Class Powerpoint Unit 4

1. Unit 4
The Mole

2. Unit 4
Measuring Matter

3. Warm-Up
Identify terms we
use everyday that
describe a quantity.
(at least 3)

4. Warm-Up
What is the common
meaning of a mole?
What is the scientific
meaning of a mole?

5. Do you ask for 12? Do you ask for 24?
OR do you ask for a DOZEN???
Counting particles

6.  Much easier for the
product we are buying
if it is sold in UNITS.

7. The Mole is…
Abbreviated mol
SI unit, used to measure the
amount of a substance
Defined as the number of
carbon atoms in exactly 12 g of
pure carbon-12

8. Common meaning
Scientific meaning
A small burrowing An SI base unit used to
animal. measure the quantity of
matter.
The damage to the
lawn was caused by a The chemist measured out a
mole. mole of the compound.

9. Avogadro’s Number
 Thenumber
6.0221367 x 10 23
 Avogadro was a
Italian physicist.

10. Avogadro’s Number
Avogadro's number is rounded to
three significant figures.
Is
used to measure extremely small
particle such as atoms.

11. Converting between Moles
and Particles
Suppose you buy 3 ½ dozen of roses
and you want to know how many
roses you have.
Example Calculation (Work problem out)
Describe how you can tell if the wrong
conversion factors has been used.

12. Moles to particles
How many particles are
in 3.50 mol of sucrose?

13. Reinforcement Problems
1. Zinc (Zn) is used to form a corrosion-
inhibiting surface on galvanized steel.
Determine the number of Zinc atoms in
2.50 mol Zn?
2. Calculate the number of molecules in
11.5 mol of water (H2O).
3. Challenge Calculate the number of
oxygen atom in 5.0 mol of oxygen
molecules. Oxygen is a diatomic
molecule O2

14. Particles to Moles
 Youcan convert between moles and
number of representative particles by
multiplying the known quantity by the
proper conversion factor.

15. Reinforcement
How many moles are in
each of the following?
a. 5.75 x 1024 atoms Al
b. 2.50 x 1024 atoms Fe

16. Unit 4
Measuring matter worksheet

17. Unit 4
Mass and the Mole

18. Calculate the number of
molecules in 5.67 mol of water
(H2O).

19. Mass and the Mole
Amole always contain the same
number of particles; however, moles
of different substances have different
masses.

20. Real world application
 Youwould not expect a dozen limes to
have the same mass as a dozen eggs.
WHY?
Eggs and Limes differ in size and composition.
One mole quantities has different masses for
the same reasons.

21. Molar Mass
The atomic weight of and
element expressed in
grams is the mass of one
mole of that element.
Unit g/mol

22. Apply
What of the mass of one
mole of copper?
63.546 g

23. Using molar mass-
Real World Application
 Imagine that we buy jelly beans
bulk, and sell them by the dozen.
We soon realize that counting out a
dozen or 10 dozen every time
someone orders some is to much
work, What can we do instead?

24. Using molar mass-
Real World Application
 The mass of…
1 dozen Jelly Beans = 35 g of Jelly Beans
Conversion Factors= ??

25. Using molar mass-
Real World Application
What mass of Jelly beans
would you measure if a
customer want 5 dozen jelly
beans?
175 g Jelly beans

26. Moles to Mass
Suppose you need to measure
out 3.00 mol of copper (Cu).
Howwould you measure that
amount? 191 g Cu
Suppose you want to go
grams back to moles?

27. Mole-to-Mass Conversion
1. Determine the mass in grams of each of
the following.
a. 3.57 mol Al
b. 42.6 mol Si
2. Challenge Convert each given quantity
in scientific notation to mass in grams
expressed in scientific notation.
a. 3.45 x 102 mol Co
b. 2.45 x 10-2 mol Zn

28. Mass-to-Mole
1. Determine the number of moles in each
of the following.
a. 25.5 g Ag
b. 300.0 g S

29. Unit 4
Mass and the Mole
Mass-to-atoms

30. Converting between mass and atoms
Back to the jellybean
example… at the end of the
day you have 550 g of jelly
beans left over, how many jelly
beans is that?
Remember 35 g= 1 dozen

31. Steps to follow
1. grams of jelly
beans 
dozens of jelly
beans
2. Dozens of jelly
beans  # of
jelly beans

32. Reinforcement
Mass to atoms/ Atoms to
mass Worksheet

33. Unit 4
Re-cap

34. Activity
Ona blank sheet of
computer paper design a
concept map, outlining the
conversion between mass,
moles, and particles

36. Warm-Up
Calculate the mass of
0.25 mol of carbon-12

37. Unit 4
Moles of a compound

39. Chemical Formulas & The MOLE
 Chemical formula=indicates the
number and types of atoms
contained in one unit of the
compound.

40. Interpret
How many of each kind of
atom- carbon, chlorine, and
fluorine-are contained in 1 mol
of CCl2F2 ?
C=1
Cl=2
F=2

41. Interpret
 How many moles of F are in
5.50 of freon (CCl2F2 ) ?
11.0 mol F atoms

42. Molar mass of Compounds
Calculate the molar mass of K2CrO4.
194.20 g

43. Reinforcement
1. Determine the molar mass of each ionic
compound.
a. NaOH
b. CaCl2
c. KC2H3O2
2. Calculate the molar mass of each
molecular compound.
a. C2H5OH
b. HCN

44. Converting Moles to mass
 How many grams of F
are in 11.0 mol F ?

45. Conversion Factors
 Conversion factors map

46. Unit 4
Measuring matter Lab

47. Warm-Up
Suppose a student measured out
8.91 g of Al2O3 Aluminum oxide…
(answer the following)
Molar mass of Al2O3
Moles of compound
Moles of each element
Atoms of each element

48. Warm-Up
 Suppose a student measured out 81.2
g of CaCO3 Calcium Carbonate…
(answer the following)
 Molar mass of
 Moles of compound
 Moles of each element
 Atoms of each element

49. Warm-Up
Suppose a student measured out
81.2 g of CaCO3 Calcium
Carbonate… (answer the
following)
Molar mass of
Moles of compound
Moles of each element

50. Measuring Matter Lab
Safety-goggles, gloves
Proper use of equipment
Procedure
Expectation –Clean Lab
area
Questions?

51. After Lab
Complete analysis
Work on review/homework
packet
Study for Quiz

52. Unit 4
Chalk Lab

53. Warm-Up
Find the molar mass
of calcium
carbonate (CaCO3)

54. Lab Expectations
Findthe following…
How many moles of chalk
How many mole of Ca… How many
atoms
How many moles of C… How many
atoms
How many moles of O… How many
atoms

55. Lab Procedure
Find mass before-record
Find mass after-record
Find molar mass
Determine the number
of mole required

56. Unit 4
Mole & the volume

57. Warm-Up
How much does 4.2
moles of Ca(NO3)2
weigh?

58. Quiz

59. Unit 4
Mole & the volume

60. Moles to Volume
In the calculations, some questions
should use conversions more than
once; however, mole is in the center
, because all conversions between
mass, particles and volume should
be converted to moles at first.

62. Moles and volume

63. Unit 4
Empirical and Molecular
formulas

64. Real World Application
 You might have noticed that some food
containers contain two or more serving
instead of the single serving you would
expect. How would you determine the
total number of calories contained in the
package?

66. Real World Application
 What do you do??

67. Percent by mass
Percent by mass: the ratio of
the mass of each element to
the total mass of compound
expressed as a percent.

69. Percent Composition
Percent by mass (element) =
mass of element
-------------------------- X100 =
mass of compound

70. Sample Problem
 You have a 100g sample of a compound
that contain 55 g of element X and 45g of
element Y.
 The compound is 55% X and 45% Y

71. The percent composition from
the Chemical Formula
Percent by mass=
mass of element in 1 mol of compound
-------------------------------------------------------------------- X 100
molar mass of compound

72. Unit 4
Empirical formula from mass
composition

73. Empirical formula
 ISTHE FORMULA WITH THE SMALLEST
WHOLE-NUMBER MOLE RATIO OF THE
ELEMENTS.
 The empirical formula might or might not
be the same as the actual molecular
formula, of different the molecular
formula will be a simple multiple of the
empirical formula.

74. Empirical formula
For example:
Hydrogen peroxide
Empirical formula= HO
Molecular formula= H2O2

75. How do you find empirical
formula?
 Percent
composition or masses of the
elements in a given mass of compound

76. List the steps needed to calculate
the empirical formula from
percent composition data…
1. Assume that the total mass of the
compound is 100.00 g the percent by
mass of each element is equal to the
mass of each element in grams
2. Convert the mass of each element to
moles using the molar mass as a
conversion factor.

77. List the steps needed to calculate
the empirical formula from
percent composition data…
3. Divide each molar amount by the
smallest mole value
4. If needed multiply each by an interger to
determine the smallest whole-number
ratio
5. Write the empirical formula using the
smallest whole-number ratio.

78. Reinforcement
 Empiricalformula from percent composition
Methyl acetate is a solvent commonly used in
some paints, inks, and adhesives. Determine the
empirical formula for methyl acetate, which has
the following chemical analysis:
 48.64% carbon
 8.16% hydrogen
 43.20% Oxygen

79. 1. Assume that each percent by mass represents
the mass of the element in 100g sample, the
percent sign CAN BE REPLACED WITH
GRAMs(g).
2. Convert each mass to moles using conversion
factors (inverse molar mass- that relates moles
to grams)
3. Calculate the simplest ratio of moles of
elements by dividing the moles of each
element by the SMALLEST value in the
calculated mole ratio.
4. Multiply each number in the ratio by the
SMALLEST number

80. Reinforcement
 Empiricalformula from percent composition
Methyl acetate is a solvent commonly used in
some paints, inks, and adhesives. Determine the
empirical formula for methyl acetate, which has
the following chemical analysis:
 48.64% carbon
 8.16% hydrogen
 43.20% Oxygen

81. Unit 4
Molecular formula

82. Molecular formula
 Specifiesthe actual number of atoms of
each element in one molecular formula
unit of the substance