Trends in the Periodic
Table
Lesson objectives
• Recognise periodicity in the table and be able to account for this in terms of
effective nuclear charge, ionisation energy, bond types etc.
• Recall and be able to explain general trends in melting and boiling points, and
acid-base behaviour
(the above using the elements Na to Ar as exemplification)
• Recognise and be able to account in terms of effective nuclear charge for trends
down groups: bond types, melting and boiling points, reactivity etc.
(the above using (as a minimum) groups 1 and 17 as exemplification)
Success criteria
• Describe and explain the trends in melting point of period 3 elements, group 1
and 7 elements.
• Describe and explain the trends in reactivity of group 1 and 7 elements.
Keywords
•
•
•
•
•
•
•
•
•
Metallic bond
Covalent bond
Metallic lattice
Charge density
Covalent structures
Giant molecular
Simple molecular
Monoatomic
Reactivity
Period 3 melting points trends
Describe all the trends you can see in the graph.
Can you explain this trend in terms of the bonding and
structure of the elements?
Melting points across period 3 elements
• They are based on the bonding and structure
present.
• Metallic bond and metallic structure
• Na, Mg and Al
• Covalent bond and giant molecular structure
• Si
• Covalent bond and simple molecular structure
• P, S, Cl, Ar- No bonding
Structure and Bonding
A summary of the variations in structure and bonding of elements across
Periods 3
Increase in melting point from Na to Al
From sodium to aluminium:
•The positive charge increases and the ionic size decreases
•The number
of are
delocalised
atom size.
increases.
The ions
drawn to electrons
reflect theirper
relative
•This results in a higher charge density(larger charge and smaller size of ions
gives rise to a larger charge density)
•This increases the strength of the metallic bond which in turn increases the
melting and boiling points.
MELTING POINT TREND - NON METALS
relative mass
melting point / K
Covalent macromolecular
P4
124
317
S8
256
392
Cl2
71
172
Ar
40
84
Covalent molecular
Covalent molecular
Covalent molecular
Covalent atomic
• Explanation for covalent structures
• Silicon has a giant covalent structure in which its atoms are
joined by strong covalent bonds.
• There are weak van der Waals forces between the molecules of
phosphorus, sulphur, chlorine and Argon.
• This force is more in sulfur (S8) due to increased molecular
mass followed by phosphorous(P4), chlorine(Cl2) and Argon
which exist as single atoms (Ar)
Trends in melting and boiling points of halogens
Melting and boiling of simple covalent molecules involves breaking
the intermolecular force that hold the molecules together
Melting points and
boiling points
increase down the
group.
Trends in melting and boiling points of halogens
Going down the group, halogen molecules increase in size and the number of
electrons increase.
This leads to greater van der Waals forces between molecules
This increases the energy needed to separate the molecules and therefore
higher melting and boiling points going down the group
van der Waals forces
fluorine
atomic radius = 42 × 10-12 m
boiling point = -118 °C
iodine
atomic radius = 115 × 10-12 m
boiling point = 184 °C
Trends in melting and boiling point of alkali
metals
Melting and boiling of metals involves the breaking the force of attraction
between the metal ions and the delocalized electrons (metallic bond)
Melting points and boiling
points decrease down the
group.
•
•
Going down the group, the
atoms become larger
This makes the attraction
between the delocalised
electron and the nuclei weaker.
Hence less energy is needed
down the group
Reactivity
• To react the atom must form an ion
• For metals
They react by losing electrons
e.g K(g)  K+(g) + e• For non-metals
React by gaining electrons e.g
Cl2(g) + 2e-  2Cl-(g)
Reactivity of halogens
F
 The atoms of each element get larger going down
the group.
 This means that the outer shell gets further
away from the nucleus and is shielded by more
electron shells.
 The further the outer shell is from the positive
attraction of the nucleus, the harder it is to attract
another electron to complete the outer shell.
 This is why the reactivity of the halogens
decreases going down group 7.
Cl
Br
decrease in reactivity
The reactivity of halogens decreases going
down the group. What is the reason for this?
Reactivity of alkali metals
• Metals tend to be have low ionization
energy. This means outer electrons are
easily removed.
Li
Na
• Form ions by losing their outermost
electrons
K
• Reactivity increase down the group
Rb
Cs
Increasing
Reactivity
Reactivity of Alkali Metals
– This is caused by increase in
Li
atomic size and shielding
– Hence the outer electrons are
easier to lose down the group.
– This means that reactivity
increases with the size of the
atom.
– Metallic character also increases
down the group due to the ease
with which electrons are lost
Na
3P
4N
11P
12N
Reactions of period 3 elements with oxygen
Element
Description
Equation
Na
burns vigorously with a yellow flame
to form a white solid
burns vigorously with a bright white
flame to form a white solid
4Na(s) + O2(g) → 2Na2O(s)
Al
burns vigorously with a bright white
flame to form a white solid
4Al(s) + 3O2(g) → 2Al2O3(s)
Si
burns with a bright white flame and
white smoke to form a white solid
Si(s) + O2(g) → SiO2(s)
P
burns spontaneously with a bright white
flame and smoke to form a white solid
burns with a blue flame to form a
colourless gas
P4(s) + 5O2(g) → P4O10( excess)
Mg
S
2Mg(s) + O2(g) → 2MgO(s)
S(s) + O2(g) → SO2(g)
Other oxides
P: if oxygen is limited, phosphorus(III) oxide can also form:
P4(s) + 3O2(g) → P4O6(s).
S: some sulfur trioxide can also form: 2S(s) + 3O2(g) → 2SO3(g)
Chlorine doesn’t react directly with oxygen but has its 3
oxides i.e Cl20, Cl2O5 and Cl2O7
Argon is unreactive.
Bonding of period 3 oxides
• Period 3 elements react with oxygen to form oxides
• Some have more than one oxides e.g P,S and Cl
• Atoms of lower electronegativity lose electrons to the
oxygen atoms
• Atoms of higher electronegativity share electrons with
the oxygen atoms
• Because of this, bonding of the oxides of elements
across period 3 becomes less ionic and more covalent
Bonding and structure in period 3 oxides
•
•
•
•
Ionic bond and giant ionic structures
Na2OMgO
Al2O3 – It has some degree of covalent character
•
•
•
•
•
Covalent bond
SiO2 - giant covalent structure
P4O10 – simple covalent
SO2/SO3 –simple covalent
Cl2O/Cl2O5/Cl2O7 -simple covalent
Acid-base properties of the oxides of
period 3
• Due to the variation in nature of bonding of the oxides, they
show different behavior towards water, acids and bases
•
•
•
•
Behaviour towards acids and bases
Basic oxides
Amphoteric- reacts with both acids and bases
Acidic oxides
Basic oxides
• Na2O and MgO are ionic
• Na2O is highly soluble while MgO is slightly soluble in water
• They react with water to form hydroxides
• Na2O(s) + H2O(l)  2NaOH(aq) (vigorous)
pH 12-14
(strongly alkaline)
•
•
•
•
• MgO(s) + H2O(l)  Mg(OH)2(aq)
pH 9-10
(weakly alkaline)
Moreover, they react with acids to form salt and water
Na2O(s) + 2HCl(l)  2NaCl(aq)+ H2O(l)
Thus they show basic properties and are basic oxides
Amphoteric oxides
•
•
•
•
Al2O3 is ionic oxide with covalent character
It is insoluble in water
It reacts with both acids and bases- amphoteric
Reactions with acid
• Reaction with base
Acidic oxides
• Silicon dioxide, SiO2, phosphorous pentoxide, P4O10, sulfur dioxide, SO2, and
dichloride monoxide, Cl2O are covalent
• All, except SiO2 react with water to form acids. SiO2 is a giant covalent
molecule
• P4O10(s) + 6H2O(l)  4H3PO4 (aq)
• SO2(g) + H2O(l)  H2SO3(aq)
• Furthermore, they react with dilute alkali to form salt and water e.g
• Hence, they show acidic properties and are acidic oxides
Summary
Oxides Na2O
MgO
Al2O3
SiO2
P4O10
SO2/SO
3
Bonding
ionic
ionic
structure Giant
ionic
Giant
ionic
Ionic with
covalent
character
Giant
Ionic
Covalent Covalent Covalent
Aci/base Basic
property
Basic
amphoteric Acidic
Simple
covalent
Simple
covalent
Simple
covalent
Acidic
Acidic