1. Bonding

2. Octet Rule
Octet Rule states that when bonding occurs atoms tend to reach an arrangement with 8 electrons in the outer shell
Exceptions:
- d-block elements, including the transition
metals (variable valency)
- small elements like He, Li (want full outer
shell)

3. Valency
The valency of an element is the number of bonds an atom of the element forms when it reacts
Valency is equal to the number of electrons an atom needs to lose or gain when forming a bond

4. Group number:
Valency: I 1 II 2
III 3 IV 4 V VI
VII
VIII

5. Transition metals have a variable valency
They can lose different amounts of electrons depending on what they bond with

6. Ionic Bonding An ion is a charged atom or group of atoms
When an atom loses an electron it forms a positive ion (more protons than electrons)
When an atom gains an electron it forms a negative ion (more electrons than protons)
An ionic bond is the force of attraction between oppositely charged ions

7. Examples:
Sodium chloride (NaCl)
Na + Cl → Na+ Cl-
2,8,1 2,8, ,8 2,8,8
Na + Cl → Na + Cl -

8. Magnesium fluoride (MgF2)
Mg + F + F → Mg2+ F- F-
2,8, , , , ,8 2,8
Mg F F → Mg F F -

9. Group ions: Name: Formula: Hydroxide ion Nitrate ion
Hydrogencarbonate ion
Permanganate ion
OH-
NO3-
HCO3-
MnO4-
Carbonate ion
Chromate ion
Dichromate ion
Sulfate ion
Sulfite ion
Thiosulfate ion
CO32-
CrO42-
Cr2O72-
SO42-
SO32-
S2O32-
Phosphate ion
PO43-
Ammonium ion
NH4+

10. Ionic structure
Ions are arranged in compounds in a crystal lattice structure
Crystal lattices are made up of unit cells linked together (repeating unit)
E.g., NaCl molecule = unit cell
NaCl compound = crystal lattice

11. Ionic compound formulas
Chemical formula represents the compound using chemical symbols and numbers
Have to be able to work out ionic formulas for 1st 36 elements
Ionic bonds usually form between Group I & II and Groups VI & VII

12. Group I & II (metals) lose electrons and form positive ions,
e.g., Li → Li+ , Mg → Mg2+
Group VI & VII (non-metals) gain electrons and form negative ions,
e.g., O → O2-, Cl → Cl-

13. These oppositely charged ions are attracted to each other forming an ionic bond,
e.g., Li+ + Cl- → LiCl
Li ion Cl ion Neutral LiCl
compound
Mg O2- → MgO
Mg ion O ion Neutral MgO
Ionic componds are neutral because the charges of the ions that form them should balance

14. Covalent bonding
A covalent bond consists of one or more shared pairs of electrons
A molecule is a group of atoms joined together. It is the smallest part of an element that can exist independently.
Need to know 5 examples:

15. (a) Hydrogen molecule
- H has 1 electron in its outer shell so needs to gain another electron to have a full outer shell
- The two H atoms share their electrons to make them stable
H H

16. (b) Chlorine molecule
- Cl has 7 electrons in its outer shell so needs to gain another electron to have a full outer shell
- The two Cl atoms share 1 of their electrons with the other to make them stable
Cl Cl

17. (c) Water molecule
- H has 1 electron its outer shell so needs to gain another electron to have a full outer shell
- O has 6 electrons in its outer shell so needs to gain another 2 to have a full outer shell
- Each H atom shares its electron with the O atom to make them stable O
H H

18. (d) Ammonia molecule
- H has 1 electron its outer shell so needs to gain another electron to have a full outer shell
- N has 5 electrons in its outer shell so needs to gain another 3 to have a full outer shell
- Each H atom shares its electron with the N atom to make them stable N
H H H

19. (e) Methane molecule
- H has 1 electron its outer shell so needs to gain another electron to have a full outer shell
- C has 4 electrons in its outer shell so needs to gain another 4 to have a full outer shell
- Each H atom shares its electron with the C atom to make them stable H
H C H

20. Types of covalent bonds
Single bonds share one pair of electrons
e.g., H Cl
Double bonds share two pairs of electrons
e.g., O O
Triple bonds share three pairs of electrons
e.g., N N

21. Sigma bond is the head on overlap of two atomic orbitals
Pi bond is the side on overlap of two atomic orbitals
Single bond = sigma bond
e.g.,
H Cl

22. Double bond = 1 sigma & 1 pi bond
e.g.,
O O
Triple bond = 1 sigma & 2 pi bonds
N N

23. Electronegativity
In a covalent bond between two atoms of the same element the pair of electrons in the bond is shared equally
But in a covalent bond between atoms from different elements one of the atoms has a greater attraction for the electrons than the other
This can cause a slight charge to form on the atoms in the bond

24. Example - HCl
The Cl atom is more attracted to the electrons than the H atom,
δ+H Cl δ-
Electronegativity is the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond

25. Very reactive elements, like K and Na, have low electronegativity values – can be called electropositive
Very unreactive elements, like F and Cl, have high electronegativity values

26. Increase across a period because:
Increase in nuclear charge –
Increase in number of protons but electrons all added to same shell
(b) Decrease in atomic radius –
Electrons are closer to the nucleus
(c) No screening effect –
Electrons are added to the same shell so there are no inner electrons to shield nuclear charge

27. Decreases down a group because:
Increase in atomic radius –
Electrons are further away from the nucleus
(b) Screening effect –
Inner electron shells shield the outer ones from the nuclear charge
(c) Decrease in nuclear charge –