Transcript Document

West Midlands Chemistry
Teachers Centre
November 2009
Structure and Bonding
Presenter: Dr Janice Perkins
Types of Bonding
1. Ionic Bonding
2. Covalent Bonding
3. Metallic Bonding
Ionic Bonding
• Negative Ions (anions)
have a negative charge because of a surplus of e-
• Positive Ions (cations)
have a positive charge because of a deficiency of e-
• Ions formed by e- transfer from one atom to
another
the number of e- lost or gained depends on the elements
involved
Ionic Bonding is the attraction between these
positive and negative ions
It is called ELECTROSTATIC ATTRACTION
Question from past paper
Magnesium and chlorine react together to form the ionic
compound magnesium chloride, MgCl2
(i) Explain how each of the ions in this compound is
formed
(ii) Explain why compounds with ionic bonding tend to
have high melting points
Two e- transferred from Mg to Cl 
one e- to each of two Cl atoms 
electrostatic attractions 
are strong 
Too easy?
No, it was very poorly answered
Covalent Bond
• Covalent bond = shared pair of e• One electron comes from each atom
Question from June 09
Two organic compounds with similar relative
molecular masses are shown below.
H
H
H
C
C
O
H
H
H
H
C
C
C
H
H
H
H
H
H
H
Ethanol
Propane
State the type of bond present between the
C and H atoms in both of these molecules.
Explain how this type of bond is formed.
Covalent
Shared pair of electrons, one from each atom 
Co-ordinate Bonding
also called Dative covalent bonding
Co-ordinate bonding is to do with how a
covalent bond is formed
Once formed it is a normal covalent bond
One atom/ion supplies both e- to another
atom/ion
Question from Jan 09
Phosphorus is in the same group of the Periodic
Table as nitrogen.
The molecule PH3 reacts with an H+ ion to form a
PH4+ ion.
Name of the type of bond formed when PH3 reacts
with H+ and explain how this bond is formed.
•Coordinate/dative 
•Both electrons/ lone pair (on P/PH3) 
•Shares/donated from P(H3)/ to H(+) 
Metallic bonding
• This type of bonding is often very poorly
explained.
• frequently see:
ionic bonds
hydrogen bonds
van der Waals’ forces
in the answer of even the better candidates
• This is wrong!
Metallic bonding
X
+
+
+
+
+
+
+
+
+
+
+
+
• OUTER electrons are
‘DELOCALISED’ i.e. free
to move through metal
X
• Metal lattice held together by the
ATTRACTION between this sea of delocalised
e- and the positive ions in the lattice
Question from June 09
State the type of bonding involved in
silver.
Draw a diagram to show how the
particles are arranged in a silver lattice
and show the charges on the particles.
• metallic bonding 
• regular arrangement of same sized particles 
• + charge in each ion 
Bond Polarity
Cl
x
Cl
• A bond is formed between atoms of the same
element
• They both attract the electrons in the bond to
the same extent
• This attraction is called electronegativity
Electronegativity
Definition:
• The power of an atom to attract electron density (or a
pair of electrons) 
• from a covalent bond. 
Trends:
Electronegativity increases across the period
• Due to more protons and smaller size, so stronger
attraction for the bonding eDecreases down a group
• Due to larger size/more shells, so weaker attraction for
bonding e-
• What’s the sporting Caption?
• ‘Throwing the puppy dog!’??
What’s the sporting Caption?
Wish we were on holiday
What’s the sporting Caption?
Well this one is clear enough
It must be ‘Tug-o-War’
Electronegativity and covalent bonds
• Like a tug-o-war
• The ‘tug’ is provided by the electronegativity of
the atom
• A tug-o-war between equal teams is like the pull
of the Cl atoms in the covalent bond because
atoms of the same element have same
electronegativity.
• perfectly even sharing of the bonding e-
Cl
x
Cl
Covalent bonds between atoms
with different electronegativities
• The attraction for the bonding e- by the two atoms
is different
• This results in unequal sharing of the bonding ebetween the two atoms
• The extent of the unequal distribution will depend
on how different the electronegativities are
- the term we use is ‘electronegativity
difference’
• The bigger the difference, the greater the
dipole
Polar Bonds
Atom B is more electronegative than atom A
Bonding e- are attracted more strongly to B
Electron
density
is low.
+
A
+A –
x
B
B
AB is a polar molecule
it has a dipole
Electron
density
is high.
Forces acting between molecules
Types of Intermolecular Forces (IMF)
• van der Waals’ forces
(temporary induced dipole - dipole
attractions)
• Permanent dipole - dipole attractions
• Hydrogen bonding
All result from attractions between the
partial charges in the dipoles
The stronger the dipole – the stronger the
IMF
Van der Waals’ forces
• It is the weakest of the three IMFs
• It result from temporary unequal distributions
of e- density
• The bigger the atom or molecule, the more
electrons there will be and the larger will be
the surface area – this increases the strength
of the van der Waals’ attractions
• It is always present but is often swamped by
stronger IMFs
• It is the only IMF present between non-polar
molecules
Permanent dipole-dipole forces
• Always refer to them as ‘dipole-dipole’, not
just ‘dipole’, attractions
• They result from attractions between the
+ and - of polar molecules
• The greater the electronegativity difference
between the two atoms in the bond, the
greater the dipole
Hydrogen Bonding
• This is the strongest IMF. It is an extreme
example of dipole-dipole attractions
• It only occurs in molecules with very large
electronegativity differences.
• N–H, O–H and F–H bonds
• It results in an attraction that is about 10% as
strong as a covalent bond
Hydrogen bonding in HF
Mark 1 = 3 lone pairs on the fluorine atoms 
Mark 2 = dipole correctly shown 
Mark 3 = hydrogen bond between lp and H 
δ+
H F
δ-
δ+
H F
δ-
Question from Jan 09
Electronegativity
H
C
N
O
2.1
2.5
3.0
3.5
State the strongest type of intermolecular force in the
following compounds.
• Methane (CH4) van der Waals
• Ammonia (NH3) Hydrogen bonding
Use the values in the table to explain how the strongest type
of intermolecular force arises between two molecules of
ammonia
Large electronegativity difference between N + H 
Forms N - / H + 
Lone pair on N attracts H (+) 
States of Matter
• The properties shown by materials are the
result of their structure and bonding
• You need to know about crystalline
materials which are:
Ionic
Metallic
Giant covalent
Molecular
Ionic Crystal Lattices
• Each ion is surrounded by a number of ions of
the opposite charge
• In sodium chloride, each ion is surrounded by
6 of the oppositely charged ions.
• This gives a cubic shape.
Remember.
Oppositely charged ions attract each other by
Electrostatic attraction
2D diagram of NaCl
Na
+
-
Na
+
-
Na
+
-
Cl
-
Na
+
-
Cl
Na
Cl
+
Cl
Na
-
Cl
+
Cl
• Each Na+ ion is surrounded by a total of 6 Cl- ions
One more is be in front of the ion; the other is behind it
3D diagram of NaCl
-
+
Cl
Na+
Na
Cl-
+
Na
Cl-
Cl
Na+
ClCl -
-
ClCl
Each Na+ ion is
surrounded by six
Cl- ions and vice
versa
Why are ionic solids brittle
animation?
+
+
+
+
+
+
+-+
---+
+
+
+
+
+
+-+
-
+
-
+
+
-
+
-
----
Ionic Lattice
OK – so what was happening
there?
Here it is again with sub-captions
for the hard-of-thinking!
Why are ionic solids brittle
animation?
+
+
+
+
+
+
+-+
---+
+
+
+
+
+
+-+
-
+
-
+
+
-
+
-
---Like charges
repel
Metals
In Mg there is an attraction between the
delocalised e and the lattice of metal ions
2+
2+
2+
2+
2+
2+
2+
2+
2+
2+
2+
2+
Most metals have high m.p. / b.p. as a lot
of energy is needed to remove an atom from
this attraction
The delocalised e- flow thorough the metal
(a current) so metals conduct electricity
So how can metals change their
shape?
+
++
++
++
++
++
++
++
++
++
++
++
+
Malleable since
the layers can
slide over one
another
Comparison of Na, Mg and Al
delocalised e-
+
2+
3+
2+
+
3+
+
2+
3+
+
2+
3+
2+
+
3+
+
2+
3+
+
2+
3+
2+
+
3+
+
2+
3+
+
2+
3+
2+
+
3+
+
2+
3+
Na+
Na+
+
Na
Mg2+
Al3+
Mg
2+
3+
Al
Mg
2+
Al3+
• All ions have the same electron configuration (isoelectronic)
• As the proton number increases from Na to Al, so does the
attraction for the outer shell electrons. Size decreases
• Higher ionic charge, more delocalised e-, smaller ionic radius
• Metallic bonding increases; so m.p. / b.p. increase Na  Al
Giant molecular crystals
Diamond
Graphite
Silica
• Physical strength
explain in terms of bond breaking (diamond/Si) and layers
sliding (graphite)
Many strong covalent bonds need to be broken in the
structure of diamond and silica so strong substances.
In graphite the layers can slide over one another so graphite
is softer and is often used as a lubricant.
• Melting points
explain in terms of bond breaking
Many strong covalent bonds need to be broken in
macromolecules so high melting points.
• Electrical conductivity
explain in terms of presence/absence of delocalised eGraphite has delocalised electrons between the layers so it
conducts electricity.
Molecular crystals
• The example you need to know is Iodine
(but you could be asked to apply your understanding
to other elements, such as sulphur)
•
•
•
•
•
•
There is only one element present
The bond is non-polar
The only IMF present is van Der Waals’
These IMFs are weak so
The melting point is relatively low
The IMFs operate between molecules
Shapes of molecules/ions
•
The basic shape is determined by
the number of electron pairs present
•
If all the electron pairs are bonding-pairs
then we get a ‘regular’ shape
•
If some of the electron pairs are lone-pairs
then we get an ‘irregular’ shape
The Electron Pair Repulsion Theory
• The shape results from repulsions between
e- pairs – NOT bonds/atoms
• The rules are:
• Lone pairs repel more strongly than
bonding pairs
• Order is
l.p/l.p. > l.p/b.p >b.p/b.p
• Remember: When lone pairs are present, the
basic shape will be distorted.
Linear Shape
B
A
B
Equal repulsion between
2 bonding pairs
Trigonal Planar shape
B
B
A
B
Equal repulsion between
3 bonding pairs
Tetrahedral shape
B
A
B
B
B
Equal repulsion between
4 bonding pairs
Trigonal Bipyramid shape
B
B
B
A
B
B
Equal repulsion between
5 bonding pairs
Octahedral shape
B
B
B
A
B
B
B
Equal repulsion between
6 bonding pairs
Regular Shapes
epairs
Formula
Bond angle
Name
2
B–A–B
180
linear
3
AB3
120
trigonal planar
4
AB4
109½
tetrahedral
5
AB5
120
90
trigonal
bipyramid
6
AB6
90
octahedral
Working out the shape
• You can work out the shape using simple maths
• Find the group number of the central atom
• Add to this number the number of bonds present
• If it is an ion
– add one if the charge on the ion is -1
– Deduct 1 if the charge is +1
• Divide total by 2 - this gives the number of e- pairs
• Number of e- pairs = basic shape
• If lone pairs present, basic shape modified
To draw the shape
•
•
•
•
•
Work out the basic shape
Lightly sketch this in pencil
Put the atoms in
Draw the lone pairs along the ‘spare’ bonds
Rub out you working pencil lines



H
N
H
Working out the shape of NH3
–
–
–
–
–
N is in Group 5, so five outer electrons
3 more electrons (1 from each H)
Makes 8 electrons
4 pairs of electrons
3 bonding pairs and 1 lone pair
Basic shape (4 e- pairs) = tetrahedral
– One lone pair, so doesn’t have basic
shape
– 3 bonding pairs + 1 lone pair
Ammonia

H
H
N
H
Unequal repulsion between
3 bonding pairs and 1 lone pair
Pyramidal shape
Working out the shape of the ion NH2N is in Group 5, so five outer electrons
– 2 bonds to H atoms, so 2 extra bonding e– 1 charge on molecule, so add extra e-
–
= 8 electrons
– Makes 2 bonding pairs and 2 lone pairs
Basic shape (4 e- pairs) = tetrahedral
– 2 bonding pairs + 2 lone pair
– so doesn’t have basic shape
The shape of the ion NH2


H
N
H
Unequal repulsion between
2 lone pair and 2 bonding pairs
‘Bent’ or ‘V-shaped’
Shapes of molecules / ions
Species
Group
number
Bonds
present
Charge
e- from central
atom
e- from other
atoms
+ charge
reduces e-'
AlCl 3
CCl 4
NH2-
NH3
H3O+
3
4
5
5
6
3
4
2
3
3
0
0
1
0
-1
-
Sum of above
6
8
8
8
8
e pairs
Divide above
by 2
3
4
4
4
4
0
2
1
1
Tetra
Tetr
Tetr
Tetr
Tetra
Bent
Total e
-
Lone
= e- pairs
0
- bonds'
pairs
Shape Closest shape Trig
based on without l.p. Planar
Actual
shape
shape with any
Trig
l.p. present
Planar
Pyramidal Pyramidal
Question from Jan 09
Arsenic is in the same group as nitrogen. It forms a
compound AsH3
Draw the shape of an AsH3 molecule and name the
shape made by its atoms.
3 bonds and 1 lp attached to As

As
H
H
H
trigonal pyramidal 
The End