1. The Mole  Just to clear up any misconceptions when we use the term “mole” we are not referring to this small blind fellow

2. The Mole  Instead when we talk about a mole in chemistry we think of this gentleman to the right: Amedeo Avogadro

3. The Mole  So what is the mole?  In chemistry we have the need to count how many atoms are in a particular substance, yet if atoms are so small how are we able to do this?  Answer: We use a measurement called the Mole

4. 6.02x10 23  The mole = 6.02x10 23 representative particles the same way that a dozen eggs = 12 eggs

5. Representative Particles  Representative particles are the species present in a substance  Atoms (molecules for diatomic molecules)  Molecular compounds (molecules)  Ionic compounds (formula units)  One mole of any substance is always said to contain 6.02x10 23 representative particles

6. Representative Particles  Examples of representative particles:  One rep. particle of carbon dioxide, a molecular compound, is made up of one carbon atom and two oxygen atoms (CO 2 )  Helium is made up of helium atoms  Diatomic molecules such as hydrogen (H 2 ) contain two atoms of that particular element (O 2, N 2, F 2, Cl 2 are all examples of diatomic molecules)

7. Making the Connection  We can physically count how many eggs are in one dozen, however it is almost impossible, and extremely impractical, to try and count how many rep. particles of a compound (e.g. H 2 O) are in 1 mL of solution.  So, instead of counting we call that number of particles a mole (mol) of that compound.

8. Conversion Factors  If we are trying to determine how many molecules are in one mole of a molecular compound, such as CO 2, we simply take one mole of CO 2 and use the conversion factor 6.02x10 23 molecules/1 mol of CO 2 to determine how many molecules we will have.

9. Conversion Factors  If we are trying to determine how many atoms are in one mole of CO 2 we must perform the conversion we just did, however, we must add another step:  How many atoms make up one molecule of CO 2 ?

10. Conversion Factors  How many atoms are in one mole of the following compounds:  H 2 O  NaCl  How many oxygen atoms are in 3 moles of the following compounds:  Al 2 O 3  C 6 H 12 O 6

11. The Mass of a Mole of an Element  Let’s first define the following term:  Molar Mass: the atomic mass of an element or compound expressed in grams.  Carbon has an atomic mass of 12 amu (atomic mass units) and, thus, a molar mass of 12 g.  One mole of C is equal to the molar mass of C..  Similarly, we can say that one mole of any element is equal to the element’s molar mass.

12. The Mass of a Mole of a Compound  How do we determine the mass of one mole of a compound?

13. The Mass of a Mole of a Compound  First, we must look at the chemical formula. For example: SO 3, which is called __________________.  Second, we must determine the type and number of atoms that make up the compound. In this case SO 3 consists of 1 sulfur atom and 3 oxygen atoms.

14. The Mass of a Mole of a Compound  Third, we must determine the molar mass of all of the atoms that make up the compound and add them together. For example:  1 S atom has a molar mass of 32.1g  3 O atoms have a molar mass of 3(16.0g)  For a total of 80.1g  So, we can say that the mass of one mole of SO 3 = 80.1g

15. The Mass of a Mole of a Compound  Calculating molar mass values is an important skill that you will use very often in chemistry, especially when determining how much of a compound that you will need to add to a chemical reaction to attain the desired product.

16. The Mass of a Mole of a Compound  Find the molar mass of each compound:  C 2 H 6  PCl 3  C 3 H 2 OH  N 2 O 5

17. Molar Mass  Remember: Molar mass can be defined as the mass (in grams) of one mole of any substance.  The unit that we use to describe molar mass is g/mol.  Let’s look at an example:  How many grams are in 9.45 mol of dinitrogen trioxide (N 2 O 3 )?

18. Molar Mass  First, we must map out our conversion problem by looking at what we are starting with and what we are looking for.  In this case we are going from moles of N 2 O 3 to grams of N 2 O 3.  Second, we must determine the molar mass of N 2 O 3, which is?

19. Molar Mass  Third, we need to set up our linear conversion (use the factor-label method like we used with metric conversions) using the molar mass as a conversion factor:  Last, we obtain our answer, which is?

20. Moles to Mass  Find the mass, in grams, of each of the following quantities:  3.32 mol K  4.52x10 -3 mol C 20 H 42  0.0112 mol K 2 CO 3

21. Mass to Moles  Find the number of moles in each of the following quantities:  0.370 g B  27.4 g TiO 2  847 g (NH 4 ) 2 CO 3

22. Conclusion  We should now know how to convert from moles to mass and from mass to moles.  Often times laboratory procedures describe chemical reactions in which the amounts of elements or compounds being added are represented in moles. Yet if we are adding a solid we must convert from moles to grams to add the correct amount of reactant so that a successful reaction takes place. Likewise, if we are adding a liquid we must convert to units of volume (usually mL) using density (D=M/V). We will go over how to convert from moles of a gas to volume of a gas in the rest of section 7.2.

23. Mole-Volume Relationships  Review: What are the units that describe molar mass?  What units do we use to describe the volume of a mole?

24. Mole-Volume Relationships  How does temperature affect the volume of a gas?  How does pressure affect the volume of a gas?  http://phet.colorado.edu/simulations/sims.php?sim=Gas_Properties http://phet.colorado.edu/simulations/sims.php?sim=Gas_Properties

25. Mole-Volume Relationships  Because of the fact that the volume of a gas is affected by temperature and pressure, we must measure gasses at STP (Standard Temperature and Pressure).  Standard temperature = 0ºC  Standard pressure = 1atm or 101.3 kPa

26. Mole-Volume Relationships  At STP one mole of any gas occupies a volume of 22.4 L. This quantity is known as the molar volume of a gas.  We can also say that because 1 mol (of gas) = 22.4L (of gas) and 1 mol = 6.02x10 23 representative particles, then 22.4L of a gas = 6.02x10 23 representative particles of that gas.

27. Moles to Liters  What is the volume at STP of these gases?  3.20x10 -3 mol CO 2  0.960 mol CH 4  3.70 mol N 2

28. Moles to Liters  Assuming STP, how many moles are in these volumes?  67.2 L SO 2  0.880 L He  1.00x10 3 L C 2 H 6