CfE Higher Chemistry Chemical Changes and Structure Bonding in … · 2016-02-03 · Unit One –...
Transcript of CfE Higher Chemistry Chemical Changes and Structure Bonding in … · 2016-02-03 · Unit One –...
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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.