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CfE Higher Chemistry

Unit One – Chemical Changes and Structure

Chapter Four – Bonding in Compounds

Types Of Bonding In Compounds

Bonding

Metallic Ionic Covalent

(Elements) (Compounds) (Elements and Compounds)

Covalent Molecular Covalent Network

Intramolecular

Intermolecular

Pure Covalent Polar Covalent

(Permanent Dipoles)

Dipole-Dipole London Dispersal Forces

Hydrogen Bonding

(Special Version of Dipole-Dipole)

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Types of Bonding in Compounds

Ionic Bonding

Ionic bonding is an electrostatic force of attraction between positive ions of metals

(and some notable non-metal complex ions) with the negative ions of non-metals ( or

groups of non-metals known as complex ions, for example SO42-).

For example Sodium Chloride

Different elements have different degrees of attraction for bonding electrons, i.e.

their electronegativity. The difference in the electronegativities between the metals

and non-metal atoms results in a transfer of electrons from the metal atom (low

electronegativity) to the non-metal atom(s) (high electronegativity),so creating a

positive metal ion and a negative non-metal ion.

Ionic bonding occurs as a result of the electrostatic force of attraction between the

positive metals ions and the negative non-metal ion.

Greater the difference in electronegativities, between the metal atoms and the non-

metal atoms, the greater the degree of ionic bonding.

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Covalent Bonding

Covalent bonding occurs in general between non-metal atoms (there are some notable

exceptions, for example, titanium chloride).

There are two types of covalent bonds;

Pure covalent

Polar covalent.

Pure Covalent Bonds

These bonds result from the non-metal elements in the molecule having the same

electronegativities.

All molecules containing only pure covalent bonds have properties associated with

London Dispersal Forces.

All molecular elements and a number of molecular compounds contain pure covalent

bonds and are therefore gases, liquids or low melting solids at room temperature.

Polar Covalent Bonds

These bonds result from the non-metal atoms having different electronegativities. As a

result of these polar bonds the molecule itself is polar (with the exception of

symmetrical molecules).

Important to note; not all molecules containing polar bonds are polar molecules a group

of molecules that have polar bonds but are symmetrical in shape and therefore non

polar molecules as their polar bonds cancel each other out.

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Compounds with a greater degree of differences in their electronegativities are

considered as most ionic, while compounds with the least degree of difference in their

electronegativities are considered most covalent.

In general differences in electronegativities indicate whether a compound is ionic, polar

covalent or indeed pure covalent but it is not precise but the ‘Bonding Continuum’ can be

used to help understand these subtle differences in the types of bonding associated

with different compounds.

To determine the true nature of the bonding present we must consider the compounds

properties.

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Intermolecular forces of Attraction

There are three intermolecular forces of attraction (two resulting from polar bonds).

1. London Dispersal Forces

The electrons of an atom that surround the nucleus are not stationary but are

constantly moving around the nucleus. This movement of electrons creates an uneven

distribution of electrons (negatively charged) around the (positively charged) nucleus

that results in the formation of a temporary dipole on the atom.

The Greek letter delta (δ) is used to denote a slight amount. Thus delta negative (δ-)

means slightly negative and delta positive (δ+) means slightly positive.

The London Dispersal Forces result from the temporary attraction between the

positive end of one dipole for the negative end of another induced dipole.

This is a temporary dipole-dipole electrostatic force of attraction.

London Dispersal Forces

Dipole – dipole attraction

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2. Permanent Dipole-Dipole Forces of Attraction

These forces of attraction result from the non-metal atoms in the molecule having

different electronegativities.

For example Hydrogen sulfide H2S

Permanent dipole

2.5Sδ-

2.2Hδ+ Hδ+ ιιιιιιιιιιιιιι S δ-

Hδ+ Hδ+

Permanent Dipole-Dipole force of attraction

(Much stronger than London Dispersal Forces of Attraction)

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3. Hydrogen Bonding (A special type of Permanent Dipole-Dipole Forces of

Attraction)

This is a special form of Dipole-Dipole force of attraction in which the non-metal atom

bonded to the hydrogen atom has a much greater electronegativity than hydrogen.

For example water H2O

Permanent dipole

Oδ-

Oδ- ιιιιιιιιιιι δ+H Hδ+

Hδ+ Hδ+

Hydrogen Bonding

(Permanent Dipole-Dipole force of attraction)

(Stronger than normal Dipole-Dipole Forces of Attraction)

Note

Hydrogen Bonding occurs when hydrogen is bonded to;

Oxygen O – H

Nitrogen N – H

Flourine F – H

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Covalent Molecular Structures

Most covalent compounds are molecular.

For Example

Chlorine Hydrogen Flourine

Cl – Cl H – H F – F

Mpt -101˚C Mpt -259˚C Mpt -220˚C

Bpt -35˚C Bpt -253˚C Bpt -188˚C

These molecular elements are gases at room temperature as they only have weak

intermolecular forces of attraction known as London Dispersal Forces between their

molecules.

London Dispersal Forces

F – F ιιιιιιιιιι F - F

Pure Covalent Bonds

( Note; Pure covalent bonds as both atoms in the molecule have identical

electronegativities).

During melting and evaporation of these elements only the very weak intermolecular

London Dispersal Bonds are broken, hence the low melting and boiling point. The

intramolecular pure covalent bonds remain intact.

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Polar Covalent Bonds

In polar covalent bonds the bonding electrons are not shared equally as the bonding

non-metal atoms have different electronegativities.

The atom with the higher electronegativity has a greater share of the bonded

electrons, so has a slight negative charge, (δ-).

The other bonded atom with the lower electronegativity has a slight positive charge,

(δ+).

This results in a polar (dipole) bond.

Examples

Hydrogen chloride (H-Cl)

δ+ δ-

H ---- Cl This is a polar molecule

2.2 3.0

Polar Covalent Bond

Water (H2O)

3.5

O δ-

δ+H H δ+ This is a polar molecule

2.2 2.2

These polar covalent bonds are also known as permanent dipoles.

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Polar Molecules

A polar covalent molecule is one that has permanent charged ends (δ+ and δ-) called

permanent dipoles.

These dipoles enable polar molecules to bond to one another by bonds known as dipole –

dipole forces of attraction.

Consider Hydrogen Iodide, HI

This intermolecular force of attraction is known as a

dipole-dipole force of attraction (bond).

δ+ δ- δ+ δ-

H ---- I ιιιιιιιιιι H ---- I

2.2 2.6 2.2 2.6

Polar Covalent Bonds

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Symmetrical Polar Molecules

Some covalent molecules contain polar bonds but are not overall polar as they are

symmetrical.

For example

Caron dioxide Carbontetrachloride

δ -

δ - Cl

O C O

δ + C δ +

δ - Cl Cl δ -

Cl δ - Polar covalent bonds

As a result of their symmetrical shape, their polarities cancel one another out resulting

in the molecules themselves being non-polar.

Their physical properties result from London Dispersal (intermolecular) forces of

attraction between their molecules, (hence both these molecules are gases at room

temperature).

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Note

The tetrahedral shape is symmetrical

Methane CH4

Subject to London Dispersal forces

H C δ - Polar bonds

Symmetrical molecule

H H δ + Therefore a non-polar molecule

H

Contains polar bonds, but as the central atom carbon is surrounded bonded to four

identical hydrogen atoms with identical electronegativities then this molecule is non-

polar. Methane molecules are subject to weak London Dispersal forces and are

therefore a gas at room temperature.

Simple alkanes like pentane, C5H12 that have polar bonds (δ –C H δ +) are non-polar

molecules as they are symmetrical.

H H H H H

H C C C C C H H H H H H

The only force of attraction associated with the pentane molecules are the

intermolecular London Dispersal forces of attraction. Due to the molecular mass

of pentane it is a liquid at room temperature.

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Non-symmetrical polar molecules Dipole – dipole forces of attraction

For example Ammonia NH3

N δ - Polar bonds (permanent dipole)

Non-symmetrical molecule

H H Therefore a polar molecule

H δ +

N δ - Dipole – dipole force of attraction

H H δ +

H

Chloroform CCl3H

Subject to dipole-dipole attraction

H C δ + Polar bonds

Non-symmetrical molecule

Cl Cl δ - Therefore a polar molecule

Cl Boiling point 61˚C, gfm 119.5g

This difference between methane (a gas) and chloroform (a liquid) is an indication of

the greater strength of dipole-dipole attractions compared with London Dispersal

Forces of attraction.

Polar molecules like chloroform CCl3H have polar bonds (i.e. permanent dipole) within

their structure and are themselves non-symmetrical in shape.

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-

-

=

+

+

-

Non-symmetrical polar molecules Hydrogen bonding

Water, H2O

Boling point 100˚C

Gfm 18g

Liquid at room temperature

H H

O Polar Bonds

Non-symmetrical molecule

Polar molecule

H H

O

This difference between chloroform (Bpt 61˚C) and water (Bpt 100˚C, gfm 18g) is an

indication of the much greater strength of dipole-dipole attraction that is hydrogen

bonding compared with normal dipole – dipole attractions.

Hydrogen bonding is a special form of permanent dipole to permanent dipole interaction

when these molecules contain a hydrogen atom bonded to a highly electronegative

element such as fluorine, oxygen and nitrogen.

2.2 δ +H – F δ - 4.0

2.2 δ + H – O δ - 3.5

2.2 δ +H – N δ - 3.0

Dipole-dipole

intermolecular force

of attraction known

as Hydrogen Bonding

Permanent dipole

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Consider

Ethanol Ether

C2H5OH C2H6O

Gfm = 46g gfm = 46g

Bpt = 79˚C Bpt = -23˚C

H H H H

H – C – C – O – H H – C – O – C - H

H H H H

Polar bonds and non-symmetrical Polar bonds and symmetrical

Therefore a polar molecule Therefore a non-polar molecule

Dipole-dipole attraction (London Dispersal Forces of Attraction)

(Hydrogen bonding)

When hydrogen bonding is present, then the molecule has much higher boiling points

compared with other molecules with similar molecular masses and no hydrogen bonding

All alcohols (-OH functional group) are subject to hydrogen bonding.

Electrical Fields

Polar molecules like ethanol and water bend in an electric field whereas non-polar

molecules like pentane do not.

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Anomalous Physical Properties of some Hydrides

Group four Hydrides

CH4, SiH4, GeH4 and SnH4

These show the expected increase in boiling points with molecular size, due to

increased London Dispersal forces of attraction.

Groups Five, Six and Seven

The boiling points for ammonia (NH3 group 5), Water (H2O group 6) and Hydrogen

fluoride (HF group 7) all have boiling points greater than expected.

The elevated boiling points indicate a stronger intermolecular forces of attraction than

expected from London Dispersal Forces or permanent dipole – Permanent Dipole forces

of attraction. This stronger intermolecular force of attraction is Hydrogen Bonding.

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Viscosity

Viscosity normally increases with increasing molecular size but molecules with hydrogen

bonding show higher viscosity than expected.

Miscibility

Miscible liquids mix thoroughly without any visible boundary eg. water and ethanol.

Immiscible liquids have a boundary between them eg. water and hexane.

Hydrogen bonding helps miscibility eg. both water and ethanol contain hydrogen

bonding. Other polar liquids are also often miscible with water eg. propanone and water

are miscible.

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Solubility

Ionic lattices and polar covalent compounds tend to be soluble in water and other polar

solvents due to the interaction of opposite charges.

When ionic compounds dissolve in water their lattice become surrounded by polar water

molecules.

The negative ions are attracted to the positive ends of the water molecules and the

positive ions are attracted to the negative ends of the water molecule

Ions surrounded by a layer of water molecules, held by electrostatic forces of

attraction are said to be hydrated.

Non-polar molecules will tend to be soluble in non-polar solvents like hexane or carbon

tetrachloride and insoluble in water and other polar solvents as they have no charged

ends to be electrostatically attracted to the polar solvent molecules.

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The Structure of Ice

Normally solids are denser than the own liquids, but, ice floats on water.

The intermolecular bonding associated with small covalent molecules is usually London

Dispersal Forces of Attraction. In ice, the intermolecular force of attraction is

hydrogen bonding. This result in a crystal lattice of water molecules that are held

together by an network of hydrogen bonds.

This arrangement not only makes the structure strong but it also spaces out the water

molecules and so prevents them from packing closely together.

If you examine the

diagram closely you will

see that each water

molecule is surrounded

by four hydrogen bonds

The structure of ice results in

the water molecules being less

densely packed together than

those of water and therefore

the ice floats

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Covalent Network structures

These structures have an infinite three-dimensional network structure of non-metal

atoms bonded together by covalent bonds. These elements have extremely high melting

and boiling points.

Two examples are;

Silicon Dioxide (sand)

Silicon Carbide

Silicon carbide has a similar network structure. It is the hardest known substance to

man. It is used on abrasive wheels for cutting rocks or on grinders for sharpening

metals

The resulting structure is almost as hard as diamond.