kPa of pressure  the conversion factor, 1 mole = 22.4 L can only be used for gases at STP!

15. Some Examples  What is the volume of 0.47 mole of carbon dioxide gas at standard temperature and pressure?  How many moles of helium fill a balloon that is 11.4 L in volume (at STP)?  What is the volume of 2.75 moles of chlorine gas (at STP)?

16. Calculating the density of a gas at STP  Molar mass and 1 mole = 22.4 L can be used to find the density of a gas at STP  For example, what is the density of helium gas at STP

17. Calculating Molar Mass from Density (for gases at STP)  A gas composed of sulfur and oxygen has a density of 3.58 g/L. What is the molar mass of this gas?  Do you think this is sulfur dioxide or sulfur trioxide or neither?

18. Percent Composition  Mass percent is the percentage of each element in that compound by mass  Calculate the mass percent of each element in NH 4 Cl

19. Empirical and molecular formulas  The empirical formula just shows the ratios of the elements – it is the lowest whole number ratio of atoms in the compound  Example, the empirical formula for benzene (C 6 H 6 ) is CH  Can use percent composition to find the empirical formula

20. Empirical and molecular formulas  The empirical formula just shows the ratios of the elements – it is the lowest whole number ratio of atoms in the compound  Example, the empirical formula for benzene (C 6 H 6 ) is CH  Can use percent composition to find the empirical formula  The molar mass is needed to turn an empirical formula into a molecular formula

21. Empirical and molecular formulas  A compound is found to be 92.26 % carbon, by mass and 7.14 % hydrogen, by mass. What is the empirical formula?  If the molar mass of the compound is 26.038 g/mol, what is the molecular formula?

22. Empirical and molecular formulas  A compound is found to have the following mass percents – 42.82 % C. 49.97 % N, and 7.21 % H. What is the empirical formula?  If the molar mass of the compound is 56.068 g/mol, what is the molecular formula?

23. Two practice problems: A sample of carbon dioxide gas occupies a volume of 55.9 L at STP. How many oxygen atoms are in this sample of carbon dioxide? What is the molar mass if a gas that has a density of 6.52 g/L at STP? The gas is known to be either SO 3, PCl 5, SF 6 or XeF 2. What is the gas?

24. Another practice problem: After decomposing an unknown compound, a chemist has found it to have the following mass percents: 40.5% carbon, 6.8% hydrogen, 9.44% nitrogen and 43.26% sulfur. Other experiments determine the compound’s molar mass to be 296.54 g/mol. What are the empirical and molecular formulas?