2. The Mole Chapter 11 – Honors Chemistry LSM High School

3. Section 11.1: Measuring Matter Objectives: Describe how a mole is used in chemistry Relate a mole to common counting units Covert moles to number of representative particles and number of representative particles to moles.

4. How do Chemists measure how much of a substance? Chemists can measure mass or volume or they can count pieces. Chemists can measure mass in grams. Chemists can measure volume in liters. MOLES Chemists can count pieces in MOLES. No, not that kind of mole!!!

5. What are MOLES? Moles are defined as the number of carbon atoms in exactly 12 grams of the carbon-12 isotope. 1 mole is 6.02 x 10 23 particles. Treat it like a very large dozen 6.02 x 10 23 is called Avogadro's number.

6. A Little History Amedeo Avogadro was born in 1776 in Turin, Italy. He went on to study molecular theory and helped other scientists distinguish between atoms and molecules. Because of his accomplishments in this field, the variable that tells the number of molecules in one mole was named after him

7. What about units on Avogadro’s number? The units of Avogadro’s number can be whatever particle you are counting. Examples: atoms, molecules, ions, etc… In chemistry these are called Representative Particles

8. What are Representative Particles? These particles are the smallest pieces of a substance. The types of representative particles that chemists generally work with are: atoms – the smallest particle of an element ions – atoms with positive or negative charges molecules – two or more covalently bonded atoms formula units – the simplest ratio of ions that make up an ionic compound

9. How Do We Use Moles? Moles are used as conversion factors. This means they are used to change units. Remember, when solving using conversion factors there are 3 questions you want to ask yourself: What unit do you want to get rid of? Where does it go to cancel out? What can you change it into?

10. Converting Moles to Particles and Particles to Moles Using Avogadro’s Number as a Conversion Factor

11. How do we write Avogadro’s number as a conversion factor? 6.02 x 10 23 particles or1 mole 1 mole6.02 x 10 23 particles

12. Practice Problem 1 How many atoms are in 2.50 mol of zinc? K:UK: Answer: 1.51 x 10 24 atoms Zn

13. Practice Problem 2 How many molecules of CO 2 are the in 4.56 moles of CO 2 ? K:UK: Answer: 2.75 x 10 24 molecules of CO 2

14. Practice Problem 3 How many moles of water is 5.87 x 10 22 molecules of water? K:UK: ANSWER: 0.0975 moles of water

15. Practice Problem 4 Given 3.25 mol AgNO 3, determine the number of formula units. K:UK: ANSWER: 1.96 x 10 24 formual units AgNO 3

16. Section 11.2: Mass and the Mole Relate the mass of an atom to the mass of a mole of atoms. Calculate the number of moles in a given mass of an element and the mass of a given number of moles of an element. Calculate the number of moles of an element when given the number of atoms of an element. Calculate the number of atoms of an element when given the number of moles of the element.

17. Let’s Look at the Periodic Table! What are some patterns that you see on the chart?

18. Atomic Numbers Always increase across a row. The atomic number is the number of protons in an atom of that element. This number identifies it as an atom of a particular element.

19. Atomic mass Usually increase across a row – but not always. Ex: Why do they have decimal values? The atomic mass (sometimes called average atomic mass) is the weighted average mass of the isotopes of that element. You will be doing a series of activities to better understand how atomic masses and Avogadro’s number was determined.

20. How Atomic Masses and the Mass of a Mole are Related: Atomic masses are a relative scale. Use isotope carbon-12 as the standard Each atom of carbon-12 has a mass of exactly 12 amu (atomic mass units) Ex: One atom of hydrogen-1 has a mass of 1 amu, that means that 1 atom of hydrogen-1 is one-twelfth the mass of one atom of carbon-12 Atomic masses are on the periodic table but they are not whole numbers b/c the values are weighted averages of the masses of all the naturally occurring isotopes of each element

21. Since mole is the number of representative particles, or atoms, in exactly 12 g of pure carbon-12, then the mass of one mole of carbon-12 atoms is 12 g. The mass in grams of one mole of ANY pure substance is its molar mass. Same value of atomic mass but has units of g/mol 12.01 grams of carbon has the same number of particles as 1.01 grams of hydrogen and 55.85 grams of iron. The number of particles is 6.02 x 10 23 atoms. Therefore, we can count things by weighing them.

22. Using Molar Mass: The molar mass can be used as a conversion factor. It relates the mass of a substance to the number of moles of that substance. To calculate the mass from the number of moles, you would use the molar mass as: # of grams 1 mole To calculate the moles from the mass, you would use the molar mass as: 1mol # of grams

23. Example Problems: 1. What is the mass, in grams, of 2.34 moles of carbon? K:UK: 28.1 g carbon 2.34 moles C 12.01 g/moles C grams C

24. 2. How many moles of magnesium are in 4.61g of Mg? K:UK: 0.190 mol Mg

25. Conversions from mass to atoms and atoms to mass There is no direct conversion between the mass of a substance to the number of representative particles of that substance. You must first convert to moles and then convert to the desired unit either using molar mass or Avogadro’s number.

26. Example Problems: 1. How many atoms are in 45.6 g Si? K:UK: 9.77 x 10 23 atoms Si

27. 2. How many atoms are in 0.120 kg Ti? K:UK: 1.51 x 10 24 atoms Ti

28. 3. What is the mass, in grams, of 1.50 x 10 15 atoms N? K:UK: 3.49 x 10 -8 g N

29. 4. What is the mass, in grams, of 1.50 x 1015 atoms uranium? K:UK: 5.93 x 10 -7 g U