Periodicity
Mr Field
Using this slide show
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The slide show is here to provide structure to the lessons, but not
to limit them….go off-piste when you need to!
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Slide shows should be shared with students (preferable electronic to
save paper) and they should add their own notes as they go along.
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A good tip for students to improve understanding of the
calculations is to get them to highlight numbers in the question and
through the maths in different colours so they can see where
numbers are coming from and going to.

The slide show is designed for my teaching style, and contains only
the bare minimum of explanation, which I will elaborate on as I
present it. Please adapt it to your teaching style, and add any notes
that you feel necessary.
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Menu:
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Lesson 1 – The Periodic Table
Lesson 2 – Physical Properties
Lesson 3 – Chemical Trends
Lesson 4 – The Period 3 Oxides
Lesson 5 – HL – The Period 3 Oxides and Chlorides
Lesson 6 – HL – Transition Metals - Introduction
Lesson 7 – HL – Coloured Complexes and Catalysts
Lesson 8 – Test
Lesson 9 – Test Debrief
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Lesson 1
The Periodic Table
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Overview
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Copy this onto an A4 page. You should add to it as a
regular review throughout the unit.
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Assessment
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This unit will be assessed by:
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A test at the end of the topic (100%)…around Lesson 8
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We Are Here
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Lesson 1: The Periodic Table
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Objectives:
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Reflect on prior knowledge of the periodic table
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Understand the structure and purpose of the periodic table
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Repeat the work of Mendeleev by constructing your own
periodic table
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Reflecting on the Periodic Table
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What is the periodic table and what is supposed to show?
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The Traditional
Based on Mendeleev’s work. Easiest to use and display.
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Dmitri Mendeleev’s Periodic Table
The one that started it all off.
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Wide Format Periodic Table
Shows true position of the f-block (lanthanides and actinides)
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Janet Periodic Table
Elements arranged in order of orbital filling. Used frequently by
physicists.
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Benfey Periodic Table
Spiral form shows the steady increase in atomic number.
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Stowe Periodic Table
Emphasises the symmetrical nature of the increase in quantum
numbers.
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Zymaczynski Periodic Table
Another way to show the symmetry in the underlying quantum
numbers.
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Giguere Periodic Table
A 3D representation emphasising the s, p, d and f blocks
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The Structure of the Periodic Table
GROUPS
PERIODS
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Groups and Periods
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Groups
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Elements show similar chemical properties
Elements show similar trends in their chemical properties
Periods
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As you move across periods, changes in the chemical and
physical properties that are repeated in the next period
This is what ‘period’ and ‘periodic’ refers to
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The periodic table and electron
configuration
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How does an element’s position in the PT relate to its
electron configuration?
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Being Mendeleev
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The first widely accepted periodic
table was produced by the Russian
chemist Dmitri Mendeleev
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It was a tremendous example of
scientists as risk-takers as it was able
to make a number of predictions
thought unlikely at the time
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Complete the exercise here in which
you will use the information available
to Mendeleev to construct your own
periodic table
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Homework – HL only!
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In two groups, you need to work together to plan a
30 minute lesson on the objectives detailed below to
be delivered in Lesson 5.
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The lesson should include:
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A presentation
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Complete with equations, diagrams, explanations etc
An activity
A review
Allocations:
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Oxides: Na2O,
MgO, Al2O3, SiO2,
P4O6 and P4O10,
SO2 and SO3, Cl2O
and Cl2O7
Group 1: Explain the physical states (under standard
conditions) and electrical conductivity (in the molten
state) of the chlorides and oxides of the elements in
period 3 in terms of their bonding and structure.
Group 2: Describe the reactions of chlorine and the
chlorides referred to above with water. State and explain
the acidity of the resulting solutions
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Chlorides: NaCl,
MgCl2, Al2Cl6, SiCl4,
PCl3 and PCl5,
and Cl2
SUPER IMPORTANT
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(If you have one) bring a laptop with spreadsheet program
to the next lesson.
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If you don’t have Microsoft Excel, download OpenOffice
for free from:
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www.openoffice.org
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Key Points
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The periodic table arranges the elements according to:
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Their chemical properties
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Their electronic structure
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Lesson 2
Physical Properties
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Refresh
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Nitrogen and silicon belong to different groups in the
periodic table.
a)
Distinguish in terms of electronic structure, between
the terms group and period.
b)
State the maximum number of orbitals in the n = 2
energy level.
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We Are Here
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Lesson 2: Physical Properties
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Objectives:
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Identify and explain the trends in the physical properties of the
first 20 elements including:
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Atomic radius
Ionic radius
First ionisation energy
Electronegativity
Melting point
Use Microsoft Excel to produce a spreadsheet to graph the
above physical data
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Atomic Radius
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This is the ‘size’ of an atom
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There is no simple measure as atoms do not have a well defined ‘edge’
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We use the: covalent radius
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Values range from are measured in picometres ( 1 pm = 1x10-12 metres…a
thousand-billionth of a metre) and range over:
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This is half the distance between the nuclei of two atoms in a covalent bond
This means we don’t have values for the noble gases as they do not form bonds
270 picometres Francium
30 picometres for Hydrogen (helium would be smaller but does not form covalent bonds to
be measured)
The main factors influencing atomic radius are:
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Number of shells (the principal quantum number)
The charge in the nucleus
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Ionic Radius
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This is the ‘size’ of an ion and is measured in a similar way
to atomic radius
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It is measured in a picometres with values ranging over:
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272 pm for the Ge4- ion
16 pm for the B3+ ion
The main factors influencing ionic radius are:
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Number of shells (the principal quantum number)…don’t forget
this can be affected by the type of ion formed
The charge in the nucleus
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First ionisation energy
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This is the energy required to remove one mole of electrons from one
mole of gaseous atoms to form positive ions i.e.:
A(g)  A+(g) + e-
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Values range over:
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393 kJ mol-1 for Caesium
1681 kJ mol-1 for Helium
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Values are positive because this is an endothermic process
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Values are influenced by:
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Number of inner electron shells (and their shielding)
Charge on the nucleus
HL: At the finest level – repulsion between electrons in their orbitals
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Electronegativity
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This is a measure of the degree to which an element attracts
the shared pair of electrons in a covalent bond
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Values range over:
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4.0 for Fluorine
0.7 for Francium
Values are influenced by:
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Again, this means there are no values for the noble gases
Number of inner electron shells (and their shielding)
Charge on the nucleus
Values are unit-less as this is a relative measure
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Melting Point
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This is the temperature (in Kelvin…i.e. Celsius + 273) at
which an element melts
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Values range over:
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3935 K for Carbon
1 K for Helium
Values are influenced by:
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Nature of bonding: giant covalent, giant ionic, metallic
Strength of bonding
Strength of intermolecular forces
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Trends in Physical Properties
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You need to produce an Excel spreadsheet to help you analyse
the physical data.
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Use the blank here and follow the instructions on the instructions
page
Once you have done this you need to use this to help you
identify and explain the following trends:
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Atomic and ionic radius, first ionisation energy, electronegativity and
melting point
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Atomic and ionic radius, first ionisation energy, electronegativity
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Down group I (alkali metals)
Down group VII (halogens)
Across period 3
The general trend in electronegativity over the whole PT
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Key Points
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Each of the following physical parameters follow trends
and patterns in the PT
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These patterns are generally explained by:
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Charge in the nucleus
Number of electron shells
Electron shielding
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Lesson 3
Chemical Properties
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Refresh
Which species has the largest radius? Do not use the data
booklet…work it out!
A.
B.
C.
D.
Cl–
K
Na+
K+
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We Are Here
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Lesson 3: Chemical Properties
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Objectives:
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Understand the following trends in reactivity:
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Alkali metals with water
Alkali metals with halogens
Halogens with halide ions
Complete an experiment to investigate the above
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Chemical Trends
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Members of a group often have very similar reactivity.
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You probably know that carbon will react with hydrogen
to form methane, CH4
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You probably did not know that silicon will also react
with hydrogen to form silane, SiH4
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Watch this demonstration to see some silane being made
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Three reactions to know
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The Group I (alkali) metals react with water as follows:
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The Group I (alkali) metals react with halogens (Group VII) as
follows:
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Metal + Water  Metal Hydroxide + Hydrogen
Metal + Halogen  Metal Halide
Halogens can react with halide ions as follows (using the
example of bromide and chlorine):
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Bromide + Chlorine  Chloride + Bromine
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Investigating chemical trends
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In this experiment you will investigate trends in the
reactions mentioned on the previous slide
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Follow the instructions here
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Key Points
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Alkali metals become more reactive down the group:
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Due to the outer shell electron becoming increasingly easy to
remove
Halogens become less reactive down the group:
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Due to the increased numbers of electron shells (and thus
shielding) causing them to attract electrons less strongly
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Lesson 4
Period 3 Oxides
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Refresh
Which properties of the alkali metals decrease going down
group 1?
A.
B.
C.
D.
First ionization energy and reactivity
Melting point and atomic radius
Reactivity and electronegativity
First ionization energy and melting point
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We Are Here
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Lesson 4: Period 3 Oxides
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Objectives:
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Understand and explain the trend in acid-base behaviour of the
period 3 oxides
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Complete an experiment to demonstrate the amphoteric
nature of aluminium oxide
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The Period 3 Oxides
Element
Formula of
oxide
Reaction of oxide with water
Acid/base nature
Sodium*
Na2O
Na2O + H2O  2NaOH
Strongly basic
Magnesium*
MgO
Slight: MgO + H2O  Mg(OH)2
Weakly basic
Aluminium
Al2O3
Amphoteric
Silicon
SiO2
Very weakly acidic
Phosphorous*
P4O10
Sulphur*
SO2
SO3
Chlorine
Argon
no direct
reaction but:
Cl2O7
P4O10 + 6 H2O  4 H3PO4
Strongly acidic
Strongly acidic
SO3 + H2O  H2SO4
Strongly acidic
Cl2O7 + H2O  2 HClO4
no oxides
There is a gradual transition from basic to acidic character, reflecting a
gradual transition from metallic to non-metallic nature
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Note: you will only be tested on the elements marked with an asterisk, *
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Amphoteric Aluminium
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Complete the amphoteric aluminium experiment here.
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This goes beyond the requirements of the syllabus but
will help deepen your knowledge and understanding.
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Homework
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Research the role of acidic oxides in the formation of acid
rain.
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Include:
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Sources of acidic oxides
Names and formulas and their reactions with water
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Key Points
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The oxides of period 3 display a gradual transition
basic to acidic character
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This reflects a gradual transition from metallic to
non-metallic nature of the elements
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Lesson 5
Period 3 Oxides and Chlorides
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Refresh
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Which oxides produce an acidic solution when added to
water?
I.
II.
III.
A.
B.
C.
D.
P4O10
MgO
SO3
I and II only
I and III only
II and III only
I, II and III
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We Are Here
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Lesson 5: Period 3 Oxides
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Objectives:
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Complete the amphoteric aluminium activity
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Deliver the lessons set for homework in the first lesson of the
topic
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Amphoteric Aluminium
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Complete the experimental work for the amphoteric
aluminium activity – you have 25 minutes
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Over to you
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Welcome to my world!
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Homework
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Complete the analysis for the amphoteric aluminium
activity
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Lesson 6
Transition Metal Complexes - Introduction
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Refresh
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By reference to the structure and bonding in NaCl and
SiCl4:
a)
State and explain the differences in electrical conductivity in
the liquid state.
b)
Predict an approximate pH value for the solutions formed by
adding each compound separately to water. Explain your
answer.
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We Are Here
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Lesson 6: Transition Metal Complexes Introduction
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Objectives:
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Describe the properties of transition metals
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Understand the term ligands
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Understand and explain the formation of transition metal
complexes
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The Transition Metals
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A transition metal is an element in which at least one ion
has a partially filled d-orbital
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For example, Cu2+: 1s2 2s2 2p2 3s2 3p2 (4s0) 3d9
Properties of the transition metals include:
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Variable oxidation states (for example iron: Fe2+, Fe3+, Fe6+)
Formation of coloured compounds (more later)
Catalytic properties (more later)
Formation of complex ions (much more later)
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Scandium and Zinc
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Although in the first row of the d-block,
these are not transition metals.
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To understand why, write the full
electron configuration for:
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Sc and Sc3+
Zn and Zn2+
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Variable oxidation numbers (ions)
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Transition metals have large numbers of electrons in d-orbitals,
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Some common oxidation states we need to know:
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This means the amount of energy required to remove the second
electron is not much different to that required to remove the first
and so on.
All of them in the +2 oxidation state
Cr(III), Cr(VI)
Mn(IV), Mn(VII)
Fe(III)
Cu(I)
Task: select 4 of these and write the electron configuration
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Ligands
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A ligand is a species with a
lone pair
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Often negative ions
Common ligands include:
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Water, H2O
Ammonia, NH3
Chloride, ClHydroxide, OHCyanide, CNThiocyanate, SCNMain Menu
Transition Metal Complexes
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The lone pair on a ligand can form a dative covalent bond to a
metal ion to form a transition metal complex
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[Fe(H2O)6]3+
[Fe(CN)6]3-
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[Cu(Cl)4]2-
[Ag(NH3)2]+
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Making transition metal complexes
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Add potassium thiocyanate solution to a solution of iron (III)
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Add conc. HCl (fume hood!) to 1 cm3 of a strong solution of
cobalt (II). Repeat but use conc. NH3 instead (fume hood!).
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Add dilute NH3 to a copper (II) solution until no further
change occurs
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Record all observations
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Suggest possible structures for the complexes you have formed and
possible reaction equations
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A Challenge
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Working in small groups, complete the following activity
on the structure of some cobalt complexes.
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Key Points
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Transition metals form ions with partially filled d-orbitals
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Ligands are species with lone pairs
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Ligands will form dative covalent bonds to transition
metals forming ‘complex ions’
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Lesson 7
Complex Colours and Catalysts
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Refresh
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By reference to the structure and bonding in NaCl and
SiCl4:
a)
State and explain the differences in electrical conductivity in
the liquid state.
b)
Predict an approximate pH value for the solutions formed by
adding each compound separately to water. Explain your
answer.
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We Are Here
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Lesson 7: Complex Colours and Catalysts
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Objectives:
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Complete the amphoteric aluminium activity
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Deliver the lessons set for homework in the first lesson of the
topic
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Marketplace – in three groups
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Each group needs to produce a learning resource to teach the other students about
their chosen topic.
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Once the resources are completed, one person should remain with the resource
whilst the remaining members circulate and learn from the other stations….you
should manage your time, taking turns manning your station to make sure everyone
makes it round class.
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The catalysts topics should include the names of the elements/compounds, an
equation for the reaction they do and the importance of the catalysis.
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There will be a test at the end.
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Topic allocations:
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Origin of colour in transition metal complexes (d-orbital splitting)
Catalysts 1: Transition metal
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Catalytic converters, conversion of alkenes to alkanes, the Haber process
Catalysts 2: Transition metal compounds:
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Decomposition of hydrogen peroxide, Contact process, haemoglobin, vitamin B12
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Test time
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You have 10 minutes
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Mwahahahahahahahahaha
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Key Points
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The formation of complexes causes d-orbitals to split
into two energy levels
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Electron transitions between these energy levels give rise to
their colour
Transition metals are hugely important for their catalytic
properties
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Lesson 8
Test
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Good Luck
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You have 80 minutes!
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Lesson 9
Test Debrief
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Personal Reflection
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Spend 15 minutes looking through your test:
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Make a list of the things you did well
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Use your notes and text book to make corrections to
anything you struggled with.
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Group Reflection
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Spend 10 minutes working with your classmates:
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Help classmates them with corrections they were unable to do
alone
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Ask classmates for support on questions you were unable to
correct
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Go Through The Paper
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Stop me when I reach a question you still have difficulty
with.
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Targeted Lesson
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PREPARE AFTER MARKING THE TEST
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SHORT LESSON ON SPECIFIC AREAS OF DIFFICULTY
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