as level chemistry

1. 4 States of matter
Physical Chemistry – CIE AS Chemistry

2. New syllabus for CIE
 This content is for exams in 2022.
 AS
 If you’re starting in September 2020 with Year 12, then you’ll need to use
the old syllabus for the AS in May/June 2021.
 Only start this syllabus with AS in September 2020 if your students will take
ALL the exams (AS & A2) at the end of two years (May/June 2022).
 A2
 Year 13 September 2020 old syllabus examined in 2021
 Year 13 September 2021 new (this) syllabus examined in 2022

3. Information for the teacher:
 Mastery learning – builds on the concept of developing automatic fluency in your learners. We drive cars on auto-
pilot so let’s aim to for our students to work in chemistry with the same level of autonomy. New material is present
in small with frequent practice and mastery of small sections before we move on all the definitions are centred on
the idea that students must achieve proficiency in each unit of work before proceeding.
 Repetition and practice are critical to mastery of different concepts as students proceed through the course. We
need also to be mindful of retention of the information to maintain mastery as a student works through the course
over the 8 or 9 months before the examinations.
 Mastery learning develops a growth mindset as the implication is that all students will understand all the work. The
variable is the time it takes to master the different concepts.
 In these AS slide decks, mastery is encouraged by different styles of questioning, practice questions and relies of
constant formative assessment by the teacher. Only move on once everyone in the class is getting over 80% of
the problems correct.
 Every lesson or at the end of a new section develop a habit of reviewing previous learning and linking new ideas
back to previous content.
 Retention is enhanced through spaced repetition in the starter activities, SRQ worksheets and the allocation of
past paper questions.
 The notes section of each slide provides links back to mastery learning ideas to provide the teacher with a
rationale of what this slide is aiming to achieve or test. It also includes answers and/or explanation of problems
and questions for the teacher to consider as their students progress.
 For additional support of questions related to mastery learning please email me at: content@chemcatalyst.co.uk

4. Retrieval Q+A
 What did you learn last lesson?
 What did you learn last week?
 What did you learn through an activity you organised yourself?
 independent reading, additional questions, research?

5. Modelling
 In pairs construct models to represent the structure of:
 ionic
 metallic
 simple covalent
 giant molecular

6. Starters for ten!
 A great worksheet from the RSC to remind students of learning from
unit 3

7. 4 States of matter
Physical Chemistry – CIE AS Chemistry

8. Subject content
 Learning outcomes 4.1 from the syllabus
1. explain the origin of pressure in a gas in terms of collisions between
gas molecules and the wall of the container
2. understand that ideal gases have zero particle volume and no
intermolecular forces of attraction
3. state and use the ideal gas equation pV = nRT in calculations,
including in the determination of Mr

9. Subject content
 Learning outcomes 4.2 from the syllabus
1. describe, in simple terms, the lattice structure of a crystalline solid which is:
a) giant ionic, including sodium chloride and magnesium oxide
b) simple molecular, including iodine, buckminsterfullerene C60 and ice
c) giant molecular, including silicon(IV) oxide, graphite and diamond
d) giant metallic, including copper
2. describe, interpret and predict the effect of different types of structure and
bonding on the physical properties of substances, including melting point,
boiling point, electrical conductivity and solubility
3. deduce the type of structure and bonding present in a substance from given
information

10. Expected learning gains
 An understanding of different types of structure
 An ability to describe phase changes and pressure using the kinetic
model
 Understanding differences and similarities in ideal and real gases
 Fluency in using the general gas equation

11. Ideal Gas
4.1 The gaseous state: ideal and real gases and pV = nRT

12. The Kinetic Model
 We are all familiar with these diagrams, most of you have been
drawing them for years!
 Now we need to look closely at the moment when one phase
changes to another.

13. Vapour pressure
 In a closed container some
particles will natural evaporate
 The gas particles exert a pressure
on the inside of the container
 This is known as the vapour
pressure
 A substance with a high vapour
pressure is termed volatile
 ethanol is a good example
 (think you can smell it easily)

14. An ideal gas?
 An ideal gas is the gas we think about when we consider the kinetic
model
 Molecules have zero volume (N.B., NOT mass)
 zero intermolecular forces
 constant random motion
 elastic collisions (no energy is lost)
 Real gases do NOT behave quite the way an Ideal gas does but Ideal
gases are easier to think about and work out information about which is
CLOSE to the ‘real’ values

15. A real gas
 Moles of gas at low pressure are spread far apart – their volume is
negligible when taken as part of the whole volume
 When the temperature is high the particles have excess kinetic energy
so the intermolecular forces are insignificant
 However the volume of the particles and intermolecular
interactions become significant at high pressures and low
temperatures.
 But high temperatures and low pressures a real gas becomes
almost IDEAL.

16. General gas equation
 The general gas equation works for any volume or temperature
p = pressure in Pascal, Pa
V = volume (m3)
(1 m3 = 1000 dm3 = 1,000,000 cm3)
n = no. of moles
R = gas constant (8.31 J K-1 mol-1) (in your data booklet)
T = temperature in Kelvin

17. Kelvin
 Named after William Thomson,
1st Baron Kelvin who was a 19th
century Scottish mathematical
physicist
 It is an absolute temperature scale

18. Magnetic resonant imaging
machines use super
conducting magnets to create
a very large magnetic field
The wires of the
electromagnet are super
cooled to reduce resistance
and allow superconductivity
in the wires and reduce
energy consumption. This is
cooling is done by using
around 1500 litres of liquid
Helium at -270

19. Worked example
 What volume is needed to store 6 moles of helium
gas at 200kPa and 350K?

20. Relative Formula Mass
 The ideal gas equation allows an interesting re arrangement
 Mr = mRT/pV
 You are rearranging and substituting n for m/Mr

21. Worked problems
 Worked examples
 At atmospheric pressure (101 kPa) 0.12g of butane vaporises as the
temperature of a vessel is raised to 120oC.
I. How many moles of butane were vaporised?
II. What volume does this gas occupy on m3 and cm3?

22. Homework
 Use the following link to access the virtual lab
 http://jersey.uoregon.edu/vlab/
 (You may need to add this website to the exceptions list in advanced
settings in the Java control panel)

23. Past Paper Questions
 Paper 1, June 2003, Q6
 Paper 2, June 2004, Q1
 Paper 2, June 2006, Q2b
 Paper 1, Nov 2007, Q31
 Paper 21, June 2011, 1(d)
 Paper 11, June 2013, Q9
 Paper 13, Nov 2013, Q7
 Paper 43, Nov 2013, Q2 (b)

24. Practice with practical
 RSC Determination of relative atomic mass experiment

25. In the following multiple choice quiz, show all you working –
you may be asked to prove your calculation on the board
and explain your rationale

26. Types of structure
4.2 Bonding and structure

27. 1. Ionic structure
 In ionic substances thousands, millions or(most likely) billions of
positive and negative ions arrange themselves one after the other in
three dimensions
 This regular arrangement of alternating ions is known as a ionic lattice
 Between the ions are
strong electrostatic
attractions
 Ionic bonds

28. Imagine zooming in…
This pattern
keeps repeating
over and over
until the edge of
the crystal

29. Types of lattice
 Depending on the size of the ions different compounds can have
different types of lattice
 One of the most common is a simple cubic lattice where each sodium
ion is surround by 6 chlorine ions and vice versa
 However others exist such as body-centered cubic and face-centered
cubic

30. MgO and NaCl
 NaCl and MgO share the same ionic structure
 What charges do the individual ions have?
 How did ionic charge effect the strength of metallic bonds in metals?
 In which of these two salts would you find the strongest ionic bond?
Why?
 Can you imagine what effect this might have on the melting point if you
compared the two substances?
 NaCl = 801oC MgO = 2,852oC
 We’l expand on this difference again later this unit

31. 2. Simple Molecular Structure
A molecule could be considered to be atoms covalently bonded together but
more importantly it is a ‘finite’ unit – it has a beginning and an end
A molecule of methane
consists of:
4 hydrogen atoms
and 1 carbon atom
ONLY – i.e. 5 atoms
These molecules are held together by
weak intermolecular forces
(see Unit 3)

32. Simple molecular structure – I2
 eg. Iodine
It also forms a face centred cube

33. Simple molecular structure – C60
 eg. fullerenes
 Sometimes known as
‘Buckyballs’ the
buckminsterfullerene is a ball of
carbon atoms similar to a football
 It has a diameter of 0.7nm and
once again each ball is held
together by weak intermolecular
forces

34. Nanotube
 Can you remember graphite from IGCSE?
 A single sheet of graphite is known as
‘graphene’
 If you roll a sheet of graphene into a tube you
create nanotubes, depending on how you
roll them they can conduct electricity very well
due to a single free electron from each
carbon atom
 A single tube is very strong and again the
individual tubes are held together with weak
intermolecular forces

35. There are hundreds of thousands of fibres in each fibre
On fibre is one ten-thousandth the diameter of a typical human hair.

36. Allotropes
 Next we will look at the structure of diamond and graphite which are
made of carbon as well.
 All of these ‘different’ carbon based substances made from the same
element are called allotropes
 We can define an allotrope as:
An allotrope is a different form of the same element, in the same
state, with an alternative arrangement of atoms

37. Shared characteristics
 Ionic substances are held together by ionic bonds and have repeating
units that continue throughout the substance – infinitely almost
 Simple molecular has discrete units and covalent bonds
 The next type of structure has a little of both
 Giant molecular has repeat units and covalent bonds

38. 3. Giant molecular - graphene
 Nanotubes are individual units but the sheets of graphene from which
they are made are giant covalent structures which can be much bigger

39. 3. Giant molecular - graphite
 Graphene can be harvested (using sticky tape!) from graphite which is
a type of metamorphic rock found in the Earth’s crust
 It consists of layers of carbon sheets with layers of delocalised
electrons in between

40. 3. Giant molecular - diamond
 Diamond has a giant
tetrahedral structure
 This is a very stable
and strong structure

41. 3. Giant molecular – silicon dioxide
 The structure of SiO2 is
very similar to diamond
 However in between each
pair of silicon atoms you
will find an oxygen atom

42. 4. Hydrogen bonded structure - ice
 In Unit 3 we looked at
hydrogen bonds a type
of intermolecular
attraction
 Some structures, like ice,
are held together by these
‘bonds’

43. 5. Metallic structure –
eg. copper
 Many metals have a closed packed
structure which is the most efficient way
for ‘sphere’s’ to pack together
 This can be a little hard to visualise so
watch this video to help

44. Match the
substance to
the structure
Copper
Nanotubes
Calcium Oxide
Ice
Chlorine
Fullerenes
Potassium Chloride
Zinc
Silicon Dioxide
Graphene
Diamond
Graphite
Ionic Lattice
Simple
Molecular
Giant
molecular/
covalent
Hydrogen
bonded
Metallic

45. Physical Properties
4.2 Bonding and structure

46. How structure effects physical properties
Melting
points
High melting and
boiling points due
to strong
electrostatic
attractions (ionic
bonds)
Smaller ions are
closer packed so
have greater
forces, hence a
higher melting
point
The greater the
charge on the ion,
the greater the
attraction and the
higher the boiling
point
Conductivity:
solid
Ionic solids DO
NOT conduct
electricity
There are no
delocalised or
free electrons to
allow a current to
flow
Conductivity:
liquid
Molten Ionic
compounds DO
conduct electricity
As a liquid the
ions are free to
move towards the
electrodes to be
oxidised or
reduce. A current
is free to flow
Conductivity:
aq. solution
Aqueous
solutions of ionic
compounds DO
conduct electricity
In a solution the
ions are free to
move towards the
electrodes to be
oxidised or
reduced. A
current is free to
flow
Solubility in
water
MOST Ionic
solids are soluble
in water
Assuming
dissolution is
energetically
favourable the
ionic solid will
dissolve in water
Solubility in
hexane
MOST Ionic
solids are
INSOLUBLE in
hexane
The bonds which
would form
between the non-
polar solvent and
the ions are not
strong enough to
pull apart the
ionic lattice

47. How structure effects physical properties
Melting
points
Low melting and
boiling points due
to weak
intermolecular
forces
A small amount of
energy is needed
to overcome
these forces
between the
molecules
Conductivity:
solid
DO NOT conduct
electricity
All electrons are
used in covalent
bonds, none are
delocalised
Conductivity:
liquid
DO NOT conduct
electricity
Even if some free
electrons are
present the
distance between
the molecules is
often too great for
electrons to flow
Conductivity:
aq. solution
n/a
Most simple
organic molecule
are insoluble in
water
Solubility in
water
INSOLUBLE in
polar solvents
The hydrogen
bonds between
the water
molecules are too
strong to be
replaced
Solubility in
hexane
SOLUBLE in
non-polar/organic
solvents
Similar
Intermolecular
forces allow the
substances to mix

48. Melting
points
Very high melting
and boiling point
Strong covalent
bonds acting in all
directions
Conductivity:
solid
DO NOT conduct
electricity
(graphite is an
exception)
The electrons are
used in covalent
bonds, none are
delocalised
Graphite
conducts due to
the layers of
delocalised
electrons
Conductivity:
liquid
DO NOT conduct
electricity
Very difficult to
melt
Conductivity:
aq. solution
n/a
Most giant
covalent
structures are
insoluble in water
Solubility in
water
INSOLUBLE in
polar solvents
The covalent
bonds between
the atoms are too
strong
Solubility in
hexane
INSOLUBLE in
non-polar
solvents
The covalent
bonds between
the atoms are too
strong
How structure effects physical properties

49. How structure effects physical properties
Melting
points
Very high melting
and boiling point
(except Grp 1)
Strong electrostatic
attractions between
the ions and the
delocalised
electrons
The greater the
charge the stronger
the attraction
Conductivity:
solid
DO conduct
electricity
The delocalised
electrons are
mobile and allow a
current to pass
Conductivity:
liquid
DO conduct
electricity
Once again
electrons are free
to move (under an
applied potential
and can conduct
electricity)
Conductivity:
aq. solution
n/a
Metals are
insoluble in water
as the electrostatic
attraction is too
great to overcome
Solubility in
water
INSOLUBLE in
polar solvents
The electrostatic
attraction is too
great
Solubility in
hexane
INSOLUBLE in
non-polar solvents
The electrostatic
attraction is too
great

50. Starters for 10
 Starters for 10 - Worksheet 3.3

51. Why ionic solids dissolve
 Breaking bonds requires energy.
 The bonds created with the water molecules when an ionic solid
dissolves must release enough energy to be energetically favourable.
 Otherwise the ionic solid will not dissolve.

52. What bonds are created?
Dative covalent bonds can
form between the Oxygen’s
lone pairs and the positive
ion
The Cl- ion is stabilised by
H-bonds formed from the δ-
side of the water molecules

53. In your own words
1. Why is potassium is a better conductor than sodium but have a lower
melting point?
2. When can the salt potassium iodide conduct electricity?
3. If Iodine atoms in iodine molecules are joined together by strong
covalent bonds why does Iodine have a low melting point?
4. Why does it dissolve in hexane?
5. Can it conduct electricity?

54. What can you remember?
IONIC SIMPLE
MOLECULAR
GIANT
COVALENT
METALLIC
Melting point
Conductivity when solid
Conductivity when liquid
Conductivity in aq
solution
Solubility in water
Solubility in hexane.
Copy and complete this table with either ‘YES’ or ‘NO’ or ‘HIGH/LOW’
Without looking back in
your notes!

55. End of unit recent exam Practice
 Answer the following Paper 1 and Paper 2 Qs from recent exams
 For Paper 2 Qs answer the whole question
Paper Session Year Question
21 May 2017 2
21 May 2019 3
11 May 2019 6
11 May 2019 7
12 May 2019 5
12 May 2019 6
12 May 2019 7
13 May 2019 4
13 May 2019 7

56. Which of these ideas has been difficult?
 Ideal gas vs. real gases
 The Ideal gas equation/calculations
 Types of structure
 ionic, simple molecular, giant molecular, hydrogen bonded, metallic
 Physical properties of different structures

57. Create 5 summary sentences using the words above to help you

58. Homework – Independent learner
 Read a recent newspaper article on an environmental issue and try to
find an alternative source of information on the same issue
 Compare the content and style of the pieces.
 Write a 100 word summary giving details of the sources of information.

59. Teacher self-evaluation
 Changes to be made to PowerPoint, teaching, resources?
 Ask the students for feedback