1. The Mole
Ch 11

2. Measuring matter
11.1

3. 11.1 Vocabulary
Review
molecule: two or more atoms that covalently bond together to form a unit
New
mole
Avogadro’s number
Main Idea - Chemists use the mole to count atoms, molecules, ions, and formula units.
NOTE – YOU WILL NEED A SCIENTIFIC CALCULATOR FOR THIS CHAPTER!

4. How do we measure items? You can measure mass, distance, or volume
You can count pieces.
We measure mass in grams.
We measure distance in meters.
We measure volume in liters.
We count pieces in MOLES.

5. We’re not talking about this kind of mole!
What is the mole?
We’re not talking about this kind of mole!

6. Counting Particles
Chemists need a convenient method for accurately counting the number of atoms, molecules, or formula units of a substance.
The mole is the SI base unit used to measure the amount of a substance.

7. Moles (is abbreviated: mol)
It is an amount, defined as the number of carbon atoms in exactly 12 grams of carbon-12.
1 mole = x of the representative particles.
Treat it like a very large dozen!
6.022 x is called Avogadro’s number.

8. Similar Words for an amount
Pair: 1 pair of shoelaces
= 2 shoelaces
Dozen: 1 dozen oranges
= 12 oranges
Case: 1 case of Dr. Pepper
= 24 cans Dr. Pepper
Gross: 1 gross of pencils
= 144 pencils
Ream: 1 ream of paper
= 500 sheets of paper

9. What are Representative Particles?
The smallest pieces of a substance:
For a molecular compound: it is the molecule.
For an ionic compound: it is the formula unit (made of ions).
For an element: it is the atom.
Remember the 7 diatomic elements? (made of molecules)

10. Types of questions How many oxygen atoms in the following? CaCO3
Al2(SO4)3
How many ions in the following?
CaCl2
NaOH
3 atoms of oxygen
12 (3 x 4) atoms of oxygen
3 total ions (1 Ca2+ ion and 2 Cl1- ions)
2 total ions (1 Na1+ ion and 1 OH1- ion)
5 total ions (2 Al SO42- ions)

11. Converting Between Moles and Particles
Conversion factors must be used.
Moles to particles
Example:
Number of molecules in 3.50 mol of sucrose

12. Converting Between Moles and Particles (cont.)
Particles to moles
Use the inverse of Avogadro’s number as the conversion factor.

13. Practice problems (round to 3 sig. figs.)
How many molecules of CO2 are in 4.56 moles of CO2?
How many moles of water is 5.87 x molecules?
How many atoms of carbon are in 1.23 moles of C6H12O6?
How many moles is 7.78 x 1024 formula units of MgCl2?
2.75 x 1024 molecules
mol (or 9.75 x 10-2)
4.44 x 1024 atoms C
12.9 moles

14. 10.1 Check What does the mole measure? A. mass of a substance
B. amount of a substance
C. volume of a gas
D. density of a gas

15. 10.1 Check
What is the conversion factor for determining the number of moles of a substance from a known number of particles? A. B.
C. 1 particle  6.02  1023
D. 1 mol  6.02  1023 particles

16. Mass and the mole
11.2

17. 11.2 Vocabulary
Review
conversion factor: a ratio of equivalent values used to express the same quantity in different units
New
molar mass
Main Idea - A mole always contains the same number of particles; however, moles of different substances have different masses.

18. The Mass of a Mole
1 mol of copper and 1 mol of carbon have different masses.
One copper atom has a different mass than 1 carbon atom.

19. Measuring Moles Remember relative atomic mass?
The amu was one twelfth the mass of a carbon-12 atom.
Since the mole is the number of atoms in 12 grams of carbon-12,
the decimal number on the periodic table is also the mass of 1 mole of those atoms in grams.

20. The Mass of a Mole (cont.)
Molar mass (MM) is the mass in grams of one mole of any pure substance.
Also called Formula Weight (FW)

21. Molar Mass
Equals the mass of 1 mole of an element in grams (from periodic table)
grams of C has the same number of pieces as
1.008 grams of H
55.85 grams of iron.
We can write this as: g C = 1 mole C
We can count things by weighing them.

22. Using Molar Mass
Moles to mass

23. Using Molar Mass (cont.)
Convert mass to moles with the inverse molar mass conversion factor.

24. Examples How much would 2.34 moles of carbon weigh?
How many moles of magnesium is g of Mg?
How many atoms of lithium is 1.00 g of Li?
How much would 3.45 x 1022 atoms of U weigh?
28.1 grams C
1 mol Mg
8.72 x 1022 atoms Li
13.6 grams U

25. Using Molar Mass (cont.)
This figure shows the steps to complete conversions between mass and atoms.

26. 11.2 Check
The mass in grams of 1 mol of any pure substance is: A. molar mass B. Avogadro’s number C. atomic mass D. 1 g/mol

27. 11.2 Check Molar mass is used to convert what? A. mass to moles
B. moles to mass
C. atomic weight
D. particles

28. Moles of compounds
11.3

29. Vocabulary
Review
representative particle: an atom, molecule, formula unit, or ion
Main Idea -The molar mass of a compound can be calculated from its chemical formula and can be used to convert from mass to moles of that compound.

30. Chemical Formulas and the Mole
Chemical formulas indicate the numbers and types of atoms contained in one unit of the compound.
One mole of CCl2F2 contains one mole of C atoms, two moles of Cl atoms, and two moles of F atoms.

31. The Molar Mass of Compounds
The molar mass of a compound equals the molar mass of each element, multiplied by the moles of that element in the chemical formula, added together.
The molar mass of a compound demonstrates the law of conservation of mass.

32. The Molar Mass of Compounds
in 1 mole of H2O molecules there are two moles of H atoms and 1 mole of O atoms
To find the mass of one mole of a compound
determine the number of moles of the elements present
Multiply the number times their mass (from the periodic table)
add them up for the total mass

33. Calculating Molar Mass
Calculate the molar mass of magnesium carbonate, MgCO3.
g g + 3 x ( g) =
Thus, grams is the formula mass for MgCO3.

34. Examples Na2S N2O4 C Ca(NO3)2 C6H12O6 (NH4)3PO4
Calculate the molar mass of the following and tell what type it is:
Na2S
N2O4 C
Ca(NO3)2
C6H12O6
(NH4)3PO4
= g/mol
= g/mol
= g/mol
= g/mol
= g/mol
= g/mol

35. Moles to Mass Conversion for Compounds
For elements, the conversion factor is the molar mass of the elements.
The procedure is the same for compounds, except that you must first calculate the molar mass of the compound.

36. Mass to Moles Conversion for Compounds
The conversion factor is the inverse of the molar mass of the compound.

37. For example 5.69 g NaOH 1 mol NaOH 39.997 g NaOH
How many moles is 5.69 g of NaOH?
1 mole NaOH = g NaOH
5.69 g NaOH
1 mol NaOH
g NaOH
= mol NaOH
= g NaOH 3 Sig Figs

38. Mass to Particles Conversion for Compounds
Convert mass to moles of compound with the inverse of molar mass.
Convert moles to particles with Avogadro’s number.
This figure summarizes the conversions between mass, moles, and particles.

39. 11.3 Check How many moles of OH— ions are in 2.50 moles of Ca(OH)2?
B. 2.50
C. 4.00
D. 5.00

40. 11.3 Check
How many particles of Mg are in 10 moles of MgBr2? A  1023 B  1024 C  1024 D  1025

41. Empirical and Molecular formulas
11.4

42. Vocabulary Review New percent composition empirical formula
percent by mass: the ratio of the mass of each element to the total mass of the compound expressed as a percent
New
percent composition
empirical formula
molecular formula
Main Idea - A molecular formula of a compound is a whole-number multiple of its empirical formula.

43. Percent Composition
The percent by mass of any element in a compound can be found by dividing the mass of the element by the mass of the compound and multiplying by 100.

44. Percent Composition
The percent by mass of each element in a compound is the percent composition of a compound.
Percent composition of a compound can also be determined from its chemical formula.

45. Calculating Percent Composition of a Compound
Like all percent problems:
part whole
x 100 % = percent
Find the mass of each of the
elements.
Next, divide by the total mass of the compound; then x 100%

46. Example
Calculate the percent composition of a compound that is made of 29.0 grams of Ag with 4.30 grams of S
(Assume you have one mol of substance)
29.0 g Ag
X 100 = 87.1 % Ag
33.3 g total
Total = 100 %
4.30 g S
X 100 = 12.9 % S
33.3 g total

47. Examples Calculate the percent composition of C2H4?
How about Aluminum carbonate?
85.7% C, 14.3 % H
23.1% Al, 15.4% C, and 61.5 % O

48. Empirical Formula
The empirical formula for a compound is the smallest whole-number mole ratio of the elements.
To calculate the empirical formula from percent by mass:
1) Assume you have g of the compound.
2) Then convert the mass of each element to moles.

49. Empircal vs Molecular
The empirical formula may or may not be the same as the molecular formula.
Molecular formula of hydrogen peroxide = H2O2
Empirical formula of hydrogen peroxide = HO

50. Molecular Formula
The molecular formula specifies the actual number of atoms of each element in one molecule or formula unit of the substance.
Molecular formula is always a whole-number multiple of the empirical formula.

51. Empirical vs. Molecular Formula
What is the empirical formula for the compound C6H12O6?
A. CHO
B. C2H3O2
C. CH2O
D. CH3O
Which is the empirical formula for hydrogen peroxide?
A. H2O2
B. H2O
C. HO
D. none of the above

52. Formulas of hydrates
11.5

53. Vocabulary
Review
crystal lattice: a three-dimensional geometric arrangement of particles
New
hydrate
Main Idea -Hydrates are solid ionic compounds in which water molecules are trapped.

54. Naming Hydrates
A hydrate is a compound that has a specific number of water molecules bound to its atoms.
The number of water molecules associated with each formula unit of the compound is written following a dot.
Example
Sodium carbonate decahydrate
= Na2CO3 • 10H2O

55. Naming Hydrates

56. Analyzing Hydrates
When heated, water molecules are released from a hydrate leaving an anhydrous compound.
To determine the formula of a hydrate, find the number of moles of water associated with 1 mole of hydrate.

57. Analyzing Hydrates (cont.)
1. Weigh hydrate.
2. Heat to drive off the water.
3. Weigh the anhydrous compound.
4. Subtract and convert the difference to moles.
The ratio of moles of water to moles of anhydrous compound is the coefficient for water in the hydrate.

58. Use of Hydrates
Anhydrous forms of hydrates are often used to absorb water, particularly during shipment of electronic and optical equipment.
In chemistry labs, anhydrous forms of hydrates are used to remove moisture from the air and keep other substances dry.

59. 11.5 Check
Heating a hydrate causes what to happen? A. Water is driven from the hydrate. B. The hydrate melts. C. The hydrate conducts electricity. D. There is no change in the hydrate.

60. 11.5 Check
A hydrate that has been heated and the water driven off is called: A. dehydrated compound B. antihydrated compound C. anhydrous compound D. hydrous compound

61. 11.5 Check
Two substances have the same percent by mass composition, but very different properties. They must have the same ____. A. density B. empirical formula C. molecular formula D. molar mass

62. Empirical and Molecular formulas
11.4

63. Formulas of hydrates
11.5