Question

Write the Lewis structure for each molecule or ion. \begin{equation}\mathrm a.{H}_{3} \mathrm{COCH}_{3} \text { b. } \mathrm{CN}^{-}\end{equation}\begin{equation} \mathrm c.{NO}_{2}^{-} \quad \text { d. } \mathrm{ClO}^{-}\end{equation}

Short answer

For H3COCH3, the central C atoms are bonded to each other with single bonds to H and a double bond to O. For CN-, a triple bond connects C and N with a lone pair on N. For NO2-, a bent structure with one N=O double bond, one N-O single bond, and a lone pair on N. For ClO-, a single bond connects Cl and O, with three lone pairs on both Cl and O.

Step-by-step solution

Counting Valence Electrons for H3COCH3

Firstly, calculate the total number of valence electrons for the molecule H3COCH3 (acetone). Each H atom contributes 1 electron, each C atom contributes 4, and the O atom contributes 6, for a total of (3*1) + (2*4) + 6 = 18 valence electrons.

Drawing the Skeletal Structure for H3COCH3

Draw the skeletal structure with the less electronegative atoms, carbon, at the center. Connect the C atoms to each other, and connect H atoms to C, and O atom to one of the C atoms. Use single bonds initially.

Completing the Octet and Distributing Electrons for H3COCH3

Place the remaining electrons starting with the outer atoms to satisfy the octet rule. Each H atom needs 2 electrons, and each C and O atom needs 8. Place lone pairs on the oxygen atom after all other atoms have satisfied their needs.

Counting Valence Electrons for CN-

Calculate the total number of valence electrons for the cyanide ion, CN-. C contributes 4 electrons, N contributes 5, and the extra negative charge contributes 1 for a total of 4 + 5 + 1 = 10 valence electrons.

Drawing the Lewis Structure for CN-

Link C and N by a triple bond, which accounts for 6 of the 10 valence electrons. Place the remaining 4 electrons as 2 lone pairs on the N atom. Add brackets with a negative sign to denote the extra charge.

Counting Valence Electrons for NO2-

Calculate total valence electrons for the nitrite ion, NO2-. N contributes 5 electrons, each O contributes 6, and the negative charge adds 1, totaling 5 + (2*6) + 1 = 18 valence electrons.

Drawing the Skeletal Structure for NO2-

Draw the N atom in the center with O atoms on the sides. Connect each O atom to the N atom with a single bond.

Completing the Octet and Distributing Electrons for NO2-

Place lone pairs on the O atoms to complete their octet. If insufficient electrons are left to complete the N atom octet, use a double bond between N and one O atom. Enclose the structure in brackets with a negative sign to represent the extra charge.

Counting Valence Electrons for ClO-

Count the total valence electrons for the hypochlorite ion, ClO-. Cl contributes 7 electrons, O contributes 6, and the negative charge adds 1 for a total of 7 + 6 + 1 = 14 valence electrons.

Drawing the Lewis Structure for ClO-

Connect Cl and O with a single bond. Complete the octets by placing 6 electrons (3 lone pairs) on the Cl atom and 6 electrons (3 lone pairs) on the O atom. Enclose the structure in brackets with a negative sign to show the charge.

Key concepts

Valence Electrons

Valence electrons play a crucial role in chemical bonding, as they are the outermost electrons of an atom and participate in forming bonds with other atoms. Understanding valence electrons is fundamental to drawing Lewis structures, which are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist.

The number of valence electrons for a given element is determined by its group number (column) in the periodic table. For instance, elements in Group 1 have 1 valence electron, while elements in Group 18 have 8 valence electrons, reflecting their noble gases' stable configuration.

Counting Valence Electrons

In the exercise provided, we first count valence electrons: hydrogen (H) has 1, carbon (C) has 4, oxygen (O) has 6, nitrogen (N) has 5, and chlorine (Cl) has 7. These values are key in predicting an element's reactivity and are utilized to ensure that the Lewis structure obeys the octet rule, which will be discussed in the next section.

Octet Rule

The octet rule is a chemical rule of thumb that states atoms tend to bond in such a way that they each have eight electrons in their valence shell, achieving a noble gas-like electron configuration. This provides stability to the atom in the form of a full octet.

There are exceptions to the octet rule, particularly for molecules containing elements that are capable of having fewer or more than eight electrons, such as hydrogen (which only requires two electrons to fill its valence shell) or sulfur and phosphorus (which can expand their octet in some compounds).

Applying the Octet Rule

In the step-by-step solutions for molecules like H3COCH3 (acetone), NO2-, and ions like CN- and ClO-, electrons are distributed to satisfy the octet rule. The central atoms in these structures are typically less electronegative and capable of forming multiple bonds, a pattern we exploit to ensure that each atom reaches an octet, as seen with the double bonds formed in some of these structures.

Electron Dot Structures

Electron dot structures, commonly called Lewis structures, are a way of representing molecular structure where dots represent valence electrons. These diagrams help us to visualize the bonding between atoms and the distribution of electrons in a molecule.

The dots in a Lewis structure are arranged to indicate paired and unpaired electrons. Paired electrons can either form covalent bonds between atoms or exist as lone pairs, which are pairs of electrons not involved in bonding but contribute to the shape and reactivity of the molecule.

Drawing Lewis Structures

In drawing Lewis structures, it's important to place the least electronegative atom at the center (except hydrogen, which always goes on the outside), and spread out the remaining atoms around it. Initially, single bonds are used to connect atoms; then electron pairs are added to complete the octet configuration. For ions, the overall charge is indicated with brackets around the structure. The steps to drawing Lewis structures are illustrated in the solutions provided for molecules like acetone (H3COCH3) and ions such as CN-, NO2-, and ClO-.