Contents
 The Periodic Table
 Physical Properties
 Chemical Properties
Objectives:
 Describe the arrangement of elements in the
periodic table in order of increasing atomic
number
 Distinguish between the electron arrangement of
elements and their position in the periodic table
up to Z = 20
 Apply the relationship between the number of
electrons in the highest occupied energy level for
an element and it’s position in the periodic table.
Activity
 Watch video from Science Bank 3 DVD Disc 2 on;
 The Periodic Table
 Electronic Structure
Objectives:
 Describe basic trends in electron arrangement,
structure and electrical conductivity in the periodic
table.
 Describe and explain the trends in melting points for
the alkali metals (Li → Cs) and the halogens (F→ I)
 Describe and explain the trends in melting points for
elements across period three.
Periodicity and properties
 Periodicity is explained by the electron structures
of the elements:
 Group 1, 2 and 3 are metals, with up to three electrons
in their outer shells. They have giant metallic
structures. They give away their outer electrons (up to
three) to form ionic compounds.
 Group 4, with four electrons in their outer shells, are
semi-metals. They have giant structures. They only form
covalent compounds.
 Group 5 to 7 are non-metals with five, six or seven
electrons in their outer shells. They either accept
electrons (up to three) to form ionic compounds, or
share their outer electrons to form covalent compounds.
 Group 0 are the noble gases – atoms that have full outer
shells and are unreactive.
Metallic Structure
 In metallic bonding each metal atom ‘loses’ its
outer (valence) electrons. Forming cations
surrounded by a ‘sea’ of delocalised electrons.
 The greater the number of valence electrons in
the metal, the greater the strength of the
metallic bond.
+
+
-
+
+
-
+
+
-
+
+
-
- - - - - - --
2+
2+
2+
2+
- - - - - - --
2+
2+
2+
2+
Giant Covalent Structure
 Group 4 elements do not form metallic bonds,
instead they tend to share electrons to obtain a
full outer shell of electrons, this type of bonding
is called Covalent bonding.
 In Giant Covalent structures the strong covalent
bonds occur in all directions making this type of
structure particularly strong.
 Diamond is a classic example of a giant covalent
structure.
Simple covalent structures.
 Elements in Groups 5 to 7 form simple covalent
molecules. This means that the atoms in each molecule
are bonded covalently, these are strong bonds formed
from the sharing of electrons.
 The molecules are held together by weak forces of
attraction. These forces are called van der Waals
forces and are produced by the electrons in the
molecules.
 The bigger the atoms in the molecule or the more
atoms in the molecule, the greater the number of
electrons and thus the greater the strength of the van
der Waals forces and the higher the melting and boiling
point of the substance.
Van der Waals
 Van der Waals forces are caused by the electrons
in a molecule.
 As molecules move the electrons get dragged
behind them, as they are much lighter than the
heavy nucleus.
 This means that a temporary imbalance of
electrons exists forming a temporary polarity
(this means that for a brief moment in time the
molecule will have a slightly positive end and a
slightly negative end).
 This temporary polarity (dipole) causes a brief
electrostatic attraction between one molecule and
another. This is a Van der Waals force.
Activity;
 Analyse the graph from the worksheet ‘Trends in
melting points and boiling points of Period 3
Elements’ then describe and explain this trend
using your knowledge of bonding.
Trends in the melting and boiling points of the
period 3 elements.
Giant covalent bonding, there
are strong covalent bonds in
each direction.
Melting and boiling points/K
3000
Metallic bonding,
2500
cations in a ‘sea’of
delocalised
electrons. As2000
the
number of outer
1500
electrons (valence
electrons) increase
the strength1000
of the
metallic bonding
500
increases.
Simple covalent
molecules. There are
weak intramolecular
forces called ;van der
Waals’ between the
simple molecules. The
strength of the van der
waals dependsTm/K
on the
size of the molecule.
Tb/K
The bigger the
molcule
the greater the van
der Waals.
0
11
12
13
14
15
16
Atomic number
only there are
17 Atomic
18
only very weak forces of
attraction between the
particles
Comparing Tm and Tb trends
down groups 1 and 7
Trends in Tm and Tb for Group 7
1800
700
1600
1400
600
1200
1000
Tm / K
800
600
Tb / K
Tm and Tb / K
Tm and Tb / K
Trends in Tm and Tb down Group 1
500
400
Tm / K
300
Tb / K
200
400
200
100
0
0
3
11
19
37
Atomic number
55
87
9
17
35
53
85
Atomic number
Notice the difference in the trends as we go down each group,
can you explain it?
Explaining Tm and Tb trends
down groups 1 and 7
decrease
 In Group 1 the melting points __________
down
the group.
bigger
 This is because, the atoms get ________
down
the group and so the forces of attraction
stronger
between them become ________.
 In Group 7 the melting points show the opposite
increase
trend, and _________
down the group.
 This is because the solid crystals of the halogens
contain non-polar diatomic molecules, which are
only attracted to each other by weak forces, the
strength of these forces increase with the mass
of the molecule (chapter 4 - bonding)
Activity
 Reread through your notes including page 47- 48
of the Chemistry Course Companion and answer
the question from the worksheet; ‘Melting point
trends in Groups 1 & 7 & Period 3’
 Find the answers by clicking here.
Trends in physical properties
 Electrical conductivity
left
 The pattern is good conductivity on the ________
and poor
poor
conductivity on the _________.
etallic
 This trend is due to bonding. M_________
bonding on the
left and c__________
right
ovalent
______
bonding on the _______.
 Melting point and boiling point
 There is a clear break in the middle of the table between
left
elements with high melting points (on the ______)
and those
right
with low melting points (on the ______).
 These trends are due to giant ionic lattice structures on the
left and simple molecular structures on the _______
right
_____
Atomic Radius
Objectives:
 Describe and explain the trends in atomic radii
for the alkali metals and the halogens
 Describe and explain the trends in atomic radii
for elements across period three.
Measuring atomic radii
 The atomic radius should be defined as the
distance from the centre of the nucleus to
the outermost electron. This is not
possible, why?
 What can be measured, by a technique
known as X-ray diffraction, is half the
distance between two nuclei of bonded
atoms.
 If they are bonded covalently it is known
as the covalent radius of an atom, whereas
if metallic bonding is involved it is the
metallic radius.
Atomic properties
Some key properties of atoms, such as size and ionisation
energy are periodic.
Activity: (in pairs)
 Look up the atomic radii of the elements on page 7 of the
book of data.
 Think of a suitable scale and draw scale drawings of the
atoms down Group 1 (don’t include Hydrogen or Francium)
and along Period 3, using compasses.
 Label each atom with its name, symbol and actual radius.
 Use a key to show the scaling you have used.
 Describe the trend that you can see; down a group and
along a period. Can you explain it?
Trends in atomic radii
 Atoms
get
larger
as we
go
down
any
group.
 And get smaller as you
go along the period
from left to right.
Trends in atomic radii
 Atomic radius is a
periodic property , it
decreases across
each period and
where there is a jump
we start the next
period.
Down a group
 As we go down group 1 in the periodic table, each
element has one extra complete shell of electrons
than the one before.
 So, for example, the outer electron in potassium
is in shell 4, while that in sodium is in shell 3.
 Both outer electrons experience the same
shielded nuclear charge of 1+.
 Shielded nuclear charge refers to the charge
from the resultant nuclear charge on the
electrons in the outer shell.
Activity:
Using your data booklet and your notes fill in the
information in the table below. Can you explain
the trend in atomic radius along period 3?
Na
Electron
arrangement
Atomic radius/
nm
Nuclear charge
Effective
nuclear charge
Mg
Al
Si
P
S
Cl
Activity:
Using your data booklet and your notes fill in the
information in the table below. Can you explain
the trend in atomic radius along period 3?
Electron
arrangement
Atomic radius/
nm
Nuclear charge
Effective
nuclear charge
Na
Mg
Al
Si
P
S
2,8,1
2,8,2
2,8,3
2,8,4
2,8,5
2,8,6
Cl
2,8,7
The Shielding effect
 The inner shells of electrons shield the outer
electrons from the pull of the nucleus.
 This doesn’t only affect the atomic radii but also
affects reactivity as you will see in the next
topic.
Across a period
 As we move from sodium to chlorine we are adding
protons to the nucleus and electrons to the outer
shell – the third shell.
 The charge on the nucleus increases from 11+ to
17+ (or +1 to +7 after allowing for shielding of the
inner shells, which remains the same).
 This increased charge pulls the electrons in closer
to the nucleus.
 So the size of the atom decreases as we go
across a period.
Ionic Properties
Objectives:
 Describe and explain the trends in ionic radii for
the alkali metals and the halogens
 Describe and explain the trends in ionic radii for
elements across period three.
Ionic properties
Activity:
 Look up the ionic radii of the elements on page 7 of the
book of data.
 Think of a suitable scale and draw scale drawings of the
atoms down Group 1 (don’t include Hydrogen or Francium)
and along Period 3, using compasses.
 Label each atom with its name, symbol and actual radius.
 Use a key to show the scaling you have used.
 Describe the trend that you can see; down a group and
along a period. Can you explain it?
 Compare these to the atomic radii. Can you explain the
differences?
Trends in ionic properties

Positive ions (cations) vs atoms
When an atom of a Group 1 element loses an
electron, the ion that is formed has a much
smaller radius – almost half the value. There are
two reasons for this.
1. One fewer electrons than there are protons, so
the nucleus attracts the remaining electrons much
more strongly.
2. One fewer energy level, because the outer shell
has effectively been removed, and the remaining
electrons have the noble gas electron
arrangement of the preceding element

Trends in ionic radii down a
group
 The size of the
ions _______
down group 1 as
the outer energy
level becomes
progressively
_________ from
the nucleus.
How to compare ionic radii
along a period
 Not all atoms form unipositive ions so we cannot make
a direct comparison of ionic radii, but we can compare
ions with the same number of electrons these are
called iso-electronic ions.
 Sodium ions, Na+, magnesium ions, Mg2+, and aluminium
ions, Al3+, all contain 10 electrons and have the
electron configuration of neon (2,8), with the second
shell completely full, these are iso-electronic ions
Trend in ionic radii along the
metals in Period 3
 As we move along the metals in period 3 the iso-
electronic ions become much smaller. Can you think
why?
 Sodium has __ protons in the nucleus, magnesium has
__ and aluminium __. The __ protons in the aluminium
nucleus will attract the eight electrons in its outer
shell ____ than the __ protons in the nucleus of the
sodium ion.
Trends in ionic properties
 Negative ions (anions)
 When the atoms of elements in group 7 gain one
electron to form a negative ion, there will be one
more electron in the outer shell and hence more
electron-electron repulsion.
 The number of protons in the nucleus is unchanged,
each of the electrons will be attracted less strongly,
and the radius of the ion increases to almost twice
the radius of the atom.
Trends in ionic radii down a
group
 The size of the
ions _______
down group 7 as
the outer energy
level becomes
progressively
_________ from
the nucleus.
Trend in ionic radii along the
non-metals in Period 3
 As we move from metals to non-metals in Period 3 the
iso-electronic ions become much _______, but from
right to left they get __________. Can you think why?
 Compare the electron arrangements and protons of
silicon and phosphorus.
Trend in ionic radii along the
metals in Period 3
 Silicon has ___ protons and ____ electrons, phosphorus
has ___protons and ___ electrons. Silicon has the
electron arrangement _____ and phosphorus has the
arrangement ______. There will be a greater pull on the
outer electrons in __________ which makes the ion
much _________.
 Phosphorus has __ protons in the nucleus, sulphur has __
and chlorine __. The __ protons in the chlorine nucleus
will attract the eight electrons in its outer shell ____
than the __ protons in the nucleus of the phosphorus ion.
Ionisation Energies
Objectives:
 Define the terms first ionisation energy.
 Describe and explain the trends in first ionisation
energies for the alkali metals and the halogens
 Describe and explain the trends in first ionisation
energies for elements across period three.
Ionisation Energies
 A graph of the first ionisation energies of the
elements against atomic number illustrates
periodicity very clearly.
 Can you state the definition for the first
ionisation energy of an element, can you write an
equation to show this?
The first ionisation energy of an element is the energy
required to remove one mole of electrons from one
mole of atoms of the element in the gaseous state.
M(g) → M+(g) + eIt is measured in Kilojoules per mole – kJ mol-1
Trends in first ionisation
energies
 Activity:
 Look at your graph of first ionisation energies.
 Describe (state) the trend in first ionisation
energies down groups 1 and 7 and across period 3.
 Using your knowledge of nuclear charge, shells, pull
of nucleus, electron arrangement try and explain
the trend in first ionisation energies down groups 1
and 7 and across period 3.
2500
2000
1500
First Ionisation
Energy / kJ mol
1000
500
Atomic Number
33
29
25
21
17
13
9
5
0
1
Ionisation Energy kJ/mol
A graph to show the First Ionisation Energies of
the first 36 elements.
Explanation of trend down a
group
 The elements in group 1 have the _______ values
in each period and the noble gases have the
________values.
 As we go down group 1 the values _______,
because the outer electron is ________ from the
nucleus and is therefore already in a higher
energy level, so ____ energy is required to
remove it.
 This is the same for all groups.
further
highest
lowest
less
increase
harder
decrease
more
easier
nearer
Explanation of trend across
Period 3.
 Across a period, as each energy level is
successively filled with electrons an equal number
of protons are also being added to the nucleus.
 As each electron is added the level (shell) is
attracted _____ to the nucleus and therefore it
becomes successively _______to remove an
electron, so that ionisation energies generally
________ across a period.
further
highest
lowest
less
increase
harder
decrease
more
easier
nearer
What does first ionisation
energy imply?
 Apart from Hydrogen most elements have more
than one electron, but will the energy required to
remove successive electrons be the same?
 The second ionisation energy of an element is
defined as the energy required to remove one
mole of electrons from one mole of unipositive
ions in the gaseous state.
M+(g) → M2+(g) + e What would be the definition of the third
ionisation energy of an element?
A Graph to show the Successive Ionisation Energies of Sodium
To remove the last
2 electrons is
incredibly hard
180000
160000
140000
Em kJ/mol
120000
100000
80000
60000
40000
20000
0
1
To remove the
first electron is
relatively easy
2
3
4
5
6
7
Number of Electrons
To remove the next 8
electrons gets
progressively harder
8
9
10
11
Electronegativity
Objectives:
 Define the terms electronegativity
 Describe and explain the trends in
electronegativity for the alkali metals and the
halogens
 Describe and explain the trends in
electronegativity for elements across period
three.
 Compare the relative electronegativity values of
two or more elements based on their positions in
the periodic table.
What is electronegativity?
 The forces that hold atoms together are all about
the attraction between positive and negative
charges
 In ionic bonding we have complete transfer of
electrons from one atom to another.
 Even in covalent bonds, the electrons shared by
the atoms will not be evenly spread if one of the
atoms is better at attracting electrons than the
other.
 The ability of an atom to attract a bonding pair of
electrons to itself is known as electronegativity.
What is the trend?
 Electronegativity is a relative value not an absolute
value, and so there are different scales of
electronegativity in use.
 The values used by the IB are attributed to the North
American chemist Linus Pauling (1901 – 1994).
 The Pauling scale is used to
measure electronegativity. It
runs from 0 to 4. The greater
the number, the more
electronegative the atom.
Elements in groups 1
and 2 are often called
electropositive
Elements in groups 5,
6 and 7 are called
electronegative
 What is the most electronegative element on the
periodic table?
 What is the least electronegative element on the
periodic table?
Explaining electronegativity
 Two factors affect electronegativity:
 The size of the atom. The smaller the atom, the
larger the electronegativity (remember your trend
in atomic radius)!
 The shielded nuclear charge. The greater the
shielded nuclear charge, the larger the
electronegativity (remember how to calculate
shielded nuclear charge, proton number minus
number of inner electrons, e.g. shielded nuclear
charge of sodium = 11 – 10 = +1)
Chemical properties
Objectives:
 Discuss the similarities and differences in the
chemical properties of elements in the same
group.
 Discuss the changes in nature, from ionic to
covalent and from basic to acidic, of the oxides
across period three.
Activity
 Watch the video on Patterns of Reactivity of
Group I from Science Bank Disc 1.
Group 1
 The group 1 elements are all:
 Shiny
 Silvery
 React vigorously with water
 Stored under oil to stop them reacting with air.
 Soft (cut easily with a knife)
 Once cut they tarnish easily
 The reactivity increases down the group.
 Can you explain the reactivity of the alkali metals
down the group?
Reaction with water
 All the alkali metals react ______________
with
vigorously
water to form hydrogen and ions of the metal
hydroxide in water:
2M(s) + 2H2O(l) → 2M+(aq) + 2OH-(aq) + H2(g)
 Where M represents Li, Na, K, Rb or Cs
 Alkali metals are considered good reducing agents
as they readily donate an electron.
 Lithium reacts the most ________, and retains its shape




as it reacts.
The heat produced when sodium reacts is enough to
_____ it into a ball so that it darts about on the surface
of the water.
Occasionally the hydrogen released may ______ and
produce a _______flame characteristic of sodium ions.
Potassium is much ______ reactive, and the hydrogen
evolved usually burns steadily with a _______ flame
because of the presence of ___________ ions.
Rubidium is even more reactive, and the reaction with
caesium may be so violent that the glass vessel _______.
ignite
melt
slowly
potassium
explodes
less
yellow
red
lilac
more
Reactivity of Group I
 All the Group I elements have just one electron in
their outer shell, and when they react they lose
this outer electron to form the unipositive ion.
 The reactivity increases down the group as the
outer electron becomes successively easier to
remove, as the attraction from the nucleus
decreases.
Activity
 Watch the demonstration of sodium reacting with
chlorine and then sodium reacting with bromine to
form the sodium halides.
 Notice the appearance of the reactants and the
products.
Reaction with Halogens
 The alkali metals also react readily with the halogens,
with the reactivity of the metals increasing down the
group. If a piece of heated sodium is lowered into a
gas jar of chlorine it will burst into flames, and white
fumes will be seen, caused by the formation of the
ionic salt sodium chloride.
 Write the reaction between sodium and chlorine.
 Similarly, if bromine vapour or iodine vapour is passed
over heated sodium, a vigorous reaction occurs and
the two elements combine to produce sodium bromide
and sodium iodide respectively.
 Write the equations for these reactions.
The halogens
 The four halogens (F2), chlorine (Cl2), bromine (Br2) and




Iodine (I2) all exist as diatomic molecules.
When they react, the single bond between the atoms in the
halogen molecule is broken and the two atoms then each
gain one electron to form a halide ion.
The halogens are strong oxidising agents, they accept
electrons readily,
Fluorine is the most oxidizing and reactivity decreases
down the group.
The halogen gases are very poisonous and this was first
utilized in the first world war.
Oxidising effect of the
Halogens.
 In aqueous solution; chlorine is a clear, colourless solution,
bromine is pale yellow to orange, (depending on the
concentration of the bromine), similarly, iodine can vary
from pale yellow to brown. Consequently, it can be quite
difficult to distinguish between dilute, aqueous solutions of
bromine and iodine simply on their colour.
 As the halogens are non-polar they dissolve readily in
hydrocarbon solvent. This forms the basis of a useful test
to distinguish between the two halogens as the colours in
hydrocarbon solvent are very different. Iodine forms a
purple solution and bromine forms an orange solution.
Activity
 Copy this table into your books, then fill in as you
watch the video on the following slide.
Water
(control)
Chlorine
water (Cl2)
Bromine
water
(Br2)
Iodine
water (I2)
Sodium
chloride
Sodium
bromide
Sodium
iodide
 Look at the following
video and try to decide
whether a chemical
reaction has taken
place, write in your
table which halogen is
present in the
hydrocarbon solvent.
 We use hydrocarbon
solvent to distinguish
between the halogens
 Chlorine is clear,
colourless
 Bromine is orange
brown
 Iodine is pink
Water
(control)
Sodium
chloride
Sodium
bromide
Sodium
iodide
Chlorine
water (Cl2)
Cl2
Cl2
Br2
I2
Bromine
water (Br2)
Br2
Br2
Br2
I2
Iodine
water (I2)
I2
I2
I2
I2
 Write a word and balanced symbol equation for
each of the reactions in the table.
Chemical Properties of
elements in the same period
 There is a very noticeable change in the
properties of the elements across the period.
 Metals can be distinguished from non-metals by
their chemical properties.
 The oxides of metals tend to consist of the metal
ions and oxide ions, and so are said to be ionic.
 Write down the formulas of sodium oxide,
magnesium oxide and aluminium oxide.
 When they are in the liquid state they will
conduct electricity, and are decomposed to their
elements in the process.
 Silicon dioxide, SiO2, has a giant covalent
structure and so has a very high melting and
boiling point, but it does not conduct electricity
when molten.
 The oxides of the non-metals sulphur, phosphorus
and chlorine are all simple covalent, and have
relatively low melting and boiling points.
 Argon does not form an oxide.
 The oxides of the period three elements also
show a pattern in their acidity/basicity copy out
the table on the next slide and fill in whilst
watching the video.
Activity
 Copy out this table and fill in as you watch the video clip ILPAC
VIDEO 4 and during discussion after.
Element Na
State
at rtp
Nature
of
bonding
Formula
of
oxide
pH
Mg
Al
Si
P
S
Cl
Ar
Element Na
Mg
Al
Si
P
S
Cl
State
at rtp
s
s
s
s
g
g
s
Nature
of
bonding
Formula Na O MgO
2
of
oxide
pH
?
10.5
Simple
covalent
Giant
Covalent
Ionic
Al2O3
SiO2
Ar
Atomic
P4O10 SO2 Cl2O
SO3 Cl2O7
9.1
6.3
1.4
1.8
?
pH 12
A graph to show the pH of period 3 oxides.
12
10
pH
8
6
pH
4
2
0
11
12
13
14
15
16
17
pH 2.5
Atomic number of the elements across period
3
 What do you predict to be the pH of Na2O and
Cl2O7?
Reaction of metal oxides with
water
 Metals on the left of the periodic table typically
have basic oxides. For example;
 Magnesium oxide reacts with water to give
magnesium hydroxide, which is sparingly soluble in
in water and produces a somewhat alkaline solution
of pH about 10:
MgO(s) + H2O(l) → Mg(OH)2(s)
Mg2+(aq) + OH-(aq)
 As we move to the right along the period we reach
aluminium
 Aluminium oxide is virtually insoluble in water. But,
aluminium oxide will react with both acids and
alkalis, it is an amphoteric oxide.
Reaction of non-metal oxides
with water
 Non-metals on the right of the table typically
have acidic oxides. For example;
 Sulphur dioxide reacts with water to give a strongly
acidic solution of sulphuric(IV) oxide acid
(sulphurous acid). This dissociates producing H+
ions, which cause acidity of the solution:
SO2(g) + H2O(l) → H2SO4(aq) →
H+(aq) + HSO3-(aq)
In summary
 The overall pattern is that;
 Metal oxides, on the left of the periodic table,
form alkaline solutions in water.
 Non-metal oxides, on the right of the periodic
table, form acidic solutions in water
 Semi-metals, in the middle of the period show
amphoteric behaviour
Trends across period
three HL
Objectives:
 Explain the physical states (under standard
conditions) and electrical conductivity ( in the
molten state) of the oxides of the elements in
period 3 in terms of their bonding and structure.
Properties of Period 3 oxides
Property
Formula
Na2O
MgO
Al2O3
SiO2
P4O10
P4O6
SO3
SO2
Cl2O7
Cl2O
Tm oC
1275
2852
2027
1610
24
17
-92
Tb oC
-
3600
2980
2230
175
45
80
good
good
good
Very
poor
none
none
none
State at r.t.p
Electrical
conductivity
when molten
Structure
Nature of
oxide
ionic
basic
amphoteric
Giant
covalent
Simple covalent
acidic
Physical properties
 Melting point;
 Sodium oxide, magnesium oxide and aluminium oxide
are all ionic, this accounts for their high melting
and boiling points.
 Silicon dioxide has a diamond like macromolecular
structure with a very high melting point and boiling
point.
 The oxides of sulphur, phosphorus and chlorine are
covalent because the relatively small difference in
electronegativity between the elements and oxygen.
Physical properties
 Electrical conductivity when molten;
 Sodium oxide, magnesium oxide and aluminium oxide
are good conductors of electricity when molten but
do not conduct when solid. This is because they are
ionic, so when molten, the ions are free to move.
 Silicon dioxide is a very poor conductor of
elelctricity.
 The covalent nature of the non-metal oxides mean
they do not conduct electricity when molten as
there are no free ions or electrons to move.
Chemical properties
 Activity:
 Read through page 53 ‘IB course companion’.
 Under suitable sub-headings write the chemical
reactions for;



the metal oxides with water.
the amphoteric nature of aluminium oxide.
the non-metal oxides with water.
Objectives:
 Explain the physical states (under standard
conditions) and electrical conductivity ( in the
molten state) of the chlorides of the elements in
period 3 in terms of their bonding and structure.
 Describe the reactions of chlorine and the
chlorides referred to above with water
Physical Properties of the
metal chlorides
 The physical properties of the chlorides are related
to their structures
 Sodium chloride Na+ Cl- and magnesium chloride Mg2+
(Cl-)2, are ionic so both;
 Conduct electricity when molten
 Have high melting points and boiling points
 Aluminium chloride has a more covalent character,
because of this;
 It is a poor conductor of electricity when molten
 It has a relatively low melting point, in fact it sublimes
Special characteristics of
Aluminium Chloride
 There is evidence that in the solid state, it exists
as aluminium and chloride ions (in an ionic lattice)
but as it melts there is a dramatic change in the
bonding, and the covalent dimer Al2Cl6 is formed
 Molten aluminium chloride is a poor conductor of
electricity.
 In the gaseous state there is an equilibrium
between aluminium chloride monomer and
aluminium dimer
2AlCl3(g)
Al2Cl6(g)
Physical Properties of the
other period 3 chlorides
 The remaining chlorides of period 3 all have
simple molecular structures.
 These molecules are held together by weak forces
of attraction called van der Waals (chapter 4),
which result in low melting points and boiling
points.
Chemical properties of the
Period 3 metal chlorides
 When sodium chloride dissolves into water it gives
a pH of 7, it is the only chloride to be neutral all
the others are acidic.
 Magnesium chloride gives a slight acidic solution,
when it dissolves in water the small, more densely
charged magnesium ion, attracts the water
molecules and causes some of them to dissociate
to form hydrogen ions, this makes the solution
acidic.
 What do you expect to happen when aluminium is
dissolved in water?
Aluminium chloride
 The aluminium ion is even smaller and even more densely
charged, so when the anhydrous aluminium chloride is
added to water, a very exothermic reaction takes place,
and hydrochloric acid is formed:
AlCl3(s) + 3H2O(l)
Al2O3(s) + 6HCl(aq)
 Strictly speaking the aluminium ion becomes hexahydrated
to form;
hexaaquaaluminium (III) ion, [Al(H2O)6]3+
 The water molecules are strongly attracted to the high
charge density on the small aluminum ion nd three of them
successively dissociate to give hydrated aluminium
hydroxide and hydrochloric acid:
[Al(H2O)6]3+
Al(H2O)3(OH)3(s) + 3H+(aq)
Chemical properties of nonmetallic period 3 chlorides
 All the other chlorides react vigorously with
water to produce acidic solutions of hydrochloric
acid, together with the fumes of hydrogen
chloride.
 Write the equations for these reactions, find
them on page 54 IB Course Companion.
 Chlorine itself reacts with water to some extent
to form an acidic solution.
 Write this equation.
A graph to show the trend in pH for the period 3
chlorides
7
6
pH
5
4
Series1
3
2
1
0
11
12
13
14
15
16
Atomic number of period 3 elements
Property
NaCl
MgCl2
AlCl3
Al2Cl6
SiCl4
PCl3
PCl5
S2Cl2
Cl2
Tm oC
801
714
178
sublimes
-70
-112
-80
-101
Tb oC
1413
1412
-
58
76
136
-35
State
solid
liquid
Liquid
(solid)
liquid
gas
Electrical
good
conductivity
when molten
poor
Structure
ionic
Simple covalent
Nature of
solution
neutral
Weakly
acidic
acidic
none
Objectives:
 List the characteristic properties of transition
elements.
 Explain why Sc and Zn are not considered to be
transition elements.
Quick questions
 How many electrons can exist in an orbital?
 How many d orbitals does the 3d level have?
 What is the maximum number of electrons that can be
held if all the d orbitals of the 3d level are filled?
 What is Hund’s rule?
 Put these levels in order of distance from the nucleus
starting with the level closest to the nucleus: 4s, 1s,
3d, 2p,
 Write the electronic configuration of manganese using
the s,p,d,f notation.
d-block elements
The d-block elements are typical metals:
 They are good conductors of heat and electricity.
 They are hard, strong and shiny, and have high melting
points and boiling points, (one notable exception is
mercury which is a liquid at room temperature).
 These physical properties, together with low chemical
reactivity, make transition metals extremely useful,
e.g. iron for vehicle bodies and to reinforce concrete,
copper for water pipes and titanium for engine parts
that need to resist high temperatures.
Definition of a transition
element
 Transition elements are defined as d-block elements with




an incomplete d level of electrons in one or more of their
oxidation states.
Write the electron configuration for the elements
scandium to zinc.
Zn, which can form Zn2+ ions, is not a transition element,
can you think why?
Although the 4s level fills up before the 3d, it is also the
first to lose electrons, so neither Zn or its ion Zn2+ have
incomplete d orbitals.
Write out the electronic configuration for the Zn 2+ ion
Is Scandium a transition
element?
 Scandium is also considered a non-transition
element, its common oxidation state is +3, and the
tripositive ion has the electronic configuration of
argon.
 What is the electronic configuration of argon?
 The electron configuration of scandium metal is
[Ar]4s23d1, so in a sense the metal is a transition
element but its compounds are not.
Other electronic
configurations
 As the atomic number increases, the 3d sub-level
fills up regularly, except for chromium and
copper.
 These have the configurations [Ar] 4s1 3d5 and
[Ar] 4s1 3d10 respectively, because it is more
energetically favourable to half-fill and
completely fill the 3d sub-level rather than spinpair two of the electrons in the 4s orbital.
Characteristic properties
of the transition elements
 Transition elements show characteristic
properties, although they are not unique to
transition elements. These properties include:
 Variable oxidation states
 Coloured compounds
 Complex ion formation
 Good catalysts
Objectives:
 List the characteristic properties of transition
elements.
 Know that the transition elements can exist in
different oxidation states.
 Explain why some complexes of d-block elements are
coloured.
Variable oxidation states
 Group 1 metals lose their outer electron to form
only 1+ ions and Group 2 lose their outer electrons
to form 2+ ions in their compounds.
 A typical transition metal can use its 3d as well as
its 4s electrons in bonding and this means that it
can have a greater variety of oxidation states in
different compounds. This is because the 4s and
3d levels are very close in energy.
 The 4s level electrons are removed before the 3d
level.
Oxidation numbers shown by the elements
of the first d-series in their compounds
Sc
+III
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
+I
+I
+I
+I
+I
+I
+I
+I
+II
+II
+II
+II
+II
+II
+II
+II
+III
+III
+III
+III
+III
+III
+III
+III
+IV
+IV
+IV
+IV
+IV
+IV
+IV
+V
+V
+V
+V
+V
+VI
+VI
+VI
Zn
+II
+VII
The most common oxidation states are shown in red,
though they are not all stable.
Coloured compounds
 Many transition elements have coloured compounds .
 This is a result of spaces in the d-orbitals.
 These orbital's are not of exactly the same energy
except in isolated gaseous atoms, so electrons can move
from one orbital to another of higher energy.
 To do this they must absorb electromagnetic energy of a
frequency in the visible region, e.g. if a substance absorbs
green light, it lets through red and blue and thus appears
purple.
 The wavelength of light that is absorbed depends on
several factors, all of them due to the amount of energy
required to promote a d electron from the lower split
level to the higher split level.
 There are 5 d - orbitals
 In some compounds these are not of equal energy
higher energy
λ
lower energy
 E.g., a compound containing the ion [Ar]3d9 has 9 electrons
 When white light passes through, a wavelength in the visible
region (e.g green) is absorbed and an electron is promoted to
the higher level.
 The rest of the light is reflected and the complimentary
colour seen (red and blue are reflected which makes purple).
Activity
 Follow the instructions of the experiment to
investigate the different oxidation states and
coloured compounds of Vanadium compounds.
vanadate(V), VO2+(aq), is yellow
vanadate(IV), VO2+(aq), is blue
vanadium(III), V3+(aq), is green
vanadium(II), V2+(aq), is mauve
Reducing Vanadium ions
 Collect all the apparatus, including safety
spectacles.
 Measure 5cm3 acidified ammonium vanadate (V)
NH4VO3 (in H2SO4) and pour into a test tube.
 Add a small piece of zinc metal and warm gently.
 Observe the different colours.
 When the solution becomes green add a spatulaful
of zinc powder to speed up the reaction, and
observe the final mauve colour of V2+
Oxidising Vanadium ions
 Carefully filter the remaining mauve solution into
a clean test tube.
 Add acidified potassium manganate (VII)
dropwise.
 Shake after each addition and note the colour
changes






Zinc metal is used as the reducing agent.
What does reduction mean?
Zinc is powerful enough to reduce the vanadate(V) all the way to
vanadium(II). The reaction is sufficiently slow that you can see the
colours of each of the intermediate oxidation states
Acidified Potassium manganate (VII) is used as the oxidising agent.
What does oxidation mean?
Potassium manganate (VII) can oxidise Vandium (II) to vanadate (V).
oxidation
reduction
 Draw arrows to show the direction of oxidation and reduction.
Objectives:
 Define the term ligand.
 Describe and explain the formation of complexes of
d-block elements.
 Define coordination number and be able to use the
coordination number to predict the shape of the
complex
Complex ion formation
 Because of their small size, transition metal ions
attract species that are rich in electrons.
 Such species are called ligands.
 Ligands are neutral molecules or negative ions
that contain a non-bonding pair of electrons.
 These electron pairs can form coordinate covalent
bonds with the metal ion to form a complex.
 Transition metal ions are good Lewis acids
because they can accept a pair of electrons.
 Ligands are good Lewis bases because they can
donate a pair of electrons.
Ligands
 Water is a common ligand, and most (but not all)
transition metal ions exist as hexahydrated
complex ions in aqueous solutions, for example,
hexaaquairon (III) ion [Fe(H20)6]3+
 Other examples of ligands include:
 Ammonia (NH3)
 Chloride ion (Cl-)
 Cyanide ion (CN-)
 Ligands can replace each other depending on the
concentration of the solution and how much
energy is given to the solution.
Coordination number
 The number of lone pairs bonded to the metal ion
is known as the coordination number.
 Compounds have a different shape depending on
the coordination number,
 Coordination number of 6 are octahedral
 Coordination number of 4 are tetrahedral or square
planar
 Coordination number of 2 are usually linear
Objectives:
 Recall what a catalyst is.
 State the difference between heterogeneous and
homogeneous catalysts
 State examples of the catalytic action of transition
elements and their compounds
Catalytic Properties
 Catalysts are substances that increase the rate
of a chemical reaction without themselves being
chemically changed at the end of the reaction.
 Essentially they work by providing an alternative
pathway for the reaction-one with a lower
activation energy.
 They do this by helping the two reacting species
to come into closer contact with each other.
 Catalysts can be; heterogeneous and
homogeneous.
Heterogeneous Catalysts
 Heterogeneous catalysts, is where the catalyst is
in a different phase from the reactants and
products, they may do this by adsorbing reacting
molecules onto the surface of the metal.
Hydrogenation using catalysts
 An example of a heterogeneous catalyst is in the use of
nickel or palladium in hydrogenation.
 Compounds containing carbon-carbon double bonds (C = C)
are said to be unsaturated, when hydrogen is added across
the double bond they become saturated.
 The catalyst works by adsorbing hydrogen and the
unsaturated compound (e,g, ethene) onto its surface and
aligning them so that they are in the correct orientation to
react.
 Transition metals are particularly good at adsorbing
small molecules, so the metals themselves make good
heterogeneous catalysts.
 Other examples include:
 Iron in the Haber process, where ammonia is manufactured from
nitrogen and hydrogen, (write the balanced symbol equation for
this reaction)
 Rhodium, platinum and palladium in
catalytic converters in cars. These
catalysts convert carbon monoxide,
oxides of nitrogen and unburnt
hydrocarbons into less polluting gases
carbon dioxide, nitrogen and water.
Homogeneous catalysts
 Homogeneous catalysts are in the same phase as
the reactants and products.
 In homogeneous catalysis often the two reacting
species bond chemically together, and then leave.
 During the process the oxidation state of the
central element in the catalyst will increase and
then decrease.
 One of the reasons why transition metal
compounds are such good catalysts is that they
have variable oxidation states: that is that they
can be relatively easily oxidized and reduced.
 Examples of compounds of transition metals that are
important catalysts are:
 The use of manganese (IV) oxide to decompose
hydrogen peroxide:
MnO
2H2O2(aq)
2H2O(l) + O2(g)
 Vanadium (V) oxide in the conversion of sulphur dioxide
into sulphur trioxide during the manufacture of
sulphuric acid in the contact process:
V O
2SO2 (g) + O2 (g)
2SO3 (g)
 Iron and cobalt in biological catalysts, e.g. cobalt in
vitamin B12. Vitamin B12 is essential for the production
of red blood cells and the for the correct functioning
of the central nervous system.
4 (s)
2
5 (s)
The structure of vitamin B12
showing the central cobalt atom
 R represents different
groups (e.g. a methyl group,
-CH3) to which the cobalt
atom can bond in different
forms of the vitamin.
 Vitamin B12 is found in foods
of animal origin (fish, meat,
liver, eggs and milk)
 A lack of vitamin B12 causes
pernicious anaemia.