1. THE MOLE

2. What is the Mole?
6.02 x 1023 OR

3. What is a dozen?
1 dozen = 12

4. Why don’t we just say 12?
It is often easier to group objects when counting them 2 dozen = dozen = 36
…and it doesn’t matter what the objects are… a dozen means 12 whether it is a dozen eggs, a dozen roses, or a dozen ugly hats

5. So what is a mole?
A mole is a counting unit. It tells us how many objects there are.
A pair and a dozen are also counting units.
Pair
means
2 objects
Dozen
12 objects
Mole
6.02x1023 objects

6. So what is a mole? It is just a term used to represent a number
2 = pair = 1 dozen x 1023 = 1 mole
…and we call 6.02 x Avagadro’s number

7. Avogadro’s number = 6.02 x =
Avogadro -
He didn’t actually identify the number
– but it was named after him to honor his many contributions to chemistry

8. So why is it such a big number?
What counting unit would you use for these objects?
PAIR

9. So why is it such a big number?
What counting unit would you use for these objects?
DOZEN

10. So why is it such a big number?
What counting unit would you use for these objects?
MOLE
Since atoms and molecules are very small, the quantity used to measure them needs to be a very large number

11. So why is it such a big number?
Since a mole is such a large number, it is useful only for counting very small objects like atoms and molecules. A mole of donuts is too many to eat.
Like a pair (2) or a dozen (12), the quantity of a mole (6.02 x 1023) never changes, but unlike a pair or a dozen, its value is very, very large

12. Avogadro’s Number as Conversion Factor
6.02 x 1023 particles
1 mole or
1 mole
Note that a particle could be an atom OR a molecule!

13. Practice Number of atoms in 0.500 mole of Al
a) Al atoms
b) x 1023 Al atoms
c) x 1023 Al atoms
a) Al atoms
b) x 1023 Al atoms
c) x 1023 Al atoms
0.500 mole Al | *1023 atoms =
|1 mole
2. Number of moles of S in 1.8 x 1024 S atoms
a) mole S atoms
b) mole S atoms
c) x 1048 mole S atoms
a) mole S atoms
b) mole S atoms
c) x 1048 mole S atoms
1.8 *1024 atoms S | 1 mole =
| *1023 atoms

14. REPRESENTATIVE PARTICLES & MOLES
ATOMIC NITROGEN
ATOM N
6.02x1023
NITROGEN GAS
MOLEC. N2
WATER
H20
CALCIUM ION
ION
Ca2+
CALCIUM FLUORIDE
FORMULA UNIT
CaF2
CHEMICAL
REP. PARTICLE
SUBTANCE
REPRESENTATIVE
PARTICLE
FORMULA
IN 1 MOL

15. What if I asked you how many atoms are in a mole of a compound?
you must know how many atoms are in a representative particle or cluster of the compound.
To do this you must know the chem formula
For example, each molecule of CO2 is composed of 1 C + 2 O’s = 3 atoms
1 mole of carbon dioxide molecules contains avogadro’s number of carbon dioxide molecules.
Thus a mole of CO2 contains three times avogadro’s number of atoms

16. How many atoms of carbon are in 2.12 mols of propane molecules (C3H8)?
To find the # of atoms in a mol of a compound,
You 1st determine the # of atoms in a representative particle of that compound
And then multiply that # of atoms by avogadro’s #
How many atoms of carbon are in 2.12 mols of propane molecules (C3H8)?
Useful Info:
1 mol C3H8 = 6.02x1023 molecules C3H8
1molecule C3H8 = 3 atoms of C

17. (1molecC3H8)(X)=(1.276x1024molec•molsC3H8)
1st we need to see how many molecules of propane we have if we have 2.12 moles.
1 mole C3H8
2.12 moles C3H8 =
6.02x1023molec C3H8
X molec C3H8
(1molecC3H8)(X)=(1.276x1024molec•molsC3H8)
(1 molec C3H8)
(1 molec C3H8)
X = 1.276x1024 molecules of C3H8

18. (1molec C3H8)(X)=(3.83x1024molec • atoms C)
2nd we need to see how many atoms of C if we have 1.276x1024 molecs of C3H8.
1 molecule C3H8
1.276x1024 molec C3H8 =
3 atoms of C
X atoms C
(1molec C3H8)(X)=(3.83x1024molec • atoms C)
(1 molec C3H8)
(1 molec C3H8)
X = 3.83x1024 atoms of Carbon

19. So the mole can be used to measure any kind of particle – atoms, ions, molecules, formula units
It is possible to have a mole of iron atoms, a mole of water molecules and a mole of sodium chloride – all contain 6.02 x 1023 particles
But do 1 mol of iron, 1 mol of water and 1 mol of sodium chloride all have the same mass?
NO!!

20. Do a dozen donuts, a dozen books and a dozen bricks weigh the same?
NO, so neither do a mole of different atoms, ions or molecules

21. Each sample contains 1 mol (6.02 x 1023 atoms)
12.0g carbon
32.0g sulfur
65.38g zinc
Which element do you think has the heaviest atoms?

22. So now that we know what a mole of atoms looks like…where did the number come from?

23. CARBON!! 6.02 x 1023 atoms Guess how many there were?
Scientists took 12g of pure carbon-12
(composed of 6 protons + 6 neutrons; mass number = 12)
and counted how many carbon atoms
were in the sample
Guess how many there were?
6.02 x 1023 atoms
Why use 12 grams? They wanted the mass of one mole to be the same as the mass number

24. 6.02 x1023 atoms of Carbon
This pile of carbon-12 weighs 12.0 grams and contains 6.02 x 1023 carbon atoms
This is one mole of carbon

25. This pile of sulfur contains 6.02x1023 atoms of sulfur.
This is one mole of sulfur.
Review: A mole is a simple way of saying we
have 6.02x1023 objects.

26. Atomic molar mass of an element = is the average mass (g) of one mole of atoms of that element

27. Back to carbon….
12.01 g/mol
atomic molar mass of an element =
mass (g) of one mole of atoms of that element
Here is one mole of carbon which weighs 12.0 g
The atomic molar mass of carbon on the periodic table is….
Which means there are g for every mol of carbon

28. Here is 65. 38 grams of the element zinc
Here is grams of the element zinc. How many moles of zinc do we have?
1 mole .
atomic molar mass of an element = mass (g) of one mole of atoms of that element

29. What is the mass, in grams, of this sample of sulfur?
This pile of sulfur contains 6.02x1023 atoms of sulfur (or 1 mole of sulfur)
What is the mass, in grams, of this sample of sulfur?
32.06 g
atomic molar mass of an element = mass (g) of one mole of atoms of that element

30. Here is 0.50 moles of copper. How many grams of copper do we have?
0.50 moles x (63.55 g/mole) = 31.8g
0.50 moles | g= 31.8g | 1 mole)

31. This pile of powdered iron contains 3.01x1023 atoms. What is its mass?
27.9 g
3.01x1023 atoms| 1 mole Fe | g = 27.9g Fe | *1023 atoms Fe | 1mole Fe
3.01x1023 atoms| 1 mole = moles | 6.02 *1023 atoms moles | g Fe = 27.9 g Fe | 1 mole

32. Molar mass = the mass of one mole of a substance
So now we’re not talking about individual elements but molecules and compounds
I.e. What is the molar mass of water?
The formula for water is:
H2O , so one molecule contains hydrogen atoms and 1 oxygen atom

33. Molar mass = the mass of one mole of a substance
Atomic molar mass of H = 1.01g/mol x 2 = 2.02 g/mol
Atomic molar mass of O = 16.00g/mol x 1 = 16.00g/mol
Molar mass of water H2O is g/mol

34. Other Names Related to Molar Mass
Molecular Mass/Molecular Weight: If you have a single molecule, mass is measured in amu’s instead of grams. But, the molecular mass/weight is the same numerical value as 1 mole of molecules. Only the units are different. (This is the beauty of Avogadro’s Number!)
Formula Mass/Formula Weight: Same goes for compounds. But again, the numerical value is the same. Only the units are different.
THE POINT: You may hear all of these terms which mean the SAME NUMBER… just different units

35. Molar mass Now you try it… What is the molar mass of methane CH4(g) ?
C = g/mol x 1 = g/mol
H = 1.01g/mol x = 4.04 g/mol
Molar mass of methane = g/mol

36. Molar mass Now you try it… What is the molar mass of methanol CH3OH?
C = g/mol x 1 = g/mol
H = 1.01g/mol x 4 = g/mol
O = g/mol x = g/mol
Molar mass of methanol = g/mol

37. Find the molar mass of SnBr
Practice
Find the molar mass of SnBr
= g/mole
1 mole of Br atoms
1 mole of Sn atoms
1 mole of SnBr formula units
= g/mole
= g/mole

38. Converting between mass, molar mass and moles
m = quantity of matter (mass) in grams
n = quantity of matter in moles
M = molar mass
m = n x M
So if you wanted to know the mass of 2 mol of carbon
m = (2 mol) x (12.01g/mol) or
2 mol C | g C
| 1 mole C
= g C

39. Molar Mass of Molecules and Compounds
Mass in grams of 1 mole equal numerically to the sum of the atomic masses
1 mole of CaCl2 = g/mol
1 mole Ca x 40.1 g/mol
+ 2 moles Cl x 35.5 g/mol = g/mol CaCl2
1 mole of N2O4 = 92.0 g/mol

40. Practice Molar Mass of K2O = ? Grams/mole
B. Molar Mass of antacid Al(OH)3 = ? Grams/mole
Al= 1*26.98 =26.98
O =3 *16. = 48.0
H= 3*1.008 = g/mol

41. Practice
Prozac, C17H18F3NO, is a widely used antidepressant that inhibits the uptake of serotonin by the brain. Find its molar mass.

42. Converting Moles and Grams
Aluminum is often used for the structure of light-weight bicycle frames. How many grams of Al are in 3.00 moles of Al?
3.00 moles Al ? g Al

43. 1. Molar mass of Al 1 mole Al = 27.0 g Al
2. Conversion factors for Al
27.0g Al or mol Al
1 mol Al g Al
3. Setup 3.00 moles Al | g Al
1 mole Al
Answer = 3.00* 27 g Al = 81.0 g Al

44. Practice
The artificial sweetener aspartame (Nutra-Sweet) formula C14H18N2O5 is used to sweeten diet foods, coffee and soft drinks. How many moles of aspartame are present in 225 g of aspartame?

45. Atoms/Molecules and Grams
Since 6.02 X 1023 particles = 1 mole AND 1 mole = molar mass (grams)
You can convert atoms/molecules to moles and then moles to grams!
(Two step process)
You can’t go directly from atoms to grams!!!! You MUST go thru MOLES.
That’s like asking 2 dozen cookies weigh how many ounces if 1 cookie weighs 4 oz? You have to convert to dozen first!

46. Everything must go through Moles!!!
Calculations
molar mass Avogadro’s number Grams Moles particles
Everything must go through Moles!!!

47. Atoms/Molecules and Grams
How many atoms of Cu are present in 35.4 g of Cu?
35.4 g Cu mol Cu X 1023 atoms Cu g Cu mol Cu
= 3.4 X 1023 atoms Cu

48. Practice
How many atoms of K are present in 78.4 g of K?

49. Practice
What is the mass (in grams) of 1.20 X 1024 molecules of glucose (C6H12O6)?

50. Practice How many atoms of O are present in 78.1 g of oxygen?
78.1 g O2 1 mol O X 1023 molecules O2 2 atoms O g O mol O molecule O2

51. X = 12.2 L of CO 1 mole CO .500 mol CO = X L CO
*Remember*
1 mol CO = 22.4 L CO
1 mole CO
.500 mol CO =
22.4 L CO
X L CO
(1mol CO)(X L CO)= (.500 mol CO)(22.4 L CO
(1mol CO)
(1mol CO)
X = 12.2 L of CO

52. atomic
1 mole
mass
Mass
22.4 L
Representative
Particles
Volume of Gas
at STP
6.02x10 23
Mole

53. Practice
If you have a 35.67g piece of Chromium metal on your car, how many atoms of Chromium are in this piece of metal?
You are given mass and asked for number of particles
Set it up as a factor label problem

54. atomic
1 mole
mass
Mass
22.4 L
Representative
Particles
Volume of Gas
at STP
6.02x10 23
Mole

55. Representative Volume of Gas Particles at STP 23 6.02x10 22.4 L Mole
atomic
1 mole
mass
Mass
22.4 L
Representative
Particles
Volume of Gas
at STP
6.02x10 23
Mole
We are given a mass

56. Representative Volume of Gas Particles at STP 23 6.02x10 22.4 L Mole
atomic
1 mole
mass
Mass
22.4 L
Representative
Particles
Volume of Gas
at STP
6.02x10 23
Mole
We are asked for atoms
It’s going to take us 2 conversions, we just need to follow the arrows
We are given a mass

57. 1st we must convert our given mass of Chromium to moles of Chromium
So we need to use the periodic table to calculate the mass of 1 mole of Chromium
1 mole Cr
X mol Cr =
51.996g Cr
35.67g Cr
(51.996gCr)(X mol Cr) =
(35.67gCr)(1 mol Cr)
X = mol Cr

58. 2nd we must convert our newly found moles of Cr to atoms of Cr
So we need to remember that 1 mole of anything there are 6.02x1023 particles
6.02x1023 atoms Cr
X atoms Cr =
1 mole Cr
.6860 mol Cr
(1 mol Cr)(X atoms Cr) =
(6.02x1023 atoms Cr)
(.6860 mol Cr)
X = 4.13x1023 atoms Cr

59. Factor label one step method
Given: 35.67g Cr
Want: atoms Cr
Conversions:
1 mole Cr = 6.02*1023 atoms
1 mole Cr = g Cr
Set up
35.67g Cr | 1 mole Cr |6.02*1023 atoms | g Cr | 1 mole Cr
= *1* 6.02*1023 atoms= 2.148*1025 atoms = 4.131*10 23 atoms

60. Percent Composition
Atomic weights can be found by using the Periodic Table of the Elements.
Find the atomic weight of each element.
Multiply the number of atoms by the atomic weight, and add all of these weights per compound together.
Divide each atom (or ion's) individual atomic mass by the total chemical mass *100

61. Percent Composition FeO = Fe = 55.845g/mol O = 15.999 g/mol

62. Percent Composition FeO = 71.844 Fe = 55.845g/mol /71.844 = 0.777
= 0.223

63. Percent Composition
Fe = 0.777*100= %
0 = *100 = %

64. Percent Composition Li3P - Lithium Phosphide Li = P = 3* 6.941 =
20.823
/ =
0.402
30.97
1* =
/ =
0.598
51.793

65. Percent Composition C=5*12.01 = 60.050 C= 60.050 / 145.140 = 0.411
What is the percent carbon in C5H8NO4 (the glutamic acid used to make MSG monosodium glutamate), a compound used to flavor foods and tenderize meats?
a) %C
b) %C
c) %C
C=5*12.01 =
H= 8*1.01 = 8.080
N = 1*14.01 = 14.01
O= 4* =64.0
C5H8NO4 =
C= / = 0.411
0.411*100 = 41.1%

66. Chemical Formulas of Compounds
Formulas give the relative numbers of atoms or moles of each element in a formula unit - always a whole number ratio (the law of definite proportions).
NO atoms of O for every 1 atom of N
1 mole of NO2 : 2 moles of O atoms to every 1 mole of N atoms
If we know or can determine the relative number of moles of each element in a compound, we can determine a formula for the compound.

67. Types of Formulas Empirical Formula
The formula of a compound that expresses the smallest whole number ratio of the atoms present.
Ionic formula are always empirical formula
Molecular Formula
The formula that states the actual number of each kind of atom found in one molecule of the compound.

68. Simplest and molecular formulae
Consider NaCl (ionic) vs. H2O2 (covalent) Cl Na Cl H O Na H O H O Cl Cl Na Na
Chemical formulas are either “simplest” (a.k.a. “empirical”) or “molecular”. Ionic compounds are always expressed as simplest formulas.
Covalent compounds can either be molecular formulas (I.e. H2O2) or simplest (e.g. HO)

69. Simplest and molecular formulae
Consider NaCl (ionic) vs. H2O2 (covalent) Cl Na Cl H O Na H O H O Cl Cl Na Na
Q - Write simplest formulas for propene (C3H6),
C2H2, glucose (C6H12O6), octane (C8H14)
Q - Identify these as simplest formula, molecular
formula, or both H2O, C4H10, CH, NaCl

70. Answers Q - Write simplest formulas for propene (C3H6),
C2H2, glucose (C6H12O6), octane (C8H14)
Q - Identify these as simplest formula, molecular
formula, or both H2O, C4H10, CH, NaCl
A - CH2
A - H2O is both simplest and molecular
C4H10 is molecular (C2H5 would be simplest)
CH is simplest (not molecular since CH can’t form a molecule - recall Lewis diagrams)
NaCl is simplest (it’s ionic, thus it doesn’t form molecules; it has no molecular formula) CH
CH2O
C4H7

71. To obtain an Empirical Formula
1. Determine the mass in grams of each element present, if necessary.
2. Calculate the number of moles of each element.
3. Divide each by the smallest number of moles to obtain the simplest whole number ratio.
If whole numbers are not obtained* (in step 3), multiply through by the smallest number that will give all whole numbers
* Be careful! Do not round off numbers prematurely

72. A sample of a brown gas, a major air pollutant, is found to contain 2
A sample of a brown gas, a major air pollutant, is found to contain 2.34 g N and 5.34g O. Determine a formula for this substance.
require mole ratios so convert grams to moles
moles of N = 2.34g of N = moles of N
14.01 g/mole
moles of O = g = moles of O
16.00 g/mole
Formula:

73. A sample of a brown gas, a major air pollutant, is found to contain 2
A sample of a brown gas, a major air pollutant, is found to contain 2.34 g N and 5.34g O. Determine a formula for this substance.
moles of N
moles of O
/0.167 =1 =2
/0.167
NO2
N1O2

74. Empirical Formula from % Composition
A substance has the following composition by mass: % Na ; % B ; % H
What is the empirical formula of the substance?
2.645
/ 2.645
= 1
Na = g
/ =
B = g
/ =
2.645
/ 2.645
= 1
H = g
/ =
10.516
/ 2.645
= =4
NaBH4

75. Calculation of the Molecular Formula
Molecular formula is actual formula of the compound.
To determine the molecular formula from the empirical formula you must have the mass of the molecular formula. (Always Given to you)
Divide the mass of molecular formula by the mass of the empirical formula.
This will give you a whole number that you must multiply the subscripts in the empirical formula by.

76. Calculation of the Molecular Formula
CH2O is the empirical formula for a compound with a molecular mass of g.
CH2O = = 30.02g/mol
Divide by 30.03
180.18/30.03 = 6
6 (CH2O) = C6H12O6

77. Calculation of the Molecular Formula
A compound has an empirical formula of NO2. The colourless liquid, used in rocket engines has a molar mass of 92.0 g/mole. What is the molecular formula of this substance?
N2O4

78. CO2 H5N = Empirical CO2 H5N = 63.050g/mol 252.3 / 63.050 = 4
A compound contains 0.80 mol C, 4.0 moles H and 1.60 mol O, and 0.80 mol of N, and has a molecular mass of g/mol. What is the molecular formula for the compound?  
CO2 H5N = Empirical
CO2 H5N = g/mol
252.3 / = 4
4(CO2 H5N) = C4O8H20N4