Question

The total number of lone pairs of electrons in \(\mathrm{N}_{2} \mathrm{O}_{3}\) is

Short answer

The total number of lone pairs of electrons in (N_{2}O_{3}) is 12.

Step-by-step solution

Understand the concept of lone pairs

Lone pairs refer to pairs of valence electrons that are not shared with another atom and are not involved in bonding. Each atom in a molecule can have a certain number of lone pairs depending on its valence electron count and its bonding behavior.

Determine valence electrons for nitrogen and oxygen

Nitrogen (N) has 5 valence electrons and Oxygen (O) has 6 valence electrons. For the molecule (NO_{2})_{3}, there are two types of nitrogen atoms: the central one and the two terminal ones, and there are three oxygen atoms.

Draw the Lewis structure for N2O3

To determine the number of lone pairs, draw the Lewis structure of (N_{2}O_{3}). For the (N_{2}O_{3}) molecule, the structure is such that the central nitrogen atom is bonded to two oxygen atoms, one with a double bond and one with a single bond. Each terminal nitrogen atom is bonded to one oxygen atom with a single bond.

Count lone pairs on each oxygen atom

Each oxygen atom bonded to nitrogen with a single bond has 3 lone pairs (6 electrons), and the oxygen atom with a double bond has 2 lone pairs (4 electrons).

Count lone pairs on each nitrogen atom

The central nitrogen atom has 1 lone pair (2 electrons), while the terminal nitrogen atoms have no lone pairs since they use their valence electrons to form bonds with oxygen and to maintain a positive formal charge to stabilize the structure.

Calculate the total number of lone pairs in the molecule

Sum the lone pairs from all the oxygen and nitrogen atoms. For (N_{2}O_{3}), there are 3 oxygen atoms with single bonds each contributing 3 lone pairs, giving a total of 9 lone pairs, the oxygen with a double bond contributes 2 lone pairs, for a total of 11 lone pairs from oxygen atoms. The central nitrogen atom contributes 1 lone pair. Therefore, the total number of lone pairs is 11 (from oxygen) + 1 (from the central nitrogen) = 12.

Key concepts

Lewis Structure

The Lewis structure is a valuable tool in chemistry for visualizing the arrangement of atoms within a molecule and how they are connected. It is named after the American chemist Gilbert Lewis, who introduced it in 1916. This representation focuses on the valence electrons, depicting them as dots surrounding the atomic symbols. Chemical bonds between atoms are shown as lines, representing shared pairs of electrons, while lone pairs are dots that belong on individual atoms.

When drawing a Lewis structure, follow these steps: Determine the total number of valence electrons in the molecule, draw the skeletal structure by connecting atoms with single bonds, distribute remaining valence electrons to complete the octets (or duets for hydrogen) of atoms, and finally, modify the structure if necessary to account for multiple bonds if any atom does not have a complete octet. This process helps in understanding the electron distribution and the geometrical shape of the molecule. In our exercise, Lewis structures help us visualize the lone pairs in the N₂O₃ molecule, crucial for identifying the total number of non-bonding electron pairs.

Valence Electrons

Valence electrons are the electrons found in the outermost electron shell of an atom. These are the electrons involved in forming chemical bonds, as they can be lost, gained, or shared in the process. The number of valence electrons is critical for determining an atom's bonding behavior, reactivity, and placement in the Lewis structure.

For example, in the exercise, nitrogen (N) has 5 valence electrons, and oxygen (O) has 6. You can easily calculate this from their positions in the periodic table. Elements in the same group have the same number of valence electrons. In the context of Lewis structures, valence electrons dictate how atoms connect. For instance, oxygen tends to form two bonds because it needs two electrons to complete its octet, leading to two, one, or no lone pairs depending on the bonding. In N₂O₃, the distribution of valence electrons explains why the central nitrogen atom has a single lone pair, while the terminal nitrogen atoms have none, and the varying number of lone pairs on the oxygen atoms.

Chemical Bonding

Chemical bonding refers to the attraction between atoms that enables the formation of chemical substances containing two or more atoms. The bond is caused by the electromagnetic force between oppositely charged ions in ionic bonds or through the sharing of electrons in covalent bonds. The nature of the bond influences the physical properties of the resultant compounds, such as boiling point, melting point, and solubility.

Covalent bonds, which are featured prominently in our N₂O₃ example, involve the sharing of valence electron pairs between atoms. There are also polar covalent and nonpolar covalent bonds; the difference lies in how equally the electrons are shared due to differences in electronegativity. Lone pairs, as mentioned previously, are valence electron pairs that are not shared between atoms and thus do not contribute to bonding but can affect the molecule's shape and reactivity. Understanding chemical bonding helps to make sense of how and why atoms come together to form molecules, and recognizing the types of bonds and lone pairs can lead to deeper insights into a molecule's characteristics and behavior.