To draw the Lewis structure for NF3 (nitrogen trifluoride), follow these steps:
1. **Count the total number of valence electrons**: Nitrogen (N) has 5 valence electrons, and each fluorine (F) atom has 7 valence electrons. Since there are three fluorine atoms, the total number of valence electrons is 5 + (3 × 7) = 26.
2. **Place the least electronegative atom in the center**: Nitrogen is less electronegative than fluorine, so it will be the central atom.
3. **Connect the central atom to the surrounding atoms with single bonds**: Draw single bonds from the nitrogen atom to each of the three fluorine atoms. This uses up 6 electrons (3 bonds × 2 electrons).
4. **Distribute the remaining electrons**: After forming the bonds, you have 20 electrons left (26 – 6). Place these electrons around the fluorine atoms to complete their octets. Each fluorine atom needs 6 more electrons to complete its octet, so you will place 6 electrons around each fluorine atom. This uses up all 20 electrons.
5. **Check the octet rule**: Nitrogen has 8 electrons around it (2 from each bond and 2 lone pairs), and each fluorine atom has 8 electrons (6 lone pairs and 2 from the bond).
**Bonds and Non-Bonding Pairs**:
– **Bonds**: There are 3 single bonds between nitrogen and each fluorine atom.
– **Non-Bonding Pairs**: Nitrogen has 1 lone pair (2 electrons) that are not involved in bonding.
**Molecular Shape**:
The shape of NF3 is trigonal pyramidal. This is because the central nitrogen atom has three bonding pairs and one lone pair, which repel each other to form a pyramid-like shape. The lone pair occupies more space than the bonding pairs, causing the fluorine atoms to be pushed slightly closer together, resulting in a bond angle of approximately 107 degrees.