1. The Mole Part I: Mole Conversions
Part II: % Composition, Empirical Formulas, & Molecular Formulas

2. Part I:
Molar Conversions

3. Counting Particles
Different units are used to count different types of objects:
DOZEN EGGS PAIR OF SOCKS
REAM OF PAPER

4. Counting Particles
Each counting unit, dozen, pair, ream, etc., is appropriate for the kinds of objects they’re used to count……
…..it is easier to buy and sell paper by the ream – 500 sheets – than by the individual sheet

5. Counting Particles
Chemists also need a convenient method for accurately counting the number of atoms, molecules, or formula units in a sample of substance
But atoms are SO small, it is impossible to count them directly – THUS, chemists created a counting units called the “MOLE”

6. THE MOLE
Chemists use the MOLE to count atoms, molecules, ions and formula units
THE MOLE abbreviated “mol” is the SI base unit used to measure the amount of a substance.
The MOLE is defined as the number of carbon atoms in exactly g of pure carbon-12.

7. THE MOLE
Through experimentation, it has been established that a MOLE of ANY SUBSTANCE contains
x 1023 representative particles
A representative particles is any kind of particle, such as an atom, a molecule, a formula unit, an electron or an ion.

8. THE MOLE
The representative particles in a mole of water is a water molecule The representative particle in a mole of copper is the copper atom The representative particle in a mole of sodium chloride is the NaCl formula unit

9. THE MOLE
x 1023 This number is known as Avogardro’s number = 602,213,670,000,000,000,000,000 In this class, we abbreviate Avogardro’s number 6.02 x 1023

10. THE MOLE
Avogardro’s number would NOT be a convenient way to count something like marbles…. A mole of marbles would cover the surface of Earth to a depth of more than SIX KILOMETERS….. If you spent a billion dollars every second you couldn’t spend a mole of dollars in your lifetime!

11. CONVERTING BETWEEN MOLES AND PARTICLES
Dimensional Analysis/Factor Label To find the number of particles in 3.5 moles of glucose….(the representative particle here is the molecule) 1. Use the relationship: 1 mol = 6.02 x Develop conversion factors: 1 mol or 6.02 x x mol

12. What we call “particles”
If you have…
We call the particles…
Element
Atoms
Ionic Compound (metal+non-metal)
Formula units
Covalent Compound
(non-metal + non-metal)
molecules

13. CONVERTING BETWEEN MOLES AND # PARTICLES
QUESTION: 3.5 moles of glucose = ? # molecules 3. Set up dimensional analysis and choose the conversion factor that will solve for the unknown 1 mol or 6.02 x 1023 molecules 6.02 x 1023 molecules 1 mol 3.5 mol glucose 6.02 x 1023 molecules = 1 mol glucose x 1023 molecules = 2.1 x 1024 molecules

14. How many molecules are in 7.5 moles of carbon dioxide?

15. A sample of lead contains 1. 5 x 1024 atoms
A sample of lead contains 1.5 x atoms. How many moles of lead atoms are there?

16. A sample of zinc chloride contains 2. 150 x 1018 formula units
A sample of zinc chloride contains x 1018 formula units. How many moles of zinc chloride are there?

17. MASS AND THE MOLE
A mole of anything always contains the same number of particles, however, a mole of different substances have different masses. A dozen apples would not have the same mass as a dozen grapes….. If you put one mole of carbon atoms and one mole of copper atoms on separate scales, they, too, would have different masses

18. USING MOLAR MASS
Suppose you were asked to count the number of candies in a large bag of skittles…… You could take a FEW minutes and count them one by one…..

19. USING MOLAR MASS
OR you could “count” by massing the candy…. If one piece of candy has a mass of 1 g, and the whole bag of skittles less the mass of the bag is 691 g, how many skittles are in the bag?

20. USING MOLAR MASS
We can do the same thing with atoms. They are TOO small to count, but if we know the mass of one particle, and the mass of the whole sample, we can “count” the number of particles present in the sample. Instead of skittles, we have atoms, molecules ions, or formula units

21. MOLAR MASS
The mass of one mole of any pure substance is called its MOLAR MASS The molar mass of any element is numerically equal to its atomic mass and has the units g/mol (A mole is defined as the number of carbon-12 atoms in exactly 12 g of pure carbon-12, thus 12 g of carbon have 6.02 x 1023 atoms present)

22. MOLAR MASS
The samples below represent one mole of several elements: 1 mol Hg = g Molar Mass of Hg = g/mol Molar mass of S = g/mol Molar mass of Fe = g/mol Molar mass of C = g/mol

23. CONVERTING MOLES TO MASS
Recall from the previous page……
1 mol Hg = g Molar Mass of Hg = g/mol
Suppose you had 2.5 mol of Hg, what is its mass?
2.5 mol Hg g = 500 g
1 mol Hg
93.8 g Hg mol Hg = mol g

24. CONVERTING MASS TO MOLES
Recall from the previous page…… 1 mol Hg = g Molar Mass of Hg = g/mol Suppose you had 341 g of Hg, how many mols of Hg do you have? 341 g Hg 1 mol Hg = 341 mol = 1.70 mol g

25. THE MOLAR MASS OFCOMPOUNDS
Recall that a chemical formula indicates the numbers and types of atoms contained in one unit of the compound. Ex. CF4 1 carbon atom 4 fluorine atoms (12.01 amu) (19.00 amu) What mass would one molecule of CF4 have? 1 C (12.01 amu) + 4 F (19.00 amu x 4) = 1 molecule would have a mass of amu

26. THE MOLAR MASS OF COMPOUNDS
What mass would one molecule of CF4 have? 1 C (12.01 amu) + 4 F (19.00 amu x 4) = 1 molecule would have a mass of amu If one molecule of CF4 has a mass of amu, how many grams would one mole of CF4 have? 1 mole of CF4 has 1 mole of C atoms and 4 moles of F atoms: g g = g/mol

27. DETERMINING THE MOLAR MASS OF A COMPOUND
Start with the CORRECT formula of the compound
Look up the molar mass of each element present in the compound
Multiply each molar mass by the number of moles of that element in the chemical formula (same as the subscript)
Add up the mass of each element

28. CONVERTING THE MOLES OF A COMPOUND TO MASS
Dimensional analysis can be used to convert between the mols and mass of a compound EX. Convert 3.25 mol H2SO4 to grams 3.25 mol H2SO g = 319 g 1 mol H2SO4

29. CONVERTING THE MASS OF A COMPOUND TO MOLES
Dimensional analysis can be used to convert between mass and mols of a compound EX. Convert g H2SO4 to mols g H2SO4 1 mol H2SO4 = mol g

30. What is the mass of 0.25 moles of chromium?

31. What is the molar mass of lithium nitrate, LiNO3?

32. What is the molar mass of calcium acetate, Ca(C2H3O2)2?

33. How many moles are in 423.1 g of calcium acetate, Ca(C2H3O2)2?

34. How many grams are in 12.8 moles of calcium acetate, Ca(C2H3O2)2?

36. MOLE RELATIONSHIPS FROM A CHEMICAL FORMULA
Al2O3: If you had 1.50 mol of Al2O3, how many mols of Al3+ ions are present? 1.50 mol Al2O3 2 mol Al3+ ion s = 3.0 mol Al3+ 1 mol Al2O3

37. If you had 4. 80 mol of Na2SO4, how many moles of Na1+ are present
If you had 4.80 mol of Na2SO4, how many moles of Na1+ are present? (SO4)2-?

38. CONVERTING THE MASS OF A COMPOUND TO NUMBER OF PARTICLES
# particles  mols  grams To convert from grams to number of particles, one has to go through “Mole-ville” Ex g Mg = ? # atoms of Mg 2.35 g Mg 1 mol Mg 6.02 x 1023 atoms Mg = g 1 mol 5.82 x 1022 atoms Mg

39. CONVERTING THE NUMBER OF PARTICLES OF A COMPOUND TO MASS
# particles  mols  grams
To convert from number of particles to grams, one has to go through “Mole-ville”
Ex X 1023 molecules of water = ? grams
8.31 x 1023 molecules 1 mol H2O g =
6.02 x mol H2O
24.9 g H2O

40. How many molecules of ammonia, NH3 are in a 68-g sample?

41. What is the mass of the oxygen atoms in 4 moles of sodium sulfate?

42. Mole Conversion Chart

43. Molar Conversions molar mass 6.02  1023 MOLES (g/mol) (particles/mol)IN
GRAMS
MOLES
NUMBER OF
PARTICLES
(g/mol)
(particles/mol)

44. % Composition, Empirical Formulas, & Molecular Formulas
Part II
% Composition, Empirical Formulas, & Molecular Formulas

45. Percentage Composition
the percentage by mass of each element in a compound
Be sure to include the total mass of the element (so for like CuCl2, include 2 Cl masses)

46. Percentage Composition
Find the % composition of Cu2S.
g Cu
g Cu2S
%Cu =
 100 =
79.852% Cu
32.07 g S
g Cu2S
%S =
 100 =
20.15% S

47. Percentage Composition
Find the percentage composition of a sample that is 28 g Fe and 8.0 g O.
28 g
36 g
%Fe =
 100 =
78% Fe
8.0 g
36 g
%O =
 100 =
22% O

48. Percentage Composition
How many grams of copper are in a 38.0-gram sample of Cu2S?
Cu2S is % Cu
(38.0 g Cu2S)( ) = 30.3 g Cu

49. Percentage Composition
Find the mass percentage of water in calcium chloride dihydrate, CaCl2•2H2O?
36.04 g g
%H2O =
 100 =
24.51%
H2O

50. C2H6 CH3 Smallest whole number ratio of atoms in a compound
Empirical Formula
Smallest whole number ratio of atoms in a compound
C2H6
reduce subscripts
CH3

51. 1. Find mass (or %) of each element.
Empirical Formula
1. Find mass (or %) of each element.
2. Find moles of each element by using the molar mass
3. Divide moles by the smallest # of moles calculated in step 2 to find subscripts.
4. When necessary, multiply subscripts by 2, 3, or 4 to get whole #’s.

52. Find the empirical formula for a sample of 25.9% N and 74.1% O.
25.9 g
1 mol
14.01 g
= 1.85 mol N
= 1 N
1.85 mol
74.1 g
1 mol
16.00 g
= 4.63 mol O
= 2.5 O

53. N2O5 N1O2.5 Need to make the subscripts whole numbers  multiply by 2
Empirical Formula
N1O2.5
Need to make the subscripts whole numbers
 multiply by 2
N2O5

54. CH3 C2H6 “True Formula” - the actual number of atoms in a compound
Molecular Formula
“True Formula” - the actual number of atoms in a compound
CH3
empirical
formula ?
C2H6
molecular
formula

55. 1. Find the empirical formula. 2. Find the empirical formula mass.
Molecular Formula
1. Find the empirical formula.
2. Find the empirical formula mass.
3. Divide the molecular mass by the empirical mass.
4. Multiply each subscript by the answer from step 3.

56. Molecular Formula
The empirical formula for ethylene is CH2. Find the molecular formula if the molecular mass is g/mol?
empirical mass = g/mol
28.1 g/mol
14.03 g/mol
= 2.00
(CH2)2  C2H4