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A Cook 2015
Chemistry Led Specification Revision
Chapter 2 - Chemical Bonding and Structure
Topic 2A - Bonding
(a) Know that ionic bonding is the strong electrostatic attraction between oppositely charged ions
Ionic Bonding: THE ELECTROSTATIC ATTRACTION BETWEEN OPPOSITE CHARGED IONS
The nature of ionic bonding:
• Ionic bonding is confined to solid materials consisting of a regular array of oppositely charged
ions extending throughout a giant lattice network
• Most familiar ionic compound is sodium chloride, NaCl
...


• In an ionic solid, there are strong electrostatic attractions between ions
...

The strength of ionic bonding:
• The strength of an ionic bond can be calculated by working out the energy required in one mole
of solid to separate the ions to infinity (to the gas phase)
...


Increasing Size of Cation

The table shows the energy required to separate to infinity the ions in one mole of various
alkali metal halides:
Increasing Size of Anion
Amount of energy required to separate the ions to infinity /kJ mol-1
F-

Cl-

Br-

I-

Li+

1031

848

803

759

Na+

918

780

742

705

K+

817

711

679

651

Rb+

783

685

656

628

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(b) Understand the effects that ionic radius and ionic charge have on the strength of ionic bonding
Smaller Ions Create Stronger Electrostatic Interactions:
For ions of the same charge, the smaller the ions the more energy is required to overcome the
electrostatic interactions between the ions and to separate them
...
Hence the
stronger attraction
...

The ionic radius of an Li+ ion is very much similar to that of the Mg2+ however, the energy required
to separate is over double
...

When both cation and anion are doubly charged then the energy required to separate the ions is
even larger
...
the energy required is 3791 kJ/mol-1
...
The chloride molecule
has gained two electrons to become two chloride ions
...
g
...

• For example it matters how many oppositely charged ions are touching it (its co-ordination
number) and also the nature of the ions
...

• So, if you are going to make reliable comparisons using ionic radii, they must come from the
same source
...
Trying to explaining
things in fine detail is made difficult by those uncertainties
...

Group 1
Ion

Electro Config

Ionic Radii

Group 7
Ion

Electro Config

Ionic Radii

Li+

2

0
...
8

0
...
8

0
...
8
...
181

K+

2
...
8

0
...
8
...
8

0
...
8
...
8

0
...
8
...
18
...
220

Why?
As you go down each group the ions have more electron shells; therefore, the ion gets larger

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Ionic radii of isoelectronic ions:
Period 2

N3-

O2-

F-

Period 3

Na+

Mg2+

Al3+

No
...
Protons

7

8

9

Electronic Config

2
...
8

2
...
8

2
...
8

Ionic radius /nm

0
...
140

0
...
147

0
...
133

The ionic radii increases with increasing proton number
Why?
1
...

2
...
As the positive charge of the nucleus increases, the electrons are attracted more strongly and
are therefore pulled closer to the nucleus
...
As a result of this, the energy required to overcome these forces of
attraction on each ion is very large
...

• Lots of oppositely charged ions in the lattice which are electrostatically attracted to one another
...

Brittleness:
• If stress is applied to a crystal of an ionic solid, then the layers of ions may slide over each other
• Ions of same charge are now next to each other and so repel one another; breaking the crystal
...

Why?
This is because there are no delocalised elections and the ions are also not free to move
under the influence of an applied potential difference
2) Molten Ionic Compounds:
However molten ionic compounds will conduct
Why?
Since the ions are now movie and will migrate to the electrodes of opposite charge, when a
potential difference is applied
3) Aqueous ionic compounds:
Aqueous solutions of ionic compounds also conduct electricity and undergo electrolysis
Why?
Since the lattice breaks down into separate ions when the compound dissolves
4) Exceptions:
Lithium Nitride (LiN3) will conduct electricity, and is used in batteries for this reason
Solubility:
Many ionic compounds are soluble in water
...

• Both positive and negative ions are attracted to water molecules because of the polarity that
water molecules possess
...

For example, when a direct electric current is passed through molten sodium chloride, sodium is
formed at the negative electrode (cathode) and chlorine is formed at the positive electrode (anode)
...

Aqueous copper(II) ions, Cu2+(aq), are blue and aqueous chromate(VI) ions, CrO42-(aq) are yellow
...


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(g) know that a covalent bond is the strong electrostatic attraction between two nuclei and the
shared pair of electrons between them
Covalent Bonding
Formation of covalent bonds:
A covalent bond is formed between two atoms when an atomic orbital containing a single electron
from one atoms overlaps with an atomic orbital, also containing a single electron, of another atom
...


Sigma bond formation: Single Bond (stronger)
Occurs when the bond is between the two nucleus
An end-on overlap leads to the formation of sigma bonds
...

Pi bond formation: (weaker)
When the bonds formed and above and below the nucleus
A sideways overlap of two p-orbitals leads to the formation of a pi bond
...
So pi bonds only exist between atoms that
are joined by double or triple bonds
Single Bond = ∂
Double Bond = ∂ + π
Triple Bond = ∂ + 2π

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Why is π stronger than a ∂?
The π bond presents the electron overlap above and below the charged nuclei which means that
there is lesser forces of attraction, whereas the overlap between a ∂ bond occurs between the two
nuclei meaning they experience more of the electronic forces of attraction from the positive nuclei
...
The electronic configuration of hydrogen is 1s1
...


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(h) Be able to draw dot-and-cross diagrams to show electrons in covalent substances, including:
i molecules with single, double and triple bonds
ii species exhibiting dative covalent (co-ordinate) bonding, including Al2Cl6 and ammonium
ion
Dot-and-cross Diagrams:
Covalent and polar covalent bonding in discrete molecules can be shown by dot-and-cross
diagrams
...

Shown by an arrow starting from the atom providing the pair of electrons and going towards the
atom with the empty orbital
...


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The strength of a covalent bond is measured in terms of the amount of energy required to break
one mole of the bond in the gaseous state
Relationship:
The general relationship between bond length and bond strength, for bonds that are of a similar
nature, is the shorter the bond, the greater the bond strength
...


SHORTER BOND LENGTH, GREATER BOND STRENGTH
(j) understand that the shape of a simple molecule or ion is determined by the repulsion between
the electron pairs that surround a central atom
Electron pair repulsion theory:
The electron pair repulsion theory (EPR) states that:

- The shape of a molecule or ion is caused by repulsion between the pairs of electrons, both bond
pairs and lone (non-bonding) pairs, that surround the central atom
...

The shapes of molecules and ions:
To obtain the shape of a mole use or ion it is first necessary to work out how many bond pairs and
lone pairs around the central atom there is
This is best done by drawing the dot and cross diagram
...
5

107
104
...

What is electronegativity?
Electronegativity is the ability of an atom to attract a bonding pair of electrons in a covalent bond
...
This is because the ability of each atom to
attract the bonding pair of electrons is identical
...
This is shown by the following density map of hydrogen chloride (HCl) molecule:

The electronegativity is more lenient towards the chlorine atom as it is more electronegative
...

A bond like this is called a polar covalent bond
...
This separation of charge is called a dipole
...
The overall dipole of a
molecule depends on its shape
...

• If cancellation is complete, the resulting molecule will have no overall dipole and said to be ‘nonpolar’
...
The two atoms in
each molecule are the same and have the same electronegativity
...
The bond in each is therefore non-polar, making the molecules non-polar
...
0) is greater than that of hydrogen (2
...

Since this is the only polar bond in the molecule, overall it is polar
...
These
ideas are:
Pure (100%) covalent
Pure (100%) Ionic
Consider polar covalent bond as a covalent bond that has some degree of ionic character
...

These interactions are described as:
• Non-bonded interactions
• ‘Intermolecular’ because they occur between molecules
The most important non-bonded interaction gives rise to force known as ‘London forces’
...

Other intermolecular interactions arise from the permanent dipoles that exist in some molecules
...
The molecules are
non-polar because they have the same electronegativity and so have symmetrical electron density
throughout them
...
If at any instant the electron density becomes
unsymmetrical in molecule A, a dipole is generated
...
So an
instantaneous dipole has been created in the molecule
...
The result is the creation of a
temporary induced dipole in the neighbouring molecule
...
It is this favourable interaction that is
responsible for the London forces of attraction between two molecules
...

Features of London forces:
• London forces increase their attractive force with the more electrons in the molecules
...
The boiling points increase down the group
...
Proof that
increased number of electrons increases the attractiveness of London forces
...

• London forces depend on the size and shape of the molecule
...

• London forces are always present between molecules, regardless of whether they have
permanent dipoles or whether or not they hydrogen bond with each other
...
If the dipoles
are aligned correctly, then there will be a favourable interaction and the two molecules will attract
one another
...

As a result, when averaged out, the interaction between permanent dipoles is usually much less
than the interaction between instantaneous-induced dipoles
...

It is possible for a molecule with a permanent dipole to induce a dipole in a nearby molecule
...


Summary of non-bonded intermolecular interactions:
Name of interaction

Origin

London Forces

Instantaneous dipole - induced dipole interaction

Permanent Diples

permanent dipole - dipole interaction

van der Waals forces: All of the intermolecular interactions between molecules
...


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The Hydrogen Bond:
The atom bonded to the hydrogen has to be more electronegative than hydrogen, and that there
must be some evidence of bond formation between the hydrogen and another atom, either within
the same molecule (intramolecular hydrogen bonding) or a different molecule (intermolecular
hydrogen bonding)
...

• The interaction is not just that of an extreme dipole-dipole interaction; there is some partial bond
formation utilising a long pair of electrons on the oxygen atom
...

• Hydrogen bonds are directional, so the three atoms involved H—O—H are at 180˚ to each other
...


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Hydrogen bonding through nitrogen:
Example: Ammonia, NH3
All organic containing —N—H group can form intermolecular hydrogen bonds
...


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(r) Understand the following anomalous properties of water resulting from hydrogen bonding:
i its relatively high melting temperature and boiling temperature
ii the density of ice compared to that of water
Anomalous properties of water:
Water has number of anomalous properties
...

1) It has a relatively high melting and boiling temperature
2) The density of ice at 0˚C is less than that of water at 0˚C
Melting and boiling temperatures:
The hydrogen bonds between water molecules are relatively strong
...

Water therefore has an abnormally high melting temperature (0˚C, 273 K) and boiling temperature
(100˚C, 373 K) at 100kPa pressure
...
The differences occur as a result of different
extents of hydrogen bonding
...
this is because of 2 factors:
1) HF forms forms an average of one hydrogen bond per molecule, whereas water molecules
form an average of two hydrogen bonds per molecule so the hydrogen bonding is much more
extensive in water
...



Ammonia falls in between the two
...
It is highly likely, therefore, that all, or nearly all, of
the hydrogen bonds between the molecules are broken on vaporisation, which would explain why
its boiling temperature is higher than that of HF
...
The density of the solid (ice) is less than the density of the
liquid at 0˚C
...
This leads to a lower
density but a larger space filled
...

(s) Understand, in terms of intermolecular forces, physical properties shown by materials,
including:
i the trends in boiling temperatures of alkanes with increasing chain length
ii the effect of branching in the carbon chain on the boiling temperatures of alkanes
iii the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes
with a similar number of electrons
iv the trends in boiling temperatures of the hydrogen halides, HF to HI
Boiling temperatures of alkanes and alcohols
Unbranched alkanes
The alkanes are a homologous series of hydrocarbons with the general formula CnH2n+2
The only significant interaction between alkane molecules is the London force
Boiling point increases with Molecular Mass:
Reasons:
1) As molecular mass increases, the number of electrons per molecule increases and so the
instantaneous and induced dipole increase
2) As the length of the carbon chain increases, the number of points of contact between adjacent
molecules increases
...
There are points of contact all along the chain
...

Branched Alkanes:
The more branching in the molecule, the fewer points of contact between adjacent molecules; i
...

they do not pack together as well
...

Alcohols:
Alcohols are homologous series of compounds with the general formula CnH2n+1OH
They contain an —O—H group and can therefore form intermolecular hydrogen bonds in addition
to London forces
...

Alcohol

Electrons

Boiling
temp /K

Alkane

Electrons

Boiling
temp /K

CH3OH

18

338

CH3CH3

18

184

CH3CH2OH

26

352

Ch3CH2CH3

26

231

CH3CH2CH2OH

34

370

CH3CH2CH2CH3

34

267

CH3CH2CH2CH2OH

42

390

CH3CH2CH2CH2CH3

36

303

Let us consider the case of methanol (CH3OH) and ethane (CH3CH3)
Both molecules have a similar chain length and also have the same number of electrons
...
However, the boiling points of the alcohol is higher than that of the alkane
...
The additional force of attraction increases the energy required
to separate the molecules
...
We have
already mentioned that this is now alway the case
...


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Boiling Temperatures of the hydrogen halides:
The graph shows the boiling temperatures of the hydrogen halides HF to HI

The steady increase in boiling temperature from HCl to HI is the result of ht increasing number of
electrons per molecule, which in turn results in an increase in London forces
...

Dissolving ionic solids:
Many ionic solids dissolve in water
...

The diagram shows this process of dissolving for sodium chloride:

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The slightly negative end of the water molecules attract the sodium ions sufficiently to remove
them from the lattice; The sodium ions then become surrounded by water molecules, as shown
...

The slightly positive end of the water molecules attract the Chlorine ions sufficiently to remove
them from the lattice
...

Dissolving Alcohols:
Compounds that can form hydrogen bonds with water
Alcohols contain —O—H group and can therefore form hydrogen with water
...
The hydrogen bonding between the ethanol and water
molecules is similar in strength to the hydrogen bonding in pure ethanol and in pure water
...

Water is a poor solvent for compounds
Compounds that cannot form hydrogen bonds with water:
Non-polar molecules such as the alkanes do not dissolve in water
...

Many polar molecules also have limited solubility in water
...

Ethoxyethane, CH3CH2OCH2CH3 is polar
...
The
forces of attraction between ethoxyethane and water molecules are not large enough to replace
the relatively strong hydrogen bonding between the water molecules
...
They are much more
soluble in ethanol, and this is why some reactions of halogenoalkanes are carried out in medium of
aqueous ethanol
...

Non-Aqueous solvents:
The ‘rule-of-thumb’ is that; ‘like dissolves like’
...

For example, alkanes are soluble in one another
...

Non-polar bromine dissolves readily in non-polar hexane, and this solution is commonly used to
test for unsaturation in molecules
...

DISSOLVING SOLUTIONS: LIKE FOR LIKE (NON-POLAR FOR NON-POLAR AND POLAR
FOR POLAR)

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(u) know that metallic bonding is the strong electrostatic attraction between metal ions and the
delocalised electrons
Metallic Bonding:
The nature of metallic bonding
Features of metals:
•
•
•
•
•

high melting temperatures
good electrical conductivity
good thermal conductivity
malleability
ductility

Metals typically have one, two or three electrons in their outer shells of their atoms and have low
ionisation energies
...

Since electrical conductivity depends on the presence of mobile carriers of electric charge, we can
build a picture of a metal as consisting of an array of atoms with at least some of their ofter shell’s
electrons removed and free to move about the structure
These free electrons are called DELOCALISED ELECTRONS and are largely responsible for the
properties of metals
...
e
...

There are forces of electrostatic attraction between the nuclei of the cations and the delocalised
electrons
...


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Explaining the physical properties of metals:
Melting temperature:
In order to melt a metal it is necessary to overcome the forces of attraction between the nuclei of
the cations and the delocalised electrons to such an extent that they cations are free to move
around the structure
...
The energy required to do this is usually very large, so the melting temperatures are
usually high
...

• Group 1 metals have low melting temperatures
• Group 2 metals have higher melting temperatures
• Metals in the d-block typically have high melting temperatures since they have more delocalised
electrons per ion
...

SMALLER CATION = CLOSER DELOCALISED ELECTRONS TO CATION NUCLEU AND
GREATER INTERACTION FORCES
The smaller the cation, the closer the delocalised electrons are to the nucleus of the cation
...

Electrical conductivity:
When a potential difference is applied across the ends of a metal, the delocalised electrons will be
attracted to, and move towards, the positive terminal of the cell
...

Thermal Conductivity:
Two factors contribute to the ability of metals to transfer heat energy
...

Ductility:
They can also be drawn into a wire (ductile)
...

When a stress is applied to a metal, the layers of cations may slide over one another
...


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(v) Know that giant lattices are present in:
i ionic solids (giant ionic lattices)
ii covalently bonded solids, such as diamond, graphite and silicon(IV) oxide (giant covalent
lattices)
iii solid metals (giant metallic lattices)
Introduction to solid lattices:
The four types of solid structures we shall deal with in the section are:
1)
2)
3)
4)

giant metallic lattices
giant ionic lattices
Giant covalent lattices
Discrete (simple) molecular lattices

Metallic lattices:
Metallic lattices are composed of a regular arrangement of positive metal ions surrounded by
delocalised electrons
...

Substances that consist of a giant ionic lattice typically have the following properties:
•
•
•
•

A fairly high melting temperature
Brittleness
Poor electrical conductivity when solid but good when molten
Often soluble in water

Giant Covalent lattices:
Giant covalent lattices are sometimes called network covalent lattices
...
All bond angles are 109
...
It
also has a very high melting temperature because a great number of strong C-C bonds have to be
broken in order to melt it
...

Graphite:
Each carbon is bonded to three others by sigma bonds, forming interlocking hexagonal rings
...
The carbon atoms are close enough for the porbitals to overlap with one another to produce a cloud of delocalised electrons above and below
the plane of rings
...
Although thee are
weak London forces of attraction between the layers, this does not account for its lubricating
properties
...
Its lubricating properties are as a result of adsorbed gases on the surface of the
carbon atoms
...

Graphite is a fairly good conductor of electricity
...
An interesting feature is that
graphite can only conduct electricity parallel to its layers
...
Comparing this to a metal which has the ability
to conduct electricity in all directions throughout the structure
...

Graphene:
Graphene is pure carbon in the form of a very thin sheet, one atom thick
...

(y) Know that the structure of covalently bonded substances such as iodine, I2, and ice, H2O, is
simple molecular
Molecule lattices:
Two common solid molecular lattices are iodine and ice
...

Iodine exists as diatomic molecules, I2
...

Physical properties of molecular solids
Molecular solids will, in general, have low melting and boiling temperatures
...

Since intermolecular forces of attraction tend to be weaker than covalent bonds, little energy is
required to break down the lattice structure of a solid and cause it to melt, or to separate the
molecules and cause the liquid to boil and vaporise
...
So, a macromolecule solid such as
poly(ethene) will have a much higher melting temperature than its monomer, ethene
...

Bonding

Structure

Examples

Metallic

Giant Lattice

Mg, Al, Cu, Zn

Ionic

Giant Lattice

NaCl, MgO, CsF

Covalent (including polar
covalent)

Giant Lattice
Molecular
Macromolecular

C(diamond, C(graphite), Si, SiO2, BN
H2O (Ice), I2, P4, S8, C60, C11H22O11 (Sucrose)
Polymers (e
...
poly(ethene)), proteins, DNA

(aa) be able to predict the physical properties of a substance, including melting and boiling
temperature, electrical conductivity and solubility in water, in terms of:
i the types of particle present (atoms, molecules, ions, electrons)
ii the structure of the substance
iii the type of bonding and the presence of intermolecular forces, where relevant
Predicting physical properties:
The physical properties of a substance are determined by the type of bonding and structure it has
The following flow diagram enables you to determine the type of bonding and structure in a
substate by considering its properties
...
g
...
g
...
g
...


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