In this example problem, we will draw the Lewis structure for phosphate. The first step is to sum the valence electrons. There is one phosphorus atom with five valence electrons, and there are four oxygen atoms, each with six valence electrons. The charge is three minus, which increases the number of electrons by three. So, in total, we have 32 valence electrons. The second step is to draw the structure with only single bonds. The central atom is phosphorus, and the four oxygen atoms are attached to it. The third step is to satisfy the octet and duet rules by adding lone pairs. First, we will satisfy the octet rule for peripheral atoms. If we run out of electrons, we can create double and triple bonds. If we have extra electrons, we can put them on the central atom if that atom is an element in the third row or below in the periodic table. So, first, satisfying the octet rule for the oxygen atoms, we'll want to add six additional electrons to each oxygen atom. If we count up the number of electrons that we've used, we see that we've used all 32 of our valence electrons. We also see that the octet rule is satisfied for all atoms. Next, we'll want to use formal charge to evaluate the Lewis structure. We can calculate the formal charge of the oxygen atoms. Oxygen has six valence electrons in a neutral atom, and each oxygen atom has one bond plus six lone pair electrons. So the formal charge of oxygen is 6 - 1 + 6, which is negative 1. We can calculate the formal charge of phosphorus. The neutral atom of phosphorus has five valence electrons, and this phosphorus has four bonds and zero lone pair electrons. So its formal charge is +1. We see that there's both a plus one and a negative one formal charge in our Lewis structure. So we can try to reduce the number of formal charges. If we create a double bond between phosphorus and oxygen using one of the lone pairs on one of the oxygens, then we can change the formal charges on both oxygen and phosphorus. We can calculate the new formal charge of this oxygen atom, and it will be 6 - 2 bonds + 4 lone pair electrons, which is zero. And to calculate the new formal charge on phosphorus, it's 5 - 5 bonds + 0 lone pair electrons, which is equal to zero. So in this Lewis structure, we have fewer formal charges. The Lewis structure that has fewer and smaller formal charges should be the better structure. Phosphorus is able to accommodate ten electrons instead of eight because it's in the third row of the periodic table or below. This means that those extra electrons are occupying the D orbitals. We know that this is one valid Lewis structure, but we could have created the double bond with any of the four oxygen atoms. So we can create these four resonance structures for the phosphate anion. To indicate that these are anions, we can put the Lewis structure in brackets and indicate the charge. So, there are four resonant structures for the phosphate anion.