1. By the end of this lesson, you should be able to….
14 April, 2017
Covalent Bonding
By the end of this lesson, you should be able to….
Describe and define a covalent bond
Describe single and multiple covalent bonding
Use dot and cross diagrams to represent covalent bonding

2. Covalent Bonds Only occurs between non-metals
Formed by a shared pair of electrons

3. Why is there an attraction?
A covalent bond is directional, it only acts on the atoms involved, unlike an ionic bond

5. Click for another example
Hydrogen and Hydrogen H H H2
Click for another example
Click for animation

6. Nitrogen and Hydrogen (Ammonia)
NH3 H
Click for another example
Click for animation

7. Hydrogen and Oxygen O H2O H H Click here to return to bonding options
Click for animation

8. Task
This is the dot-and-cross diagram for a water molecule. Draw dot-and-cross diagrams for the following: F2 H2
HCl
CH4
SCl2

9. Lone pairs
When an electron pair is not used for bonding, it is known as a lone pair.
A water molecule has 2 lone pairs.
An ammonia molecule has 1 lone pair.
Lone pairs can affect the chemistry of molecules in quite significant ways.

10. Multiple covalent bonds
Some atoms can share more than one pair of electrons to form a multiple bond. O O

11. Carbon Dioxide O O C

13. Task
Draw dot and cross diagrams for the following:
C2H4
HCΞN
c) H2C=O

14. Naming covalent compounds
Two nonmetals
Name each element
End the last element in -ide
Add prefixes to show more than 1 atom
Prefixes
mon penta 5
di hexa 6
tri 3
tetra 4

15. Naming Covalent Compounds
Fill in the blanks to complete the following names of covalent compounds.
CO carbon ______oxide
CO2 carbon _______________
PCl3 phosphorus _______chloride
CCl4 carbon ________chloride

16. Answers CO carbon monoxide CO2 carbon dioxide
PCl3 phosphorus trichloride
CCl4 carbon tetrachloride

17. Dative Bonding
A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.
Lone pair

18. Dative Bonding
You can represent a dative covalent bond with an arrow, AB
The direction of the arrow, shows the direction that the electron pair has been donated.

19. Dative Bonding
When an acid is added to water, oxonium ions can form (H3O+)

20. Dative Bonding
Draw dot-cross diagrams for:
PCl4+
H3O+
H2F+

21. Complications……(as usual)
When covalent bonds form, unpaired electrons pair up to obey the Octet Rule, however this isn’t always possible because:
there may not be enough available electrons to reach an Octet
there may be more than four electrons that pair up during bonding – this is called ‘Expansion of the Octet’

22. Not enough electrons….
Eg: Be and B both form compounds with covalent bonds (despite being metals) – they do not have enough unpaired electrons to reach an octet……

23. Expansion of the octet In groups 5-7 something unusual happens.
Moving down the groups means that more of their outer-shell electrons are able to take part in bonding – this can break the Octet Rule
e.g. phosphorous can form two chlorides, PCl3 and PCl5 it just depends on how much chlorine is available.

24. Task Draw dot-cross diagrams for the following compounds: BF3 PF5 SO2
Extension:
5. BF3NH3 (This compound is formed when BF3 and NH3 react together)

25. Shapes of Molecules
Now we know how covalently bonded compounds form we can look at their shape. On paper they look flat, but many are 3D.
How to predict the shape of a molecule:
Draw the dot and cross diagram
Count the number of electrons surrounding the central atom
Divide this by 2 to find out the number of electron pairs.
These will arrange themselves so that they are as far apart as possible – REMEMBER: They are negative and so repel

26. Use the molymods to make the following 3D structures
CH4
PCl3
NH3
CO2
SF6
NO2+

27. Describing shapes of molecules
The shape of a molecule can be described in terms of its bond lengths and bond angles.
Bond length is the distance between the nuclei of two bonded atoms.
bond angle
Bond angle is the angle between two covalent bonds.
Teacher notes
Bond length depends on which two atoms are bonded together and on the structure of the rest of the molecule. It can be found experimentally using X-ray diffraction. Bond angle can also be measured using X-ray crystallography techniques.
bond length
Counting electrons enables the basic shape of the molecule and its approximate bond angles to be predicted.

28. How to draw 3D molecules

29. Teacher notes
Students could be reminded that the double bond in carbon dioxide means that, although the shape of the molecule is linear, it contains four bonding pairs, not two.

31. For the following covalent compounds do the following:
Draw the dot-cross diagram
Predict and draw the shape of the molecule
Predict the bond angle
CH4
BeCl2
BF3
NH4+
SF6
SiF4

32. Bonding and lone pairs
A pair of electrons in a covalent bond are called a bonding pair. Pairs of electrons that are not involved in bonding are called lone pairs.
bonding pair
lone pair
Electron pairs are clouds of negative charge, so there is mutual repulsion between them, forcing them as far apart as possible.
This means the number of electron pairs around the central atom(s) determines the basic shape of the molecule.

33. Effect of lone pairs on shape
The number of lone pairs in a molecule is calculated by subtracting the number of bonding pairs from the total number of electron pairs in the outer principal energy level.
The shape of a molecule with lone pairs is based on the basic shape for the total number of outer electron pairs, but with a lone pair replacing one of the bonds.
tetrahedral
pyramidal
V-shaped
Teacher notes
Students could be taught the ‘square planar’ molecular shape, as exemplified by xenon tetrafluoride (XeF4). In this structure, four bonding pairs are arranged in the same plane around the central atom with a bond angle of 90°, with a lone pair both above and below the central atom.
V-shaped molecules may also be described as “non-linear”.
replacing one bonding pair with a lone pair
replacing another bonding pair with a lone pair

34. Effect of lone pairs on bond angles
Teacher notes
It may be worth pointing out to students that the dotted line represents a bond extending behind the plane of the screen, and the wedge-shaped bond represents a bond extending in front of the plane of the screen.

35. SiCl4 PCl4+ PCl6- SiCl62- H2S XeF4 QUESTIONS
For each of the following ions/molecules, state the number of bond pairs, state the number of lone pairs, state the bond angle(s), state, or draw, the shape
SiCl4
PCl4+
PCl6-
SiCl62-
H2S
XeF4

38. Nitrogen and Boron can form the chlorides NCl3 and BCl3.
Draw dot cross diagrams to show the bonding in both NCl3 and BCl3.
Draw the shapes of both of these molecules. Show the approximate values of the bond angles and the name of the shape.
Explain why the shape of both of these molecules is different.

39. Effect of lone pairs on bond angles

41. Strength of metallic bonding: ion charge
The strength of metallic bonding depends on two factors:
1. the charge on the metal ions
2. the size of the metal ions.
1. The charge on the metal ions
The greater the charge on the metal ions, the greater the attraction between the ions and the delocalized electrons, and the stronger the metallic bonds. A higher melting point is evidence of stronger bonds in the substance.
Element
Charge on ion
Melting point (K) Na Mg Al 1+ 2+ 3+
371
923
933

42. Strength of metallic bonding: ion size
2. The size of the metal ions
The smaller the metal ion, the closer the positive nucleus is to the delocalized electrons. This means there is a greater attraction between the two, which creates a stronger metallic bond.
Element
Ionic radius (nm)
Melting point (K) Li Na K Rb Cs
0.076
0.102
0.138
0.152
0.167
Teacher notes
The ionic radius of an ion is not fixed, but depends on several factors, such as the co-ordination number of the ion.
454
371
337
312
302

43. Types of bonding

44. Discuss the bonding, structure and properties of the following substances. Points to consider and an example follow.
sodium chloride
aluminium
water
iodine
diamond carbon
graphite carbon
hydrogen chloride
substances:

45. C (diamond) C(graphite) HCl
NaCl Al H2O I2
C (diamond) C(graphite) HCl
Which types of chemical bonding are present within the substance in the solid, liquid and gaseous states.
The name given to the type of substance.
The nature of the bonding present.
The arrangement of particles in space within the solid, liquid and gaseous states.
How strongly the particles are held together.
What happens to the particles and bonds present as the substance is gradually warmed from below its melting point to above its boiling point.
Points to consider:

46. Aluminium oxide
Which types of chemical bonding are present within the substance in the solid, liquid and gaseous states.
Ionic bonds are present in the solid and liquid states
No bonds are present in the gaseous state
The name given to the type of substance.
Aluminium oxide is an ionic compound

47. Aluminium oxide The nature of the bonding present.
Ionic bonds are the electrostatic attraction between oppositely charged ions
The arrangement of particles in space within the solid, liquid and gaseous states.
In the solid the ions are fixed in place within a crystal lattice
In the liquid the ions can move through the substance but are still held close together
In the gas the ions are completely separate from each other and move completely freely

48. Aluminium oxide How strongly the particles are held together.
Ionic bonds are strong attractions
What happens to the particles and bonds present as the substance is gradually warmed from below its melting point to above its boiling point.
In the solid the ions vibrate about fixed points
At the melting point some attractions between ions break
The melting point is relatively high because ionic bonds are strong

49. Aluminium oxide
What happens to the particles and bonds present as the substance is gradually warmed from below its melting point to above its boiling point.
In the liquid the bonds between ions are constantly breaking and reforming so that they can move through the substance
At the boiling point all ionic attractions break
The boiling point is relatively high because ionic bonds are strong
In the gas the ions move completely freely of each other
In all three states the ions move faster when heated