1. UNIT 2 – QUANTITATIVE CHEMISTRY
(I.E. IN COMES THE MATHS!! Bring a calculator!!)

2. Scientific Notation… A (very) brief introduction

3. Stoichiometry
The study of the relationships between the quantities of reactants and products involved in chemical reactions.
Stoikheion = element
metron = measure

4. Proportions in compounds
The composition of a compound depends on two things
1. The elements they are composed of
2. The quantities of each element.

5. Law of definite proportions
A specific compound always contains the same elements in definite proportions by mass.
This is always constant even if produced other ways.

6. Relative atomic mass and isotopic abundance (p80)
Relative atomic mass = the mass of an element that would react with a fixed mass of a standard element, currently carbon-12.
One atomic mass unit (u) is described as 1/12 the mass of a carbon-12 atom.

7. Atomic mass unit Now we know that it weighs
1 u = x kg.
This is the standard used world wide

8. Isotopic abundance
Some elements have isotopes. In calculating the relative atomic mass of an element with isotopes, the relative mass and proportion of each is taken into account. For example, naturally occurring carbon consists of atoms of relative isotopic masses C-12 (98.89%) and-C 13 (1.11%). Its relative atomic mass is u.

9. Isotopic abundance mc = (98.89/100 x 12) + (1.11/100 x 13) =
mc = = u
No atom will contain u. it is an average.

10. Example
Cl-35 = 75.53%
Cl-37 = 24.47%
(35 x .7553) + (37 x .2447) = u
Therefore the average atomic mass is 35.49u

11. Page 81 # 1

12. Atomic Mass and Molecular Mass
The mass of one atom of an element expressed in atomic units, u.
Molecular mass
The mass of one molecule, expressed in atomic mass units, u.

13. Example MNH3 = 1(mN) + 3(mH) = 1(14.01 u) + 3(1.01 u) = 17.04 u .
Therefore, one molecule of ammonia has a mass of units.

14. Formula unit of mass
The mass of one formula unit of an ionic compound, expressed in atomic mass units, u

15. Example NaCl = 1(mNa) + 1(mCl) = 1 (22.99 u) + 1(35.45 u) = 58.44 u
Therefore, one formula unit of sodium chloride has a mass of units.

16. Page 81 # 1-3

17. The Periodic Table and the Mass of Elements
The atomic number of a substance is written near the top of the symbol. This number tells the number of protons in the element.
The Atomic mass is usually written near the bottom. This number is given as a numerical number of the elements mass. The mass of one mole of the element measured in grams.

18. The Mole

19. It's not a spy, a machine for digging tunnels, a burrowing animal, or a spot of skin pigmentation - it's
- a unit of measurement containing about
6.02 x 1023
particles.
Chemists created the mole as their own counting unit.
The Chemical Mole

20. (This is the number of Carbon atoms there are in 12.01 g or Carbon)
1 mole of atoms is
602, 214,199, 000, 000, 000, 000, 000
atoms!
(This is the number of Carbon atoms there are in g or Carbon)

21. Avogadro’s Number 6.02 x 1023 particles / mole
When you say you have a mole, it’s kind of like saying you have ‘a dozen’ of something…
A dozen donuts = 12 donuts
A dozen apples = 12 apples
A mole of CO2 molecules = 6.02 x 1023 molecules
A mole of Na atoms = 6.02 x 1023 atoms

22. Just How Big is a Mole?
Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles.
If you had Avogadro's number of unpopped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles.
If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.

23. Molar mass (g/mol) = M
the mass, in grams per mole, of one mole of a substance. The SI unit for a mol is g/mol
Molar mass is conventionally abbreviated with a capital M. For elements it is numerically equal to the atomic mass and is often called atomic mass units.

24. Steps to calculating molar mass
Write the chemical formula of the substance
Determine the number of atoms or ions of each element in one formula unit of the substance.
Us the atomic molar masses from the periodic table and the amounts in moles to determine the molar mass of the chemical
Write the molar mass in units of grams per mol (g/mol).

25. Example Molar mass of copper(Cu) = mass of 1 mol of Cu atoms
M = g/mol Cu
M = mass of x atoms

26. Example 2
Molar mass of glucose (C6H12O6) = mass of exactly 1 mol of glucose molecules
M = (6 x C) + (12 x H) + ( 6 x O)
M = ((6 x (C)) + (12 x 1.01(H)) + (6x (O)) ) / mol
M = ( (72.06 g) + (12.12 g) + (96.00g) ) / mol
M = ( g) / mol
M = g/mol
Therefore, 1 mol glucose has a mass of g and contains x 1023 glucose molecules(C6H12O6)

27. Important things to note
Specify what is being measured
Ex. molar mass of oxygen atom = g/mol
Molar mass of oxygen molecules = g/mol
When talking about oxygen we are normally referring to the diatomic molecule
Since electrons have negligible mass compared to protons and neutrons, ions will essentially have the same mass as the atoms.

28. Moles of Elements Why that number?
1 mole element = (atomic mass) in grams of that element
1 mole C = g C
1 mole B = g B
1 mole Cu = g Cu
i.e. atoms are teeny, teeny, tiny, teeny, itty, bitty things and it is not reasonable to talk about them in their atomic weights, but if scale them up to moles, then working with balanced chemical equations etc. becomes easier.

29. Determine the mass one mole of the following in grams:
CO2
NaCl
K2CrO4
Fe(NO3)3

30. PAGE 85 # 4

31. The Mole and Molar Mass The Mole = n
Avogadro’s constant (NA) = the number of entities in one mole = 6.02 x1023
Eg. 1 dozen oxygen atoms = 12 oxygen atoms
1 mole oxygen atoms = x 1023 oxygen atoms

32. Calculations involving Molar mass
Amount in moles (n) = mass(g)
Molar mass(g/mol)
The Magic Triangle
n = m / M
m = n x M
M = m / n

33. GRASP method 1 Given = What do you know?
2 Required = What do you want to know? 
3 Analysis = What formula can you use? 
4 Solution = Show your work 
5 Paraphrase = State the answer in a sentence.

34. Example
Lets say that we are given 22 grams of graphite. How many moles do we have?

35. Lets say we want to know how much CaCl2 we need for a reaction that asks for 13.52 moles

36. You are given an unknown sample that has a mass of 32. 04 grams, 2
You are given an unknown sample that has a mass of grams, moles. Find M.

37. Practice: Determine the number of mol in 13.7 g of CO
What is the mass of 2.75 mol of C3H8

38. Homework
Page 86 # 5-7
Isotopic Abundance and Mole calculations Worksheet

39. Calculating Number of Entities

40. Calculating Number of Entities (N) from Mass
We are not able to count the number of atoms in reactions; we must count them indirectly by measuring out a certain mass of a substance and calculating the number of moles.
In order to estimate the number of entities in a large sample we need to take a small sample and find out the mass of a small number of entities.

41. Calculating Number of Entities and Avogadro’s number
There are times when need to know exactly how many atoms are in something. This is especially true when dealing with gases as they change in pressure with temperature.
To find the number of entities we use the formula

42. N = nNA
Number of entities = (# of mols) x(Avogadro's #)

43. Example Find the number of atoms in 0.004 moles of water.
Find the number of molecules in 30.0 g of Cl2 (g)

44. Determining the number of particles in a mole

45. 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!

46. Particles
Molecules
Compound
Formula unit
Atom
Element

47. Equation to find number of particles!
number of moles x = number of particles
6.02 x 1023 particles
1 mole
Example: How many particles of sucrose are in 3.50 moles?

48. Equation to find number of moles!
number of particles x = number of moles
1 mole
6.02 x 1023 particles
Example: How many moles are contained in 4.50 x 1024 atoms of zinc?

49. Learning Check 1. Number of atoms in 0.500 mole of Al
2.Number of moles of S in 1.8 x 1024 S atoms

50. Molar Mass

51. Molar Mass The Mass of 1 mole (in grams)
Molar mass is equal to the atomic mass (look on the periodic table)
1 mole of C atoms = g
1 mole of Au atoms = g
1 mole of Cu atoms = g
Always round to the hundredths place.

52. Mole to Mass Conversion
Where will you find the number of grams per mole?
Mole to Mass Conversion
Number of moles x number of grams = mass
1 mole
Example: Calculate the mass in grams of moles of chromium.
moles Cr x g Cr = g Cr
1 mol Cr
The Periodic Table
Look at your Periodic Table

53. Mass to Mole Conversion
Where will you find the number of grams per mole?
Mass x mole = number of moles
number of grams
Example: How many moles of calcium are there in 525 grams of calcium?
525 g Ca x mol Ca = mol Ca
40.1 g Ca
The Periodic Table
Look at your Periodic Table

54. Mass to Atoms Conversion
Two steps:
Use the mass to moles conversion
Now calculate the number of atoms using the moles to particles equation.

55. Example:
How many atoms of gold are in a pure gold nugget having a mass of 25.0 g?

56. 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

57. Learning Check! = 1.20 X 1024 atoms K
How many atoms of K are present in 78.4 g of K?
78.4 g K mol K X 1023 atoms K g K mol K
= 1.20 X 1024 atoms K

58. Empirical and Molecular Formulas
11.4

59. Percent Composition Mass of element x 100 = percent by mass
Mass of compound
Example: Determine the percent composition of NaHCO3.
Determine the mass of each element first!
%mass element = mass of element in 1 mol compound x 100
molar mass of compound
Percent Na = g Na x 100 = 27.37% Na
84.01 g NaHCO3
You have to complete this step for each element in the compound!!!!

60. Determine the Percent Composition
from the Formula C2H5OH
Divide the mass of each element by the molar mass of the compound and multiply by 100%

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

62. Determine the Empirical Formula of
Acetic Anhydride if its Percent Composition is
47% Carbon, 47% Oxygen and 6.0% Hydrogen
Convert the percentages to grams by assuming you have 100 g of the compound
Step can be skipped if given masses

63. Determine the Empirical Formula of
Acetic Anhydride if its Percent Composition is
47% Carbon, 47% Oxygen and 6.0% Hydrogen
Convert the grams to moles

64. Determine the Empirical Formula of
Acetic Anhydride if its Percent Composition is
47% Carbon, 47% Oxygen and 6.0% Hydrogen
Divide each by the smallest number of moles

65. Determine the Empirical Formula of
Acetic Anhydride if its Percent Composition is
47% Carbon, 47% Oxygen and 6.0% Hydrogen
If any of the ratios is not a whole number, multiply all the ratios by a factor to make it a whole number
If ratio is ?.5 then multiply by 2; if ?.33 or ?.67 then multiply by 3; if ?.25 or ?.75 then multiply by 4
Multiply all the
Ratios by 3
Because C is 1.3

66. Determine the Empirical Formula of
Acetic Anhydride if its Percent Composition is
47% Carbon, 47% Oxygen and 6.0% Hydrogen
Use the ratios as the subscripts in the empirical formula
C4H6O3

67. Molecular Formulas
The molecular formula is a multiple of the empirical formula
To determine the molecular formula you need to know the empirical formula and the molar mass of the compound

68. Determine the Molecular Formula of Benzopyrene
if it has a molar mass of 252 g and an
empirical formula of C5H3
Divide the given molar mass of the compound by the molar mass of the empirical formula
Round to the nearest whole number

69. Determine the Molecular Formula of Benzopyrene
if it has a molar mass of 252 g and an
empirical formula of C5H3
Multiply the empirical formula by the calculated factor to give the molecular formula
(C5H3)4 = C20H12