1. West Midlands Chemistry Teachers Centre November 2009
Structure and Bonding
Presenter: Dr Janice Perkins

2. Ionic Bonding Covalent Bonding Metallic Bonding
Types of Bonding
Ionic Bonding
Covalent Bonding
Metallic Bonding

3. 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

4. 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

5. Covalent Bond Covalent bond = shared pair of e-
One electron comes from each atom

6. Question from June 09
Two organic compounds with similar relative molecular masses are shown below 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 

7. also called Dative covalent bonding
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

8. 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(+) 

9. 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!

10. Metallic bonding + + + + + + + + X X X X
OUTER electrons are ‘DELOCALISED’ i.e. free to move through metal
Metal lattice held together by the ATTRACTION between this sea of delocalised e- and the positive ions in the lattice

11. 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 

12. 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

13. Electronegativity Definition: Trends:
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 e-
Decreases down a group
Due to larger size/more shells, so weaker attraction for bonding e-

14. What’s the sporting Caption?
‘Throwing the puppy dog!’??

15. What’s the sporting Caption?
Wish we were on holiday

16. What’s the sporting Caption? Well this one is clear enough
It must be ‘Tug-o-War’

17. 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

18. 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 e- between 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

19. AB is a polar molecule it has a dipole
Polar Bonds
Atom B is more electronegative than atom A
Bonding e- are attracted more strongly to B + -
Electron density is low.
Electron density is high.
A x B
+A – B -
AB is a polar molecule it has a dipole

20. 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

21. 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

22. 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 + -

23. 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

24. Hydrogen bonding in HF δ+ H F δ- δ+ H F δ- δ+ H F δ- H F H F 
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- δ+
H F δ-
H F
H F

25. Question from Jan 09 H C N O
Electronegativity
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 (+) 

26. 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

27. 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

28. 2D diagram of NaCl
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

29. 3D diagram of NaCl Cl
Cl- -
Each Na+ ion is surrounded by six Cl- ions and vice versa

30. Why are ionic solids brittle animation?

31. Ionic Lattice OK – so what was happening there?
Here it is again with sub-captions for the hard-of-thinking!

32. Why are ionic solids brittle animation?
Like charges repel

33. Metals
In Mg there is an attraction between the delocalised e and the lattice of metal ions
Most metals have high m.p. / b.p. as a lot of energy is needed to remove an atom from this attraction

34. The delocalised e- flow thorough the metal (a current) so metals conduct electricity

35. So how can metals change their shape?
Malleable since the layers can slide over one another

36. Comparison of Na, Mg and Al
delocalised e-
Na Mg 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

37. Giant molecular crystals
Diamond Graphite
Silica

38. 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 e-
Graphite has delocalised electrons between the layers so it conducts electricity.

39. 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

40. 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

41. 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.

42. Equal repulsion between
Linear Shape
B A B
Equal repulsion between
2 bonding pairs

43. Equal repulsion between
Trigonal Planar shape
Equal repulsion between
3 bonding pairs

44. Equal repulsion between
Tetrahedral shape B A
Equal repulsion between
4 bonding pairs

45. Trigonal Bipyramid shape
Equal repulsion between
5 bonding pairs

46. Equal repulsion between
Octahedral shape
Equal repulsion between
6 bonding pairs

47. Regular Shapes 2 B – A – B 3 AB3 4 AB4 5 AB5 6 AB6 180 linear 120
e- pairs
Formula
Bond angle
Name 2
B – A – B
180
linear 3
AB3
120
trigonal planar 4
AB4
109½
tetrahedral 5
AB5 90
trigonal bipyramid 6
AB6
octahedral

48. Working out the shape If it is an ion
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

49. 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

50. 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

51. Ammonia Pyramidal shape Unequal repulsion between
3 bonding pairs and 1 lone pair
Pyramidal shape

52. Working out the shape of the ion NH2-
N 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

53. The shape of the ion NH2- ‘Bent’ or ‘V-shaped’
Unequal repulsion between
2 lone pair and 2 bonding pairs
‘Bent’ or ‘V-shaped’

54. Shapes of molecules / ions

55. 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
trigonal pyramidal  As H

56. The End