(B) Periodicity
ATOMIC STRUCTURE
Bonding in the first 20 elements
After completing this topic you should be able to :
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Learners should be familiar with the Periodic Table in terms of elements being
arranged in order of increasing atomic number.
The first 20 elements in the Periodic Table are categorised according to
bonding and structure:
- metallic (Li, Be, Na, Mg, Al, K, Ca)
- covalent molecular (H2, N2, O2, F2, Cl2, P4
- S8 and fullerenes (eg C60))
- covalent network (B, C (diamond, graphite), Si)
- monatomic (noble gases)
Learn about these weak intermolecular forces of attraction which exist
between all atoms and molecules (London dispersion forces.
Learn how to explain differences in physical properties such as viscosity,
melting point and boiling point in terms of differences in strength of
intermolecular forces.
The Chemical Bond
Chemical Bond
Intramolecular
(Within)
Intermolecular
(between)
Van der Waals’
INTERMOLECULAR FORCES OF ATTRACTION
This animation describes and explains the key
intermolecular forces of attraction (Van der Waals
forces of attraction) including London dispersion
forces, permanent dipole-permanent dipole
attractions and Hydrogen Bonding.
Types of bonding in elements
• Metallic
• Covalent
• Londons forces
The Metals
Metallic bonding
Groups I, II, III
Metallic elements
Delocalised
electron
+
+
+
+
+
+
+
+
Positive nucleus (core)
Electron shells
The outer shell in metals is not full and so the outer electrons in
metal atoms can move randomly between these partially filled
outer shells. The electrons are delocalised (sometimes called a
‘sea’ or ‘cloud’ of
Electrons) i.e. they are held in common by all the atoms.
Metallic bonding is the strong electrostatic force of attraction,
between the positive charged ions, formed by the loss of the
outer shell electrons of a metal atom and these delocalised
electrons.
The positive metal ions are held together by this
electron “Glue”.
The outer electrons are delocalised and free to move
throughout the lattice.
The greater the number of electrons in the outer
shell the stronger the metallic bond.
So the melting point of Al>Mg>Na
Metallic Character
The strength of the metallic bond depends on
1) The elements tendency to lose electrons (ionise)
2) The packing arrangement of the metal atoms.
3) The size of the atom.
4) The number of valence electrons in the outer most shell.
5) The number of shells.
Physical properties of metals
A. Metals are malleable and ductile
Metal atoms can ‘slip’ past each other because the metallic
bond is not fixed and it acts in all directions.
Physical properties of metals
B. Conduction of electricity and thermal energy.
Solid and liquid metals conduct heat and electricity.
The delocalised electrons are free to move in the
solid lattice. These mobile electrons can act as
charge carriers in the conduction of electricity or
as energy conductors in the conduction of heat.
Physical properties of metals
C. Change of state
In general, metals have high melting and boiling points
because of the strength of the metallic bond.
When a metal is molten the metallic bonds are still
present.
B.p.’s are much higher as you need to break the metallic
bonds throughout the metal lattice.
Metal boiling point trends
Down a group
Boiling point of alkali metals
1600
Boiling point /oC
1400
1200
1000
800
Series1
600
400
200
0
Lithium
Sodium
Potassium
Metal
Boiling point across a period
Boiling point /oC
3000
2500
2000
1500
Series1
1000
500
0
Potassium
Calcium
Metal
Gallium
The strength of metallic
bonding decreases the
forces of attraction get
weaker.
Across a period
The covalent radius
decreases as the positive
core is increasing in charge,
this has the effect of pulling
the outer closer to the
nucleus the forces of
attraction increase.
The Non-metals
Nobel gases
Group 0
Noble gases
He
Noble gases have full outer electron shells
++
They do not need to combine with other atoms.
They are said to be monatomic.
Group 0 are all gases and
exist as individual atoms.
However, the monatomic gases do form weak inter-atomic
bonds at very low temperatures.
Monatomic elements
++
Sometimes the electrons can end up on one
side of the atom, i.e. the electron cloud can wobble
++
This means that one side of the atom is more
negative than the other side.
δ+
δ-
δ-
δ+
These charges are given the symbol δ ‘delta’
A temporary dipole is therefore formed.
δ+
A dipole can induce other atoms
to form dipoles, resulting in
dipole –dipole attraction.
δLondons forces
Monatomic elements
London dispersion forces are very weak attractive forces
Noble gases b.p.’s
180
160
166
140
120
121
100
b.p / K
80
87
60
40
20
4
Helium
Neon
Argon
Krypton
Xeon
27
0 K = -273.15o C 0
B.p.’s increase as the size of the atom increases.
This happens because the London forces increases with
increasing number of electrons. The more electrons the
bigger the dipole the stronger the London forces.
The Non-metals
Covalent molecules
Covalent molecular elements
Most non-metals exist as discrete covalent
molecules held together by covalent bonds.
Discrete molecules have a definite formula with
a definite number of atoms bonded together
Covalent Bond
Chlorine atom
2,8,7
17+
Chlorine molecule Cl2
2,8,8
17+
17+
diatomic
A covalent bond is formed when a pair of electrons
are shared.
The atoms in a covalent bond is the mutual
attraction of two positive nuclei for a shared pair of
electrons.
Group VII
Fluorine F2
Strong covalent
bond
F
Weak London forces
Strong intra-molecular bonding and
weak inter-molecular bonding exist
in this diatomic molecule.
F 2 m.p. -220o C
F
F
F
Chlorine Cl2
Cl
Strong covalent
bond
Weak London forces
Cl
Cl
Cl
Strong intra-molecular bonding
and weak inter-molecular bonding
exist in this diatomic molecule.
Cl 2 m.p. -101oC
Halogens b.p.’s
500
450
457
400
b.p./ K
Fluorine
350
Chlorine
300
332
250
200
238
Bromine
Iodine
150
100
50
85
0
As the size of the halogen atom increases (more electrons),
so does the size of the London forces between the halogen
molecule.
Group VI
Strong double
covalent bond
O
Weak Londons force
O
O
Sulphur S8
Weak Londons
forces
O
Strong intra-molecular bonding and
weak inter-molecular bonding exist
in this diatomic molecule.
O 2 m.p. -218o C
Higher m.p. because there ar
stronger Londons’ forces
between larger molecules
(more electrons).
m.p. 113oC
Group V
Strong triple
covalent bond
N
Weak Londons force
N
N
N
Strong intra-molecular bonding and
weak inter-molecular bonding exist
in this diatomic molecule.
N 2 m.p. -210o C
Phosphorus P4
m.p. 44oC
Strong covalent
bonds
Weak Londons
forces
Group IV
Buckminster fullerene (Bucky Balls) were discovered
in the 1980’s.
C60
C70
Due to the large molecules , fullerenes have stronger
London forces between their molecules, compared to
elements made from smaller molecules.
Fullerenes are a family of carbon molecules made up
of rings with definite formula.
They are discrete covalent molecules
Covalent Network Elements
In the first 20 elements, only Boron, Carbon and
Silicon have covalent network structures.
Carbon Diamond
Diamond forms an infinite 3D
network structure.
Each carbon atom forms 4
covalent bonds to 4 other
carbon atoms.
Very rigid strong structure.
Diamond is one of the hardest
materials known to man.
C sublimes 3642oC
Graphite
Carbon bonded to only 3 other Carbons
So the spare electrons are delocalised
and so free to move. Graphite is a conductor.
London forces between the
layers allows layers to slide over
each other.
Graphite can be used as a lubricant
Silicon
Silicon has the same infinite 3D network
structure as diamond Si mp 1410oC.
Density change across a period
3
2.5
Density
g/cm3
2
1.5
1
0.5
0
Sodium
Magnesium
Aluminium
Silicon
Phosphorus
Sulphur
Chlorine
Argon
Na Mg Al Si P S Cl Ar
Na to Al the atom size decreases leading to greater packing in metal lattice.
Si is a covalent network, tightly packed atoms in covalent lattice.
P and S are covalent molecular solids with quite densely packed
molecules.
Cl and is a covalent molecular gas at room temperature.
Ar and is a monomolecular gas at room temperature.
Bond Strengths
Bond Type
Strength (kJ mol –1)
Metallic
80 to 600
Ionic
100 to 500
Covalent
100 to 500
Hydrogen
40
Dipole-Dipole
30
Londons forces
1 to 20
Bonding in the first 20 elements
Covalent molecular gases
These elements occur as diatomic
Covalent networks
(two atom) molecules with strong
covalent
bonds
Giant network
of atoms
withbetween
strong the atoms
(intramolecular
bonds) and weak
covalent bonds
between the atoms.
Londons forces between the
Very high melting
and boiling
points. bonds).
molecules
(intermolecular
The weak Londons forces mean low
melting and boiling points.
H
Li
Be
B
Na
Mg
Al
K
Ca
Metallic bonding.
Giant network of positively
charged nuclei surrounded
by delocalised electrons.
Delocalised electrons make
these elements good
conductors.
CC
N
Si
P
O
He
F
Cl
S
Ne
Ar
Covalent molecular solids
Polyatomic (many atom) molecules.
Fullerenes C60 C70
Monatomic elements
P4 and S8. These molecules have many
The
noblelarger
gasesLondon
exist as individual
electrons and this
produces
(monatomic)
atoms. There are only weak
forces than the diatomic
molecules.
Londons forces between the atoms.
The stronger London forces (temporary
Very
energy
is needed to break
dipoles) gives these
twolittle
elements
higher
these
forces and so the noble gases
melting and boiling
points.
have very low melting and boiling points.
These two elements are solids at room
temperature.
Diagram shows part of the
covalent
network
of carbon
Uneven
distribution
of the electrons in
atoms
in
diamond.
the electron cloud create temporary
dipoles
(d atom
+ and is
d-)covalently
which result in a weak
Each
carbon
attraction
between
atoms which come
bonded
to 4 other
carbon
close
to
each
other.
These weak
atoms.
attractions are called Londons forces.
d+
dd+
d-
H
He
Li
Be
B
CC
N
O
F
Ne
Na
Mg
Al
Si
P
S
S
Cl
Ar
K
Ca
Metallic bonding with a network of
positively charged nuclei
surrounded by a ‘sea’ of
delocalised electrons.
Diagram
Strong
showscovalent
S8
bonds
molecules
between
in sulphur
the atoms inside
-thelondons
- forces
-molecules.
with the
diatomic
shown
by
the
dotted
+
+
+
+
+
+
+
lines.
+
+
+
+
+
+
+
These large
Weakmolecules
van der Waal’s
have +stronger
+
+ Londons
+
+ molecules
+
+
forces
between
forces than
the
which come close to each
diatomic
molecules.
-other.
- - - -
Bonding in the first twenty elements
This interactive animation provides a visual
representation of the bonding and structure of the
first twenty elements in the periodic table, taking into
account both the intra- and inter-molecular forces
involved.
Questions on elements – bonding and
structure
1. Explain why the covalent network elements have high
melting and boiling points.
2. Explain why the discrete molecular and monatomic
elements have low melting and boiling points.
3. Does diamond conduct electricity? Explain.
4. Does graphite conduct electricity? Explain.
5. How does the hardness of diamond compare with
graphite? Explain.
6. Give a use for both diamond and graphite.
7. Complete the following table:
Questions on elements – bonding and structure
7. Complete the following table:
Type of bonding and
structure
Metallic solids
Properties
……………. of electricity
Covalent network solids
……….. …. melting points
……………. of electricity
exception ……………….
Covalent molecular
solids
………….. melting points
…………… of electricity
Covalent molecular
(diatomic) gases
and monatomic gases
…………… boiling points