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Ch 11: The Mole.

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Presentation on theme: "Ch 11: The Mole."— Presentation transcript:

1 Ch 11: The Mole

2 Section 11.1 Measuring Matter
Dozen= 12 Ream= 500 Pair= 2 Gross= 144

3 What is a mole? It is the SI base unit used to measure the amount of a substance Abbreviated mol Represents particles Is also called Avogadro’s number 6.02 x 1023 (3 significant figures)

4 Remember conversion factors?
12 roses = 1 dozen So, 3.5 dozen= ? Roses = 42 Well, 6.02 x 1023 atoms (or any representative paticles) = 1 mol

5 Review Scientific Notation
Review Rounding

6 Work practice problems 1-3
Example: 3.50 mol sucrose has how many molecules? Work practice problems 1-3 How many moles are in 3.58 X 1020 atoms of Ca? Work practice problems 4-7

7 closure How is a mole similar to a dozen?
What is the relationship between avagadro’s number & one mole? Why do chemists use moles? **worksheet: The mole & Avogadro’s number STOP

8 Section Mass & the Mole Different particles (atoms) have different masses. Remember atomic mass. Each element has its own specific mass. Therefore each compound has its own specific mass.

9 Molar mass (g/mol)-mass in grams of any pure substance
The molar mass of any element is numerically equal to its atomic mass. Thus…1 mol Mn = g/mol Mn = 6.02 x 1023 atoms Mn

10 Example: 1 mol Zn = 65.4 g/mol 1 mol O2 = 32.0 g/mol

11 Practice 3 mol Zinc = 196 g/mol Zn 1 mol H2O = 18.0g/mol H2O 1 mol sulfuric acid = 98.1 g/mol H2SO4

12 Mass  Mole conversions
Example: Calculate the mass in grams of mol Cr. Work Practice problems 1-4

13 Hydrate: CuSO4·5H2O 4) Calculate the mass in grams of 2.45 mol of CaCl2·2H2O

14 Mole Mass conversions
Example: Determine the number of moles for 25.5 g Ag. (mass mol) Work practice problems 5-7. **worksheet: moles & mass

15 Mass  Atom Conversion Example:
How many atoms of gold (Au) are in a pure nugget having the mass of 25.0g? Practice problems 1-5

16 Atom  Mass Conversion Example
A party balloon contains 5.50 x 1022 atoms of helium (He) gas. What is the mass in grams of the helium? Practice problems 6-10.

17 Mole  Mole Example According to the following balanced equation, how many moles of O2 is produced from 3.00 moles of CuO? 2CuO  2Cu + O2 Practice problems 1 & 2.

18 Section 11.4 Empirical & Molecular Formula
Percent Composition is the percent by mass of each element in a compound. Example: If we had 100 g of a sample of some new compound contains 55g of element X & 45 g of element Y. What is the % of element X & Y?

19 Mass of element in 1 mol of compound X 100
If we already know the _chemical formula for a compound, you can calculate its percent composition. % by mass= Mass of element in 1 mol of compound X 100 Molar mass of compound

20 Ex. Determine the percent composition of H2O.
(If you had 350. g of water, then how much is oxygen?) Practice problems 1-3. 311g

21 Empirical Formula is the smallest whole number ratio of the elements.
Calculating Empirical formula from percent composition: Directions: The percent should be assumed to be 100 g & converted to moles. Then we figure out the mole ratio by dividing each by the smallest mole.

22 POEM: Percent to Mass Mass to Mole Divide by small Multiply til whole.

23 Example 1 The percent composition of an oxide of sulfur is 40.05% S & 59.95% O. Example 2 Determine the empirical formula for methyl acetate which has the following chemical analysis:48.64% C, 8.16% H, & 43.20% O. Practice Problems

24 **worksheet: determining empirical formulas

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