1. UNIT 1: Particles AICE CHEMISTRY
Relative isotopic mass, Relative atomic mass, Relative molecular mass, The mole

2. Relative isotopic mass
The relative isotopic mass (Ir) of an isotope is the mass of an atom of that isotope relative to the mass of an atom of 12C taken as 12 units exactly

3. Relative atomic mass
An element can have several naturally occurring isotopes.
These isotopes of a element behave in the same way.
In calculating the relative atomic mass of an element with isotopes, the relative mass and proportion or percentage of each is taken into account.

4. Calculating relative atomic mass
Isotope
Relative isotopic mass
Relative abundance (%) Cl
34.969
75.80
36.966
24.20
Ar = (relative isotopic mass X1 % abundance) + relative isotopic mass X2 % abundance)
100
Ar (Cl) = ( X 75.8) + ( X 24.2)
Ar (Cl) =
Ar (Cl) = amu (atomic mass unit)

5. Your turn simulation for isotopes-and-atomic-mass
Calculate the relative atomic mass of Boron to two decimal place
Isotope
Relative abundance (%)
Relative isotopic mass
10B
19.91
11B
80.09
11.009

6. Significant figures are important because they tell us how good the data we are using is.   
100 grams
This number has only one significant figure   Because this digit is in the “hundreds” place, this measurement is only accurate to the nearest 100 grams
100. grams
This number has three significant figures (the decimal makes all three digits significant.  Because the last significant figure is in the “ones” place, the measurement is accurate to the nearest gram.

7. grams
The number has five significant figures.  Because the last significant figure is in the “hundredths” place, the measurement can be considered to be accurate to the nearest grams

8. Significant Figures
The rules for identifying significant digits when writing or interpreting numbers are as follows:
1. All non-zero digits are considered significant. Example: 1, 20, and 300 all have one significant figure. Their significant figures are 1, 2, and 3 respectively.
has five significant figures: 1, 2, 3, 4 and 5.
2. Zeros appearing anywhere between two non- zero digits are significant.
Example: has five significant figures: 1, 0, 1, 1 and 2.
3. Leading zeros are not significant. For example, has two significant figures: 1 and 2.

9. 4. Trailing zeros in a number containing a decimal point are significant. For example, has six significant figures: 1, 2, 2, 3, 0 and 0.
The number still has only six significant figures. In addition, has five significant figures.
5. Any zero that’s after all of the nonzero digits is significant only if you see a decimal point
has 1 significant figure has 6 significant figures.
precision and accuracy on a balance

10. Scientific Notation
Scientific Notation is a way of writing numbers that accommodates values too large or small to be conveniently written in standard decimal notation.
Ordinary decimal notation
Scientific notation
,570 5,720,000,000 −
= 3×102 = 4×10-5 = 4.57×103 = 5.72×109 = −6.1×10−9

11. Relative molecular mass (Mr )
The relative molecular mass of a compound is the mass of one molecule of that substance relative to the mass of a 12C
This is calculated by taking the sum of the relative atomic masses of the elements in the molecular formula (i.e. covalent compounds)
It is called relative formula mass for ionic compounds

12. Calculating relative molecular mass
Normally to find the relative atomic mass you just look in the periodic table!!
Oxygen (O2)Mr = 2 X Ar(0)
= 2 X 16.0
= 32.0 amu
Carbon Dioxide (CO2)Mr = Ar(C) X Ar(0)
= (2 X 16.0)
= 44.0 amu

13. Mass Spectrometer
is an analytical technique that can be used to determine the chemical composition of a sample
It Is a machine that separates the individual isotopes in a sample
From this you can
calculate the relative
atomic mass.

14. Mass spectrum – graph of a mass spectrometer.
Number of peaks = number of isotopes
Position of peaks = relative isotopic mass
Height of peaks = relative abundance of
the isotope,
in comparison to the
other isotopes of that
element.

15. Questions from the book
Pg 5 Question 2 part a and b

16. x

17. The Mole
Review the mole unit!

18. What is a mole? What is a dozen? 12 of something
The mole is like a dozen, it is a fixed number of something
The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 12 g (0.012 kilograms) of carbon 12.
That number is 6.02 X1023 or
a mole is a fixed number of things just like a dozen.

19. Why have the mole?
It can become very cumbersome to deal with such large numbers, especially during calculations.
To overcome this impracticality, chemists deal with multiples of particles, instead of individually. This is why we have a mole!

20. Where did the mole come from?
The mole is a unit of measurement based on work done almost 200 years ago by Amadeo Avogadro as he studied gas behavior.
His work led to the association of a number, x 1023, with the mole.
6.02 x 1023, is also called Avogadro’s constant or number
It allows particles to be "counted.“
The word "mole" is derived
from "gram molecular weight"

21. How big is the mole
The mole can be applied to anything: number of trucks, number of balloons, number of Twilight fans.
In a mole we know there are 6.02 X1023 of stuff
it can be anything trucks, twilight fans etc
To have one mole of peas, they would have to be a metre deep around earth and you would need 250 earths to fit them all on

22. How big is the mole
If you believe the universe is 15 billion years old we have not yet had a mole of seconds yet!!!
So the mole is HUGE!
But you can hold a mole of NaCl in your
hand because molecules and atoms are
so small

23. Examples of moles
1 mole of iron contains the same number of atoms as 1 mole of gold
1 mole of sodium chloride contains the same number of molecules as 1 mole of water
the number of atoms in 1 mole of iron is equal to the number of molecules in 1 mole of water.
1 mole of water (H2O) has 1 mole of oxygen and 2 o hydrogen.

24. Lets deal with that big numbers first
If I have 3 mol of trucks. How many trucks do I have?
Its all about setting up your units correctly
If you look at the units they cancel out
3 mol of trucks x 6.02x1023 trucks
1 mol of trucks
= x 1024 trucks

25. Moles of atoms in molecules
How many moles of oxygen atoms are in 5 mol of O2 ?
1 mol of O2 contains 2 mol of O atoms
5 mol of O2 contains 10 mol of O atoms
How many moles of oxygen atoms are in 5 mol of H2SO4?
1 mol of H2SO4 contains 4 mol of O atoms
5 mol of H2SO4 contains 20 mol of O atoms

26. Atoms to Moles
I have x 1011 atoms of Zinc. How many mole do I have
Is it a mole?
4.673 x 1011 atoms of Zinc x 1 mol of Zinc
6.02 x 1023 atoms of Zinc
= 7.7 x10-11 mols of zinc
Always check the units have cancelled out
This number is less than one therefore it is less than one mol as it should be!
I have x 1011 atoms of Zinc. How many mole do I have
Is it a mole?
4.673 x 1011 atoms of Zinc x 1 mol of Zinc
6.02 x 1023 atoms of Zinc
= 7.7 x10-11 mols of zinc
Always check the units have cancelled out
This number is less than one therefore it is less than one mol as it should be!

27. Moles to atoms 1 mol = 6.02 x 1023 atoms of copper
I have 3.01 x 1023 copper atoms. How many mole of copper atoms do I have?
1 mol = 6.02 x 1023 atoms of copper
Mol of copper = 3.01 x 1023copper atoms
6.02 x 1023 atoms of copper
mol
Mol of copper = 0.5 mol
Check calculation!

28. Your turn
CHALLENGE
How many moles of seconds have we experienced (if the world is 15 billion years old)
Show units and unit cancellation

29. Molar Mass

30. Molar Mass
Molar mass (M) is the mass of one mole of a substance (chemical element or chemical compound).
The unit of molar masses is:
grams per mole ( g or g/mol or g mol–1)
mol

31. How do you find the molar mass of an element?
In general the molar mass of an element is the relative atomic mass of the element expressed in grams.
The molar mass of a compound is the relative formula mass of the compound expressed in grams.

32. What is the molar mass of oxygen? What is the molar mass of bromine?

33. Example Calculate the molar mass of table sugar, sucrose (C12H22O11)
Ar Carbon = 12.01
Ar Hydrogen = 1.00
Ar Oxygen = 16.00
M = (12 x 12.01)+ (22 x 1.00)+ (11 x 16.00)
M =
M = g/mol

34. Grams to moles Use the molar mass it has grams and moles
How many moles of Zinc chloride do I have in 2.6g ZnCl2?
Molar mass of ZnCl2 = (2 x 35.45)
= g/mol
2.6 g of ZnCl2 x 1mol of ZnCl2
g ZnCl2
= mol of ZnCl2

35. Grams to moles I have 30.0g of NH3. How many moles of NH3 do I have?
Ar (NH3)= (1.00 x 3)
= 17g/mol
= 30.0g x 1 mol
17g
= 1.76 mol of NH3

36. Moles to grams I have 1.973 mol of I2
How many grams of Iodine do I have?
Molar mass of I2 = x 2
= g/mol
1.973 mol of I2 x 253.8g
1 mol of I2
= g

37. Moles to Grams
I have mol of water how many grams of water do I have?
Ar (H2O) = (1.00 x 2) =
= g/mol
= mol H2O x 18.00g of H2O
1 mol of H2O
= 5.4 g of H2O

38. Moles to Grams
Find how many moles are given in the problem.
Calculate the molar mass of the substance
Multiply step one by step two.
Grams to Moles
Find the number of grams given in the problem.
Calculate the molar mass of the substance.
Divide step one by step two.

39. How many atoms of Iron? I have 5.5 g of Iron. How many atoms do I have
Ar = 55.45g/mol
= 5.5g x 1 mol
55.45g
= mol
= mol x 6.02 x 1023 atoms of Iron
1 mol of Iron
= 5.96 x 1022 atoms of Iron

40. Your turn Pg 6 Questions 3 and 4 Worksheet given by your teacher!
CHALLENGE
Using your knowledge of mole calculations and unit conversions, determine how many atoms there are in 1 L of petrol. Assume that the molecular formula for petrol is C6H14 and that the density of petrol is approximately grams/mL.

41. STOICHIOMETRY

42. What is stoichiometry?
Stoichiometry is the quantitative study of reactants and products in a chemical reaction.

43. 2 A + 2B 3C What You Should Expect Given : Amount of reactants
Question: how much of products can be formed.
Example
2 A + 2B C
Given 20.0 grams of A and sufficient B, how many grams of C can be produced?

44. What do you need?
You will need to use
molar ratios,
molar masses,
balancing and interpreting equations, and
conversions between grams and moles.
Note: This type of problem is often called "mass-mass."

45. Steps Involved in Solving Mass-Mass Stoichiometry Problems
Balance the chemical equation correctly
Using the molar mass of the given substance, convert the mass given to moles.
Construct a molar proportion (two molar ratios set equal to each other)
Using the molar mass of the unknown substance, convert the moles just calculated to mass.

46. Mole Ratios
A mole ratio converts moles of one compound in a balanced chemical equation into moles of another compound.

47. Example Mole Ratios: 2 : 1 : 2
Reaction between magnesium and oxygen to form magnesium oxide. ( fireworks)
2 Mg(s) + O2(g) MgO(s)
Mole Ratios:
2 : 1 : 2

48. Practice Problems Write the mole ratios for N2 to H2 and NH3 to H2.
1) N2 + 3 H2 ---> 2 NH3
Write the mole ratios for N2 to H2 and NH3 to H2.
2) A can of butane lighter fluid contains 1.20 moles of butane (C4H10). Calculate the number of moles of carbon dioxide given off when this butane is burned.

49. Mole-Mole Problems Equation of reaction 2C4H10 + 13O2 8CO2 + 10H2O
Using the practice question 2) above:
Equation of reaction
2C4H O CO H2O
Mole ratio
C4H CO2
1 : [ bases]
: X [ problem]
By cross-multiplication, X = 4.8 mol of CO2 given off

50. Mole-Mass Problems
Problem 1: 1.50 mol of KClO3 decomposes. How many grams of O2 will be produced? [K= 39, Cl = 35.5, O = 16]
2 KClO KCl + 3 O2

51. Three steps…Get Your Correct Answer
Use mole ratio
Get the answer in moles and then
Convert to Mass. [Simple Arithmetic]
Hello!
If you are given a mass in the problem, you will need to convert this to moles first. Ok?

52. Let’s go!
2 KClO KCl + 3 O2
2 :
: X
X = 2.25mol
Convert to mass
2.25 mol x 32.0 g/mol = 72.0 grams
Cool!

53. Try This:
We want to produce 2.75 mol of KCl. How many grams of KClO3 would be required?
Soln
KClO3 : KCl
2 :
X : 2.75
X = 2.75mol
In mass: 2.75mol X g/mol
= 337 grams zooo zimple!

54. Mass-Mass Problems
There are four steps involved in solving these problems:
Make sure you are working with a properly balanced equation.
Convert grams of the substance given in the problem to moles.
Construct two ratios - one from the problem and one from the equation and set them equal. Solve for "x," which is usually found in the ratio from the problem.
Convert moles of the substance just solved for into grams.

55. Just follow mass- mass problem to the penultimate level
Mass-Volume Problems
Just follow mass- mass problem to the penultimate level

56. Like this: There are four steps involved in solving these problems:
Make sure you are working with a properly balanced equation.
Convert grams of the substance given in the problem to moles.
Construct two ratios - one from the problem and one from the equation and set them equal. Solve for "x," which is usually found in the ratio from the problem.
Convert moles of the substance just solved for into Volume.

57. Conversion of mole to volume
No of moles = Volume
Molar volume
Can you remember a similar equation?

58. Molar volume
The molar volume is the volume occupied by one mole of ideal gas at STP. Its value is: 22.4dm3

59. Practice Problems
Calculate the volume of carbon dioxide formed at STP in ‘dm3' by the complete thermal decomposition of g of pure calcium carbonate (Relative atomic mass of Ca=40, C=12, O=16)
Solution:
Convert the mass to mole:
Molar mass of CaCO3 = (16 x 3) = 100gmol-1
Mole = mass/molar mass
3.125/100 = mol

60. Practice Problems As per the equation, Mole ratio 1 : 1
problem mol : X
X = mol of CO2
Convert mole to volume
Volume = ( mol x 22.4)dm3/mol
= 0.7dm3

61. Percentage Composition

62. Percentage composition
Percentage composition of a compound is a relative measure of the mass of each different element present in the compound.
E.g. The percentage of Zn in ZnCl2
or The percentage of Al in Al2O3

63. Formula and Example Example: Calculate the percentage of Al in Al2O3
% by mass of = mass of element in 1 mole of compound x 100
the element mass of 1 mole of the compound
Example: Calculate the percentage of Al in Al2O3
Step 1: Find the molar mass of the compound
Mr (Al2O3) = (2 x 27) + (3 x16)
= 102 g/mol
Step 2: Find the molar mass of the element in the compound
mass of Al in Al2O3 = (2 x 27)
= 54g/mol
Step 3 : Determine the % of the element in the compound
% Al in Al2O3 = 54g/mol x 100
102g/mol
= 52.9%

64. Your turn Calculate the percentage composition of Zinc (Zn) in ZnCl2
Calculate the percentage composition of oxygen (O) in Fe(NO3) 3
Calculate the percentage composition of Carbon in C7H14O2

65. Empirical Formulas

66. Empirical formula
The empirical formula is a simple expression of the relative numbers of each type of atom in it.
A person has two hands and ten fingers, or H2F10. The empirical formula for that would be HF5
Benzene, C6H6. The empirical formula is CH.
This empirical formula tells us that the ratio of C to H is 1 to 1; there is one H atom for every C atom.

67. Empirical formula Write the empirical formula of the following. C8H16

68. How to calculate empirical formulas

69. A compound of carbon and oxygen is found to contain 27
A compound of carbon and oxygen is found to contain 27.3% carbon and 72.7% oxygen by mass.
Carbon
Oxygen
Calculate the amount in moles
27.3 = 2.27
12.0
72.7 = 4.54
16.0
Divide the number of moles by smallest
2.27 = 1
2.27
4.54 = 2
Obtain the simplest whole number mole ratio 1 2

70. 9.0g of a compound of only C,H,O contains 4.8 g of O and 3.6g of C
m(g)
3.6
9 – ( ) = 0.6
4.8
n(mol)
3.6 =0.3
12.0
0.6 = 0.6
1.0
4.8 = 0.3
16.0
Divide by smallest mol
0.3 = 1
0.3
0.6 = 2
Ratio 1 2

71. Your turn
Pg 71 Q 40, 41 and 42

72. Molecular Formulas
A molecular formula gives the actual number of atoms in one molecule of the compound
The empirical and molecular formula can be the same
The molecular formula is always a whole number multiple o the empirical formula

73. A compound has the empirical formula of CH
A compound has the empirical formula of CH . The molar mass of the compound is 78 g/mol. What is the molecular formula of the compound?
The molar mass of CH is 13g/mol
Compound molar mass is 78g/mol
No. of CH unit in a molecule = 78g/mol
13g/mol
= 6
CH x 6 = C6H6 = molecular formula

74. A sample of hydrocarbon was found to contain 7. 2g of C and 1
A sample of hydrocarbon was found to contain 7.2g of C and 1.5g of hydrogen. The molar mass of the compound is 58g/mol. What is the molecular formula?
1. Determine empirical formula C H
7.2
1.5
7.2 = 0.6
12.0
1.5 = 1.5 1
0.6 = 1
0.6
1.5 = 2.5 2 5

75. 2. Determine the molecular formula
Empirical formula is C2H5
The molar mass of C2H5 is 29g/mol
Compound molar mass is 58g/mol
No. of C2H5 unit in a molecule = 58g/mol
29g/mol
= 2
C2H5 x 2 = C4H10 = molecular formula

76. Experiment: Decomposition of Baking Soda