The number of valence electrons for an atom in a Lewis structure is equal to the number of electrons in its outermost energy level, also known as the valence shell. For nitrogen, this number is 5.
The general rule for drawing Lewis structures is to start by arranging the atoms in a molecule, connecting them with single bonds. Then, distribute the valence electrons around the atoms in order to fulfill the octet rule, which states that most atoms tend to form bonds in order to have 8 electrons in their outermost energy level. For nitrogen, this means it will typically form 3 covalent bonds and have one lone pair of electrons.
In some cases, atoms such as nitrogen may need to form double or triple bonds in order to fulfill the octet rule. This is especially true for molecules with elements from the third period of the periodic table and beyond. To determine when to use double or triple bonds, count the number of valence electrons in the molecule and distribute them accordingly, keeping in mind that each bond consists of 2 electrons.
While most atoms in a Lewis structure strive for a complete octet, there are some exceptions, including nitrogen. In certain molecules, nitrogen may have an incomplete octet, meaning it does not have 8 electrons in its outermost energy level. This is common in molecules such as nitric oxide (NO) or nitrogen dioxide (NO2).
Typically, the most electronegative atom in a molecule is placed in the center, with the other atoms surrounding it. This helps to minimize the formal charges on each atom, making the Lewis structure more stable. For nitrogen-containing molecules, nitrogen is usually placed in the center, unless it is bonded to a more electronegative atom such as oxygen or fluorine.