Bonding
All atoms bond together with other atoms to achieve a full outer energy level to become stable. Group 18 however, already have a full valence shell so they are an exception. Atoms bond by sharing electrons to form covalent bonds or they give away or accept electrons to form ionic bonds.
Lewis diagrams
Lewis diagrams are used to represent the covalent bonds present in molecules and compounds. Only the electrons in the outer energy level are represented because it is only these electrons involved in the bonding.
Steps to draw lewis diagrams:
1. Draw atoms around the central atom. The central atom is the one that has the most incomplete valence shell.
2. Count the total number of valence electrons. You must use all of these electrons.
3. You must arrange this number of electrons around the atoms so that each atom has 8 electrons around it. The only atom that should not have 8 electrons around it is hydrogen - this is because it only needs 2 electrons to fill it's valence shell.
4. The two electrons between two atoms can be replaced with a bond - a straight line
5. If any atom has an incomplete valence shell, move electrons to create double bonds or triple bonds.
There are different steps you can follow to draw lewis diagrams, I prefer the simple steps above but if you would like more details checkout the following: ABA workbook page 43, or ESA guide page 102.
Steps to draw lewis diagrams:
1. Draw atoms around the central atom. The central atom is the one that has the most incomplete valence shell.
2. Count the total number of valence electrons. You must use all of these electrons.
3. You must arrange this number of electrons around the atoms so that each atom has 8 electrons around it. The only atom that should not have 8 electrons around it is hydrogen - this is because it only needs 2 electrons to fill it's valence shell.
4. The two electrons between two atoms can be replaced with a bond - a straight line
5. If any atom has an incomplete valence shell, move electrons to create double bonds or triple bonds.
There are different steps you can follow to draw lewis diagrams, I prefer the simple steps above but if you would like more details checkout the following: ABA workbook page 43, or ESA guide page 102.
H2O
1. Oxygen has the most incomplete valence shell so should be the central atom
2. Oxygen has an atomic number of 8, this means there are 2 electrons in the first shell and 6 in the second. This means oxygen has 6 valence electrons. Each hydrogen has 1 valence electron. This gives a total of 8 valence electrons
3. These 8 electrons need to be arranged so the oxygen has 8 electrons arranged around it and each hydrogen has 2
4. The 2 electrons between the oxygen and hydrogen atoms are replaced with bonds
Final check - have I used all 8 electrons?
1. Oxygen has the most incomplete valence shell so should be the central atom
2. Oxygen has an atomic number of 8, this means there are 2 electrons in the first shell and 6 in the second. This means oxygen has 6 valence electrons. Each hydrogen has 1 valence electron. This gives a total of 8 valence electrons
3. These 8 electrons need to be arranged so the oxygen has 8 electrons arranged around it and each hydrogen has 2
4. The 2 electrons between the oxygen and hydrogen atoms are replaced with bonds
Final check - have I used all 8 electrons?
SCl2
1. S has the most incomplete valence shell so is in the middle. 2. The electron configuration of S is 2,8,6 therefore it has 6 valence electrons. Cl has an electron configuration of 2,8,7 therefore each Cl has 7 valence electrons. This gives a total of 20 valence electrons. 3. Arrange these 20 electrons around the atoms so each atom has 8 electrons around it. 4. Replace the 2 electrons between every 2 atoms with a bond. |
|
CO2
1. Carbon has 4 valence electrons and each oxygen has 6. This gives a total of 4+6+6 = 16 valence electrons. 2. Carbon has the least filled valence shell so it goes in the middle 3. Arrange these 16 electrons around the atoms so each atom has 8 electrons around it. 4. Replace every 2 electrons between every 2 atoms with a bond. |
|
Electronegativity
Within a covalent bond, different atoms attract the shared electrons within the bond more strongly than others - this is called electronegativity. Each of the atoms have been given an electronegativity value as shown above, except group 18 as they have stable electron configurations.
Non-polar bonding
If the 2 atoms that are covalently bonded together are identical or have VERY similar electronegativities (<0.4 difference), they share the electrons equally - this is called non-polar bonding.
For example, O2, the four shared electrons between the two oxygen atoms spend an equal amount of time moving around each oxygen atom. Remember, electrons are always on the move! They are not stuck how we draw them. |
Polar bonding
If the 2 atoms that are covalently bonded together are different, it is very likely that the two atoms do not share the electrons within the bond equally. The more electroneagtive atom has a greater pull on the electrons and they spend more time around that atom than the least electronegative atom.
For example, chlorine is more electroneagtive than hydrogen therefore Cl has a greater attraction for the shared electrons and they spend more time orbiting the Cl than the H. This creates polarity and what we call a dipole - this shows that one of the atoms is slightly more negative and the other is more positive. |