Question

Draw Lewis structures for (a) the peroxide ion, \(\mathrm{O}_{2}{ }^{2-}\), and (b) \(\mathrm{POCl}_{3}\) (the central atom is phosphorus, which is attached to the oxygen and the three chlorines).

Short answer

The Lewis structure for the peroxide ion, \textrm{O}_2^{2-}, consists of two oxygen atoms double-bonded to each other, each with a -1 charge. For \textrm{POCl}_3, the central phosphorus atom is single-bonded to three chlorine atoms and one oxygen atom, with all atoms satisfying the octet rule.

Step-by-step solution

Lewis Structure for Peroxide Ion (\textrm{O}_2^{2-})

Start by calculating the total number of valence electrons available, which is the sum of the valence electrons of all the atoms plus the extra electrons due to the negative charge. Oxygen atoms have 6 valence electrons each, so for two oxygens, that's 12 electrons. The charge of 2- adds 2 more electrons, for a total of 14 valence electrons. Draw two oxygen atoms side by side and connect them with a single bond (2 electrons). Distribute the remaining 12 electrons to satisfy the octet rule for both oxygen atoms, resulting in each oxygen atom having 8 electrons around it. As the oxygen atoms share two bonds (4 electrons), they form a double bond. The structure will show two oxygen atoms double-bonded to each other, each with a -1 charge, resulting in the overall 2- charge.

Lewis Structure for \textrm{POCl}_3

The total number of valence electrons is the sum of the valence electrons of phosphorus (5), oxygen (6), and each chlorine (7), which gives us 5 + 6 + (7 x 3) = 32 electrons. The central phosphorus atom is drawn and single bonds are connected from it to the oxygen and the three chlorine atoms, using 8 electrons. This leaves 24 electrons to distribute. Complete the octets of the chlorine atoms and oxygen atom by placing the remaining electrons around them, with 6 around oxygen and 18 around the chlorines (6 for each). Ensure that each chlorine and the oxygen have 8 electrons each. Since all the valence electrons are used up and every atom has a complete octet, no double or triple bonds are necessary in this structure.

Key concepts

Valence Electrons

Valence electrons play a pivotal role in chemical bonding as they are the outermost electrons of an atom and therefore the most likely to be involved in bonds. They are responsible for the chemical properties of elements and determine how an atom can bond with others. For example, oxygen has six valence electrons while phosphorus has five and chlorine has seven.

When drawing Lewis structures, as seen in the exercise with the peroxide ion \(\mathrm{O}_2^{2-}\) and \(\mathrm{POCl}_3\), the first step is to calculate the total number of valence electrons that are available. This includes the sum of the valence electrons from each atom and any additional electrons that may be indicated by a charged species.

Octet Rule

The Octet Rule is a fundamental concept in chemistry that reflects the tendency of atoms to prefer having eight electrons in their valence shell, replicating the electron configuration of noble gases. This rule is the guiding principle for determining the arrangement of electrons in many molecules and ions.

When we applied this rule to our exercise, we made sure that the oxygen atoms in the peroxide ion \(\mathrm{O}_2^{2-}\) and the atoms in \(\mathrm{POCl}_3\) ended up with complete octets. In the peroxide ion, despite the additional negative charge, each oxygen atom was depicted with eight electrons to fulfill the octet rule through the formation of a double bond.

Ionic and Covalent Bonding

Ionic and covalent bonds are two major types of chemical bonds that hold atoms together in a compound. Ionic bonding occurs when there is a transfer of electrons between atoms, typically between metals and nonmetals, leading to the formation of positively and negatively charged ions. In contrast, covalent bonding involves the sharing of electrons between atoms, usually between nonmetal atoms, allowing them to achieve a stable electron configuration.

Our exercise features covalent bonding, as evidenced by the peroxide ion and \(\mathrm{POCl}_3\). In the Lewis structures for these molecules, atoms share electrons to fulfill their octets, which is a hallmark of covalent bonding. Notably, in the peroxide ion, each oxygen shares two pairs of electrons, thus forming a double bond between them.