2. The Mole Chapter 11 – Chemistry L1 LSM High School 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.

3. 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!!!

4. What are MOLES? Moles are defined as the number of carbon atoms in exactly 12 grams of the carbon-12 isotope. 1 mole of _____ = 6.02 x 10 23 particles Mole: unit = “mol” Avogadro’s number dozen, baker’s dozen, pi

5. 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

6. 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

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

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

9. Practice Problem 2 How many molecules of CO 2 are there in 4.56 moles of CO 2 ? K:UK: Answer: 2.75 x 10 24 molecules of CO 2 How many atoms is this?

10. 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

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

12. Section 11.2: Mass and the Mole Objectives: 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.

13. Let’s Look at the Periodic Table! 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. Atomic Mass - usually increase across a row Why do they have decimal values? The atomic mass (sometimes called average atomic mass) is the weighted average of the masses of all the naturally occurring isotopes of that element. A relative scale: Uses 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, meaning 1 atom of hydrogen-1 is one-twelfth the mass of one atom of carbon-12

14. The mass in grams of one mole of ANY pure substance is its molar mass. Same value as atomic mass - has units of g/mol Occasionally referred to as: Gram atomic mass (gam) – for atoms Gram molecular mass (gmm) – for molecules Gram formula mass (gfm) – for formula units (ionic compounds) 12.01 grams of carbon has the same number of particles as 1.01 grams of hydrogen and 55.85 grams of iron. The molar mass is found in the periodic table! Avogadro’s number tells us the number of particles.

15. Using Molar Mass as a Conversion Factor: # of grams or 1mol 1 mol # of grams Practice Problems: 1.What is the mass, in grams, of 2.34 moles of carbon? K:UK: 28.1 g carbon

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

17. Section 11.3: Moles of Compounds Objectives: Recognize the mole relationships shown by a chemical formula. Calculate the molar mass of a compound. Calculate the number of moles of a compound from a given mass of the compound, and the mass of a compound from a given number of moles of the compound. Determine the number of atoms or ions in a mass of a compound.

18. Enough about atoms: What about compounds?  The chemical formula for a compound tells us the types of elements and the number of each element contained in one unit of the compound. Ammonia (NH 3 ) 1 molecule contains:  1 atom of nitrogen and 3 atoms of hydrogen Baking soda (sodium hydrogen carbonate, NaHCO 3 ) 1 formula unit contains:  1 atom of sodium, 1 atom of hydrogen, 1 atom of carbon, and 3 atoms of oxygen

19. Example Problems: 1. Calculate the number of moles of hydrogen found in 3.50 moles of NH 3. K:UK: 10.5 mol Hydrogen

20. 2. Calculate the number of moles of carbon found in 9.85 moles of C 6 H 12 O 6 (sugar). K:UK: 59.1 mol carbon

21. The Molar Mass of Compounds

22. Summary of Getting the Molar Mass of Compounds: The mass of a mole of a compound equals the sum of the masses of every particle that makes up the compound. Use the formula to tell you how many of each element that is in the compound Use the periodic table to get the masses of each element Add them all up and you get the molar mass of the compound in units of g/mol (NH 4 ) 2 SO 4

23. Example Problems of Molar Masses: 1) What is the molar mass of NH 3 ? 17.04 g/mol 2) What is the molar mass of Sr(NO 3 ) 2 ? 211.64 g/mol

24. Example Problems of Mole- Mass Conversions: 1) How many moles is 4.56 g of CO 2 ? K:UK: 0.104 moles CO 2 2) How many moles is 46.8 g of CH 4 ? K:UK: 2.92 mol CH 4

25. 3) How many grams is 9.87 moles of H 2 O? K:UK: 178g H 2 O 4) How many grams is 0.157 mol Fe 2 O 3 ? K:UK: 25.0 g Fe 2 O 3

26. Using Molar Volume as a Conversion Factor:  Molar Volume: for any gas at STP, 1 mol = 22.4 L  Standard Temperature and Pressure is 0°C or 273 K and 101.3 kPa or 1 atm 22.4 L or 1mol 1 mol 22.4 L Practice Problem: What is the volume of 1.5 moles of nitrogen gas? K:UK: 34 L N 2 (17 2-Liter bottles!)

27. REVIEW: What types of particles are contained in covalent compounds? What types of particles are contained in ionic compounds?

28. Multi-step Conversions You must first convert to moles and then convert to the desired unit either using molar mass or Avogadro’s number or molar volume.

29. Example Problems: 1. What is the volume of 45.6 g of water vapor? K:UK: ? L H 2 O (g)

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

31. 4. What is the mass, in grams, of 1.50 x 10 15 atoms uranium? K:UK: 5.93 x 10 -7 g U

32. 3. What is the mass, in grams, of 1.50 x 10 15 formula units of NaCl? K:UK: ? g NaCl

33. More Example Problems… 1) How many molecules in 6.8 g of CH 4 ? 2a) How many formula units are there in 4.9 g of NaNO 3 ? 2b) How many ions if the compound is made of Na + and NO 3 - ions?