Write the Lewis structure for each ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. MISSED THIS? Road Soction 10.9 ; Watch KCV \(10.9,\) HE 10.10 a. \(\mathrm{PO}_{4}^{3-}\) b. CN" c. \(\$ O_{3}^{2-}\) d. \(\mathrm{ClO}_{2}\)

The formal charge of an atom in a molecule or ion can be calculated by taking the number of valence electrons it originally has, subtracting the number of non-bonding electrons, and then subtracting half the number of bonding electrons. This concept is a bookkeeping device that helps us predict the most likely structure for a molecule or an ion. To calculate the formal charge, use the formula:
\[ \text{Formal Charge} = (\text{Valence Electrons}) - (\text{Non-bonding Electrons}) - \frac{1}{2}(\text{Bonding Electrons}) \]
Understanding formal charges is crucial for determining the most stable Lewis structure. Structures with formal charges closest to zero are generally the most stable. If there are differences in electronegativity between bonded atoms, the more electronegative atom is expected to take on a more negative formal charge.

Resonance structures are different ways of drawing the Lewis structure for a molecule or ion when there is more than one valid way to position the electrons. Resonance structures do not represent distinct entities, but rather contribute to a single, hybrid representation of the molecule, known as a resonance hybrid.
When drawing resonance structures, the positions of the atoms remain the same, but the locations of the electrons, particularly the pi electrons or non-bonding pairs, change. To represent resonance, a double-headed arrow is used between the structures. Considering all possible resonance structures allows us to understand the distribution of electron density in the molecule better and to predict its reactivity and properties. It's important to note that the resonance hybrid is more stable than any individual resonance structure.

The octet rule is a simple chemical principle that reflects the tendency of atoms to bond in such a way that they each have eight electrons in their valence shell, giving them the same electronic configuration as a noble gas. The rule applies to main-group elements, although there are exceptions, such as hydrogen, which can hold a maximum of two electrons.
In Lewis structures, atoms form enough bonds to obtain eight electrons through shared pairs, achieving what is called stable electron configurations. Elements from periods 3 and beyond can expand their octets because they have d-orbitals available for bonding. Understanding the octet rule is critical when predicting the arrangement of atoms in a molecule and in identifying the need for lone pairs or multiple bonds to satisfy the octet requirement.

Valence electrons are the outermost electrons of an atom and are critical in determining how the atom can bond with others. In the case of main-group elements, the valence electrons are found in the outermost s and p orbitals. These are the electrons used to form chemical bonds, and they can be readily lost, gained, or shared during chemical reactions.
When drawing Lewis structures, the total number of valence electrons gives insight into the overall bonding scheme, including how many bonds can form and whether there will be non-bonding (lone pair) electrons. In ions, valence electrons include the electrons gained or lost in addition to the neutral atom's valence electrons. Accurately counting valence electrons is the first step in creating the right Lewis structure and understanding the chemical properties and bonding of the molecule or ion.