1. Solids, Liquids and Gases
Solids (s): Molecules are held in place by intermolecular interactions.
Liquids (l): Molecules are held next to one another by intermolecular interactions, however, these interactions are not strong enough to prevent the molecules from flowing
past one another.
Gases (s): The intermolecular
interactions are too weak
to hold the molecules
next to one another,
so the molecules wander
off on their own.

2. Diffusion
Diffusion is the spontaneous spreading out of a substance, and is
due to the natural movement of its particles.
Cotton wool
soaked with hydrochloric acid
Ammonium chloride
Cotton wool
soaked with
ammonia solution
Diffusion of ammonia and hydrogen chloride gases
Ink being diffused in water

4. The mass of atoms are very small
Therefore scientists cannot count individual atoms or molecules directly
Instead a scientist count there particles by weighing a certain number of them
Same as a bank cashier, by knowing the weight of one coin they can calculate the number of coins by dividing the total mass by the mass of one coin.

5. A. What is the Mole? VERY A large amount!!!!
A counting number (like a dozen)
Avogadro’s number (NA)
1 mol = 6.02  1023 items
A large amount!!!!
VERY

6. The Mole
The mole is simply a chemist’s measurement for the amount of a substance.
It is a very large number (6 x 1023) as it ‘counts’ very small ‘objects’ such as atoms and molecules which have a very small masses.
This rather strange number was not selected deliberately – instead, it is the number of atoms in a sample of any element that has a mass in grams that is numerically equal to the elements atomic mass.
It allows chemists to work with masses in the lab which are feasible.

7. The Mole Definition
A mole of any substance is defined as the amount of the substance that contains as many particles as 12 g of carbon – 12.

8. B. Molar Mass Mass of 1 mole of an element or compound.
Atomic mass tells the...
That 6.01 *1023 atoms will weigh this amount
6.01 *1023 atoms of carbons will weigh 12 grams
grams per mole (g/mol)
1 mole of carbon is 12 grams per litre
Round to 2 decimal places

9. B. Molar Mass Examples carbon aluminum zinc 12.01 g/mol 26.98 g/mol

10. B. Molar Mass Examples water sodium chloride
2(1.01) = g/mol
H2O
NaCl
= g/mol

11. B. Molar Mass Examples sodium bicarbonate sucrose NaHCO3
(16.00) = g/mol
C12H22O11
12(12.01) + 22(1.01) + 11(16.00) = g/mol

12. If a magnesium atom is twice as heavy as an atom of carbon
Then 100 moles of magnesium
Are twice as heavy as 100 moles
Of carbon
Therefore if we have a piece of
Magnesium that is twice as heavy
As a piece of carbon they have an
Equal number of atoms

13. Atom of hydrogen weighs Mr=1
Atom of carbon weighs
Mr=12
Atom of hydrogen weighs
Mr=1
One gram of hydrogen atoms and twelve grams of
Carbon atoms both have the same number of atoms

14. 108grams Atom of silver weighs Ar=108 Atom of carbon weighs Ar=12
If we have 12g of carbon how do we calulate the mass of silver if it has the same number of atoms
108grams

15. All these masses contain the same number of atoms
12g
24g
27g
40g
56g
63.5
108g
carbon
magnesium
aluminium
calcium
iron
copper
silver
All these masses contain the same number of atoms
Avogadro`s number = 6 X Symbol L
It is the number of atoms of carbon in 12g of the carbon-12 isotope

16. The amount of a substance which contains 6 X 1023 particles is called a mole of that substance
Symbol mol

17. One mole of carbon weighs 12grams
48g of Mg (one mole weighs 24 g)
One mole of carbon contains 6 X 1023 atoms
Is 2 moles of Mg

18. Converting moles to grams
Mass of one mole of an element= relative atomic mass in grams
Element
Relative atomic mass H 1 C 12 N 14 Mg 24 Cu
63.5 Fe 56 I
127

19. The mass of one mole of an element
Relative
Atomic
Mass
in grams =

20. The relative molecular mass of a compound is the sum of the relative atomic masses of all the atoms in the compound
H2O = 2(1) + (16) = 18 (relative molecular mass of H2O)
CuSO4.H2O =(63.5) + (32) +4(16) +5 (18)=249.5

21. Converting Grams to moles
12g of carbon = 1 mole
1g of carbon = 1/12 moles
36g of carbon= 36/12=3 moles

22. How many moles of sulfuric acid (H2SO4) are there in 12
How many moles of sulfuric acid (H2SO4) are there in g of sulfuric acid
Relative mol mass =2(1) + (32) + 4(16) = 98
Total mass/mass of one mole
12.25/98=0.125 mole

23. Converting moles to number of atoms or molecules
How many atoms of sodium are present in 0.25 moles of a metal
1 mole of Na contained 6x1023 atoms
0.25 x 6x1023 atoms= 1.5 x1023 atoms

25. Molar Volume
One mole of any gas should occupy the same volume as one mole of any other gas at the same conditions of temperature and pressure.
Gases are often compared at STP, standard temperature and pressure.
Standard pressure = Pa
Standard temperature = 273 K
Molar Volume = 22.4 L = 22,400 cm3 = 2.24 x 10–2 m3

27. Relative Molecular Mass
The relative molecular mass is the average mass of a molecule relative to one twelfth the mass of the carbon 12 atom. The relative molecular mass has no units as it is the ratio of two masses.

29. Molar Mass
The molar mass of a substance is the mass in grams of one mole of the substance.
The molar mass has the same numerical value as its relative molecular mass, but its units are grams (g).
1 mol of carbon, 12 g

31. More Mole Calculations
Volume of X
(if a gas at STP) in litres at STP
Number of particles of X
÷ 22.4
÷ 6 x 1023
x 22.4
x 6 x 1023
Moles of X
Notice that there is no direct link from particles to grams, you must first convert to moles. Likewise going from volume to mass.
x molar mass
÷ molar mass
Mass of X in g

35. Chemical Formulas

36. Empirical Formula
The empirical formula of a compound is the formula that gives that simplest whole number ratio in which the atoms of the elements in the compound are present.
The empirical formula of glucose, C6H12O6 (ratio of atoms is 6:12:6) is CH2O, as this is the lowest ratio (1:2:1) of the atoms C, H and O.

40. Molecular Formula
The molecular formula of a compound is the formula that gives the actual number of atoms of each element present in a molecule of the compound.

42. Percentage Composition by Mass
If the empirical formula of a compound is known, the percentage by mass of each element present can be calculated. It can be useful to calculate the percentage of a certain element in a compound, for example the percentage of nitrogen in fertilisers.

44. Structural Formulas
The structural formula of a compound shows the arrangement of the atoms within a molecule of the compound.
Some Examples:
You will come across many more structural formulas in the organic section of the course.
Methane, CH4
Ethene, C2H4
Ethanol, C2H5OH

45. Chemical Equations

46. Chemical Equations
A chemical equation is a way of representing a chemical change.
It shows reactants and products.
To balance an equation means to change the numbers of each molecule involved, so that the same number of atoms of each element appear on the reactants side and on the products side.
Chemical equations balance on an atomic level, not molecular.
You cannot change the formula of a substance, i.e. if the equation has NH3 you cannot change this you can only put a number in front of it, 2NH3, increasing the number of N’s and the number of H’s.
Never change the subscripts (small numbers).
It is possible to write balanced equations for lots of reactions but that does not mean that the reaction actually takes place.

47. Formation of Water (Oxygen does not exist as single atoms)
(This reaction does not take place)

48. Balancing an Equation

49. Balance the following

50. Balancing Redox Equations
In working out what is oxidised and what is reduced in a reaction, it is important to remember that oxidation numbers are not a charge.
Write the oxidation numbers below each atom to which it applies, as shown in the examples.
Oxidation is an increase in oxidation number.
Reduction is a decrease in oxidation number.

51. Assign oxidation numbers to all
the atoms in the equation.
Write down half equations of what is oxidised
and reduced.
Attach subscripts to atoms oxidised and reduced
and balance the half equations.
Balance the half equations.
Attach species that were attached to the oxidised
and reduced atoms.
Include all the original species and complete the balancing by inspection.

52. Calculations using Chemical Equations
A balanced equation for a chemical reaction gives the relative amounts of each reactant and each product involved in the reaction.
If the amount of one substance is known, based on the molar ratios in the equation, the amounts (masses, particles or volume)of other substances can be calculated.
Make sure to always work in moles.