Bonding1

Report

Some atoms are very reluctant to combine with other atoms and exist in
the air around us as single atoms. These are the Noble Gases and have
very stable electron arrangements i.e. their outer shells are full.

All other atoms bond together to become electronically more stable
(they want noble gas configuration).
Bonding produces new substances and usually only the valence electrons
are involved.



The physical properties of a substance depend on its structure and type of
bonding present. Bonding determines the type of structure.
CHEMICAL strong bonds
 ionic (or electrovalent)
 covalent
 dative covalent (or co-ordinate)
 metallic

PHYSICAL weak bonds
 van der Waals‘ forces (weakest)
 dipole-dipole interaction
 hydrogen bonds (strongest)

There are 2 major ways in which atoms can bond:
 Ionic ally
 Covalently

In ionic bonding, electrons are completely transferred from one atom to
another. In the process of either losing or gaining negatively charged
electrons, the reacting atoms form ions. The oppositely charged ions are
attracted to each other by electrostatic forces, which are the basis of the
ionic bond.

Covalent bonding occurs when two (or more) elements share electrons.

Ions are electrically charged particles formed when atoms lose or gain
electrons. They have the same electronic structures as noble gases.

Metal atoms form positive ions, while non-metal atoms form negative
ions.
 e.g. Na+, Cl-, Mg2+, O2- etc .
 If negative electrons are lost the excess charge from the protons
produces an overall positive ion. If negative electrons are gained there
is an excess of negative charge, so a negative ion is formed.

The strong electrostatic forces of attraction between oppositely charged
ions are called ionic bonds.
ONE



to form
Example 1: Sodium Chloride (NaCl)
1s22s22p63s1

combines with ONE
1s22s22p63s23p5
1s22s22p6
1s22s22p63s23p6
2:8:1
2:8:7
2:8
2:8:8
An electron is transferred from the 3s orbital of sodium to the 3p orbital of
chlorine; both species end up with the electronic configuration of the
nearest noble gas the resulting ions are held together in a crystal lattice by
electrostatic attraction
What are the ionic equations?
What about the bonding in MgCl2?

Example 2: Magnesium Chloride (MgCl2)
e¯
Cl
Mg
Cl

e¯
Because magnesium have two outer shell electrons, they can combine
with two chlorine atoms by the transfer of one electron to each atom to
form one Mg2+ and two Cl- ions

What are the ionic equations?


Oppositely charged ions held in a regular 3-dimensional lattice by
electrostatic attraction
The Na+ ion is small enough relative to a Cl¯ ion to fit in the spaces so that
both ions occur in every plane.
ClChloride ion
Na+
Sodium ion
Each Na+ is surrounded by 6 Cl¯
(co-ordination number = 6)
Each Cl¯ is surrounded by 6 Na+
(co-ordination number = 6)

Melting point very high
A large amount of energy must be put in to overcome the strong
electrostatic attractions and separate the ions.

Strength Very brittle
Any dislocation leads to the layers moving and similar ions being adjacent.
The repulsion splits the crystal.
-
+
-
+
+
-
+
-
-
+
-
+
-
+
-
+

Electrical
Don’t conduct when solid - ions held strongly in the lattice conduct when
molten or in aqueous solution - the ions become mobile and conduction
takes place.
Cl-
Na+
ClNa+
Cl-
Na+
ClNa+
ClNa+
Cl-
Na+
Cl-
Cl-
Na+
ClNa+
Na+
Na+
ClDISSOLVING AN IONIC COMPOUND
IN WATER BREAKS UP THE
STRUCTURE SO IONS ARE FREE TO
MOVE TO THE ELECTRODES

Covalent bonds are formed by atoms sharing electrons to form molecules.

One single covalent bond is a sharing of 1 pair of electrons, two pairs of
shared electrons between the same two atoms gives a double bond and it
is possible for two atoms to share 3 pairs of electrons and give a triple
bond.

This kind of bond or electronic linkage does act in a particular direction
i.e. along the 'line' between the two nuclei of the atoms bonded together,
this is why molecules have a particular shape. In the case of ionic or
metallic bonding, the electrical attractive forces act in all directions
around the particles involved.

Example 1: Two hydrogen atoms form the molecule of the element
hydrogen (H2)
H

H
H
H
Example 2: One atom of hydrogen combines with one atom of Chlorine to
give you hydrogen chloride (HCl)
Cl
H
Cl
H

Draw the bonding in ammonia (NH3)
H
H
H
N
H
N
H
H

Draw the bonding in methane (CH4)
H
H
H
H
C
H
H
C
H
H

What about oxygen?
O
O
each atom needs two electrons
to complete its outer shell
O
O
each oxygen shares 2 of its
electrons to form a
DOUBLE COVALENT BOND

Atoms share electrons to get the nearest noble gas electronic
configuration

Some don’t achieve an “octet” as they haven’t got enough electrons e.g.
Al in AlCl3

Others share only some - if they share all they will exceed their “octet”
e.g. NH3 and H2O

Atoms of elements in the 3rd period onwards can exceed their “octet” if
they wish as they are not restricted to eight electrons in their “outer
shell” e.g. PCl5 and SF6
For simple covalent molecules (not giant covalent)



Electrical
Don’t conduct electricity as they have no mobile ions or electrons
Solubility
Tend to be more soluble in organic solvents than in water
Boiling point
The forces between molecules (intermolecular forces) are weak and
known as van der Waals forces. Attractions between molecules increases
as the molecules get more electrons.
e.g.
CH4 -161°C
C2H6 - 88°C
C3H8 -42°C
as the intermolecular forces are weak, little energy is required to separate
molecules from each other so boiling points are low

Recap

CHEMICAL strong bonds
 Ionic
 Covalent
Today’s session:
 PHYSICAL weak bonds
 van der Waals‘ forces
 dipole-dipole interactions
 hydrogen bonds

Intermolecular attractions are attractions between one molecule and a
neighbouring molecule. Weak

All intermolecular attractions are known collectively as
van der Waals forces.

There are 2 types:
1. Dispersion forces (London forces) or instantaneous dipole-induced
dipole forces
2. Dipole-dipole interactions

The forces of attraction which hold an individual molecule together (for
example, the covalent bonds) are known as intramolecular attractions.
Strong

In the context of this session, the word dipole means an asymmetric
distribution of electron electrical charge to give partially positive and
partially negative regions in the same molecule.

In a simple sense its a molecule with a partially positive end and a partial
negative charge at the other end.

Electric dipoles may be permanent or transient (temporary) and the
molecules discussed here are electrically neutral overall.

These attractive forces can operate between ANY particles whatever
their constitution including free atoms in a gas or ions in a crystal etc.
Instantaneous dipole-induced dipole forces/Dispersion Forces
 Because electrons move quickly in orbitals, their position is
constantly changing; at any given instant they could be
anywhere in an atom. The possibility will exist that one
side will have more electrons than the other. This will give
rise to a dipole...
 The dipole on one atom induces dipoles on nearby atoms
 Atoms are now attracted to each other by a weak forces
Original temporary
dipole
Induced dipole

There is no reason why this has to be
restricted to a few molecules. As long as
the molecules are close together this
synchronised movement of the electrons
can occur over huge numbers of
molecules.

Although the bonding within molecules is
strong, between molecules it is weak.
(Intramolecular vs. intermolecular)



The boiling points of the noble gases are
 Helium -269°C
 Neon -246°C
 Argon -186°C
 Krypton -152°C
 Xenon -108°C
 Radon -62°C
What is happening to the boiling point as you go down the
group?
Why?

Group 18 (Noble Gases)

The boiling points increase as you go down the group

All of these elements have monatomic molecules.

The number of electrons increases and hence the radius
of the atom.

The more electrons you have, and the more distance
over which they can move, the bigger the possible
temporary dipoles and therefore the bigger the
dispersion forces

Neon molecules will break away from each other at
much lower temperatures than argon molecules - hence
neon has the lower boiling point.
Will molecular shape have an effect on the
strength of dispersion forces?
Use butane and 2-methyl propane to
explain your answer.

Long thin molecules can develop bigger temporary dipoles due to electron
movement than short fat ones containing the same numbers of electrons.

Long thin molecules can also lie closer together - these attractions are at their
most effective if the molecules are really close.

Butane has a higher boiling point because the dispersion forces are greater. The
molecules are longer (and so set up bigger temporary dipoles) and can lie closer
together than the shorter, fatter 2-methylpropane molecules.

The ability of an atom to attract the electron pair in a covalent bond to
itself

Non-polar bond e.g. Cl2, O2
 similar atoms have the same electronegativity they will both pull on
the electrons to the same extent the electrons will be equally shared

Polar bond e.g. HCl
 different atoms have different electronegativities one will pull the
electron pair closer to its end it will be slightly more negative than
average, and the other end slightly more positive. A dipole is formed
and the bond is said to be polar.
 Greater electronegativity difference = greater polarity

The Pauling Scale is a scale for measuring electronegativity values
increase across periods values decrease down groups fluorine has the
highest value
H
2.1
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
K
Br
0.8
2.8
Predict the polarity for the following:
S – Cl ; C – O and C – C

A molecule like HCl has a permanent dipole because chlorine is more
electronegative than hydrogen. These permanent, in-built dipoles will
cause the molecules to attract each other rather more than they
otherwise would if they had to rely only on dispersion forces.


Dipole-dipole interactions occur in addition to Dispersion forces .
Molecules that have permanent dipoles will therefore have higher boiling
points than those that have temporary dipoles.

Use ethane and fluoromethane to explain why a molecule
with a permanent dipole has a higher boiling point.
Ethane
Fluormethane

Both have the same number of electrons so we expect the dispersion
forces to be the same.

Fluromethane has a higher boiling point due to the large permanent
dipole on the molecule (because of the highly electronegative Fluorine).

the attractive force between hydrogen in a polar bond (particularly H-F,
H-O, H-N bond) and an unshared electron pair on a nearby small
electronegative atom or ion.
Hydrogen Bonding & Water

This other atom may be in the same molecule or in a nearby molecule, but
always has to include hydrogen.
INTERMOLECULAR HYDROGEN BONDING
INTRAMOLECULAR HYDROGEN BONDING


Hydrogen Bonds have about 5% of the strength of an average covalent bond
Hydrogen Bond is the strongest of all intermolecular forces
One of the most remarkable consequences of H-bonding is found in the lower
density of ice in comparison to liquid water, so ice floats on water.
The molecules in the solid are more densely packed than in the liquid.
A given mass of ice occupies a greater volume than that of liquid water because of
an ordered open H-bonding arrangement in the solid (ice) in comparison to
continual forming & breaking H-bonds as a liquid.

generally much weaker than covalent or ionic bonds. Less energy is thus
required to vaporize a liquid or melt a solid. Boiling points can be used to
reflect the strengths of intermolecular forces (the higher the Bpt, the
stronger the forces)

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