1. Lewis Dot Structures of Covalent Compounds Atoms are made up of protons, neutrons, and electrons. The protons and neutrons are located at the center of the atom, the nucleus. These electrons can be divided into core electrons and valence electrons. The valence electrons are the outermost electrons and are the ones involved in chemical reactions

2. Rules for writing Lewis Dot structures Rule 1 Add together the number of valence electrons for each atom in the molecule. For example, CF 4 Carbon has four valence electrons and each fluorine has seven valence electrons = 4 + 4(7) = 32

3. Rule 2 Write out the elements of the molecule so that the least electronegative elements is in the center surrounded by the other elements. For example, CF 4

4. Rule 3 Place a covalent bond between the central atom and the outside atoms. Remember each covalent bond contains two electrons.

5. The four covalent bonds use eight of the 32 valence electrons in CF 4 Rule 4 There are 24 valence electrons remaining. Add electrons to the outer atoms as lose pairs to satisfy the Octet Rule. This uses 24 electrons. There Are no electrons left, so this is The Lewis dot structure for CF4

6. Rule 5 for example, NH 3 First apply Rules 1-4 to the molecule Rule 1: Count the valence electrons Rule 2: Place the least electronegative element at the centre, except for H which is always an outer atom Rule 3: Add covalent bonds between the centre atom and the outer atoms Rule 4: Add lone pairs to the outer atoms Rule 5: Add lone pairs to the centre atom

7. Rule 1 Nitrogen has 5 valence electrons and each hydrogen has 1 valence electron The total number of valence electrons = 5 + 3 (1) = 8 Rule 2 Hydrogen is always an outer atom and is never at the centre of a molecule

8. Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons. Rule 4 The 2 remaining valence electrons are not added to the outer atoms, because each H has its maximum of 2 valence electrons.

9. Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons. Rule 4 The 2 remaining valence electrons are not added to the outer atoms, because each H has its maximum of 2 valence electrons.

10. Rule 5 Place the remaining 2 Valence electrons on the central nitrogen atom Rule 6 Check all atoms in the molecule to ensure that each has 8 electrons(2 for hydrogen). If an atom has fewer than 8 electrons, create double or triple bonds. (Note: Double bonds only exist between C,N,O and S atoms) This is the Lewis structure For NH 3

11. Apply rule 6 to the following; CH 4, CF 4, Hydrogen : 1 bond = 2 electrons (stable) Carbon : 4 bonds = 8 electrons (stable) Fluorine : 1 bond + 3 lone pairs = 2 + 3 (2) = 8 electrons (stable) Carbon : 4 bonds = 8 electrons (stable)

12. Example; CH 2 O Apply Rules 1-5 to the molecule Rule 1: Count the valency electrons Rule 2: Place the least electronegative element at the centre, except for H, which is always an outer atom Rule 3: Add covalent bonds between the centre and the outer atoms Rule 4: Add lone pairs to the outer atoms Rule 5: Add lone pairs to the centre atom

13. Rule 1 Carbon has 4 valence electrons, each hydrogen has 1 valence electron, and oxygen has 6 valence electrons. Total number of valence electrons : 4 + 2(1) + 6 = 12 Rule 2 Carbon is at the centre of the molecule because it is less electronegative than oxygen. Hydrogen is always an outer atom and is never at the centre of the molecule.

14. Rule 3 Add the bonding electrons. This uses 6 of the 12 valence electrons Rule 4 Add the remaining 6 lectrons to the outer atom. Hydrogen does not need any more electrons, but Oxygen needs 6 to complete its octet.

15. Rule 6 Oxygen shares one of its lone pairs with C and O and give the desired 8 electron total Rule 5 There are no valence electrons left to add to the centre This is the Lewis dot Structure for CH 2 O

16. Exceptions to the Octet Rule The Octet Rule applies to Groups IVA through VIIA in the second row of the Periodic Table, but there are exceptions to the rule among some other elements. The following two cases are an example Example BF 3 Rule 1 Boron has 3 valence electrons and each Fluorine has 7 valence electrons Total number of electrons = 3 + 3 (7) = 24

17. Rule 2 Boron is at the centre of the molecule because it is less electronegative than fluorine Rule 3 Add the bonding electrons. This uses 6 of the 24 valence electrons

18. Rule 4 Add the remaining electrons to the outer atoms. Each Fluorine has the required 8 electrons Rule 5 This uses the remaining electrons leaving none to add to the Boron central atom

19. Rule 6 Check the number of electrons around each atom. Each Fluorine atom has 8 electrons, but the Boron Atom has only 6. This is an exception to the Octet Rule. A B=F bond is not an option, because double bonds exist only between C,N,O, and S atoms This is the Lewis dot structure BF 3

20. Example PF 5 Rule 1 Phosphorus has 5 valence electrons and each fluorine has 7 valence electrons Total number of electrons = 5 + 5(7) = 40 Rule 2 Phosporus is at the centre because it is less electronegative than fluorine

21. Rule 3 Add the bonding electrons. This uses 6 of the 24 valence electrons. Rule 4 Add the remaining electrons to the outer atoms. Each Fluorine requires 6 more electrons

22. Rule 5 This uses the remaining electrons leaving none to the central P atom Rule 6 Check the number of electrons around each atom. Each Fluorine atom has 8 electrons, but the phoshorus atom has 10. This is an exception to the Octet Rule.

23. Rule 6 Check the number of electrons around each atom. Each fluorine atom has 8 electrons, but the phoshorus atom has 10. This is an exception to the Octet Rule.