1. Chemical Bonding

2. Why the noble gaes are so noble…
The elements in group 18 are known as the noble gases.
They are completely unreactive
The other elements desperately want to be like them but to achieve this they need to change their electron configuration.
What do all of group 18 have in common in terms of their electron configuration?

3. Formation of ions
The other elements try to achieve full outer shells by either gaining or losing electrons when they react to form compounds.
Metal Elements in groups 1 and 2 only have 1 or 2 outer shell electrons, they are able to lose these to become positively charged ions called cations.
Most non-metal elements have more than three electrons in their outer shell. They are able to gain electrons to become negatively charged ions called anions.
There is a limit to how many electrons can be lost or gained and 3- anions or 4+ cations are highly unusual. Why?

4. Ionic Bonding Metals react with non-metals in a chemical reaction.
Electrons transferred from metals to non-metals.
Resulting cations and anions are held together by an electrostatic bond due to their opposite charges. Also known as an ionic bond.
Electron dot-cross diagrams are used to represent the way the atoms bond together.

5. Properties of ionic compounds
High melting points
Usually dissolve in water
Conduct electricity when molten or dissolved in water.
Now attempt 3.1 Problem 1 a-h
In a solid each ion may attract many other ions of the opposite charge and this can lead to giant ionic structures such as the giant lattice below:

6. Covalent Bonding
Covalent bonds are formed between two non-metallic elements
Electrons are shared instead of transferred
Shared electrons count as part of the outer shell for both elements in the covalent bond.
The resulting compounds are more stable than the individual elements

7. Covalent Bonding
The two atoms are held together because their positive nuclei are attracted to the pair of negative electrons
Each atom donates one electron to the pair
A bond involving one electron pair (2 single electrons) is called a single covalent bond, or just a single bond
Two pairs = Double bond
Three pairs = Triple bond

8. Covalent Bonding
Pairs of electrons involved in bonding are called ‘Bonding pairs’ and pairs of electrons not involved are called ‘Lone Pairs’.
Electron dot-cross diagrams can be used to show covalent bonding but a simpler way is using lines to represent the bonding pairs.
What would (a) Water
(b) Oxygen
(c) Nitrogen look like?
Now try 3.1 problem 2 a-i

9. Dative Covalent Bonding
This is a type of covalent bonding where one atom supplies both electrons to the covalent bond.
An example is carbon monoxide where the third bonding pair of electrons are both donated by the oxygen atom. This can be represented by an arrow.
E.g. C O
Now try 3.1 problems 3-5

10. Why are bonds like bears?
But what if the atoms involved are not identical?

11. Size matters…
A smaller atom will have its ‘atomic core’ closer to the electron pair. This will make the attraction greater.
The atomic core is made up of everything except the outer core electrons.
The ‘core charge’ also has an effect.
Fluorine has an electron configuration of 2.7, the nucleus charge is +9 and the inner electrons give a charge of -2. Overall core charge is +7.
The greater the core charge, the greater the attraction to the bonding pair of electrons.

12. And the bears?
Generally, different atoms attract to the bonding electrons unequally.
One atom gets slightly more negative charge from the bonding pair as it is attracted more strongly and conversely the other atom gets a slightly more positive charge as it is attracted less strongly to the bonding pair.
This is known as a ‘polar bond’…and just like bears, some bonds are polar and some are not!

13. Polar Bonds
A polar bond is represented by symbol ‘δ’ or delta. This means ‘small amount of’.
E.g. O δ-
H δ+ H δ+

14. Electronegativity
In order to decide the polarity of a covalent bond we need a measure of its ‘electron pulling power’. This is known as ‘electronegativity’.
The better the atoms ‘pulling power’, the higher its electronegativity.
Some examples are shown below:

15. Polarity
We use these electronegativity values to predict whether a bond will be polar.
A highly polar bond such as C-F will have a large difference in electronegativity between the two atoms.
E.g. C = 2.6 and F = 4.0
A non-polar bond will have two atoms with similar electronegativity values such as C-H
E.g. C = 2.6 and H = 2.2

16. The Bonding Continuum
Covalent and Ionic bonds are two extremes in the bonding continuum. You can think of polar bonds as somewhere in between.
Polar bonds are essentially covalent bonds where the electrons are unequally shared.
The bigger the difference in electronegativity , the more polar the bond and the greater the ionic character of the bond.
Now try 3.1 Problems 7-8

17. Metallic Bonding
Metals cannot share or transfer enough electrons to have full outer shells so they need another way to do it…
Metal atoms lose their outer electrons to form a ‘pool of delocalised electrons’. The remaining positive ions form a regualrly spaced lattice.
The positive ions are attracted to the pool of delocalised electrons and these attractions are called metallic bonds.

18. Metallic Bond Strength
Strength of metallic bonds depends upon a number of factors including:
number of electrons per atom available for delocalisation
E.g. Magnesium has twice as many outer shell electrons as sodium and therefore has higher MP and BP.
Now try 3.1 Problems 8-9

19. And finally…Writing Chemical Formulae….
Table 3 on page 42 of CI lists a number of common ions. There is no way round it…these need to be learnt!!
They will be used throughout the course for constructing chemical formulae.
When constructing chemical formulae you need to remember to balance the charges as shown in examples on pages 42-43
Now try 3.1 problems 10-11