Question

Write plausible Lewis structures for the following molecules that contain only single covalent bonds. (a) \(\mathrm{FCl} ;\) (b) \(\mathrm{I}_{2} ;\) (c) \(\mathrm{SF}_{2} ;\) (d) \(\mathrm{NF}_{3} ;\) (e) \(\mathrm{H}_{2} \mathrm{Te}\).

Short answer

The plausible Lewis structures for the molecules with single covalent bonds are: F-Cl for FCl, I-I for I2, F-S-F for SF2, and H-Te-H for H2Te. The structure for NF3 is a bit more complicated: F is linked to N through a single bond, and this N is then linked to two other Fs through single bonds (forming a Y-like shape).

Step-by-step solution

Draw Lewis structure for FCl

F and Cl each has seven valence electrons. They can share one electron forming a single covalent bond. The remaining six electrons on each atom will be shown as lone pairs. The Lewis structure is: F-Cl.

Draw Lewis structure for I2

Both I atoms have seven valence electrons. They share a pair of electrons forming a single covalent bond with each other and the remaining six electrons on each atom are lone pairs. The Lewis structure is: I-I

Draw Lewis structure for SF2

S has six and each F has seven valence electrons. S shares one electron with each F atom forming single covalent bonds. The remaining four electrons on S are lone pairs and the six electrons on each F atom are lone pairs. The Lewis structure is: F-S-F

Draw Lewis structure for NF3

N has five and each F has seven valence electrons. N shares one electron with each F atom forming single covalent bonds. The remaining two electrons on N are a lone pair and the six electrons on each F atom are lone pairs. The Lewis structure is: F\n| \nN - F\n| \nF

Draw Lewis structure for H2Te

Each H has one and Te has six valence electrons. Te shares one electron with each H atom forming single covalent bonds. The remaining four electrons on Te are lone pairs. The Lewis structure is: H-Te-H

Key concepts

Understanding Covalent Bonds

Covalent bonds are a type of chemical bond where two atoms share one or more pairs of electrons. This sharing allows each atom in the bond to reach a more stable electron configuration, often achieving a full outer electron shell similar to that of the nearest noble gas. Covalent bonds are essential in the formation of molecules and are the most common form of chemical bonding.

• Covalent bonds share electrons, unlike ionic bonds which transfer them.
• They often occur between non-metal atoms.
• Single covalent bonds involve the sharing of one pair of electrons.

For example, in the molecule \(\)\(\mathrm{FCl}\), fluorine (F) and chlorine (Cl) each contribute one unpaired electron to form a single covalent bond. This allows both atoms to achieve stable electron configurations by sharing the electrons.

Exploring Valence Electrons

Valence electrons are the outermost electrons of an atom that are involved in forming bonds with other atoms. These electrons determine the chemical reactivity and bonding capabilities of an element. In the periodic table, the number of valence electrons generally increases from left to right across a period.Understanding valence electrons helps predict how atoms will bond.

• Atoms with similar numbers of valence electrons often form similar types of bonds.
• Most stable electron configurations are achieved by having 8 valence electrons, known as the octet rule.
• Hydrogen is an exception, aiming for 2 valence electrons.

Looking at the molecule \(\mathrm{NF_3}\), nitrogen has five valence electrons and each fluorine has seven. By sharing electrons through covalent bonds, nitrogen can achieve a stable configuration with its neighbors.

The Role of Lone Pairs

Lone pairs refer to valence electrons that are not shared with another atom in a covalent bond. These pairs of electrons are localized on one atom and play a significant role in the shape and properties of molecules. Lone pairs can influence molecular geometry, polarity, and reactivity.When considering lone pairs:

• They can repel bonding pairs of electrons, affecting the molecular shape.
• Lone pairs are responsible for most of the interesting chemistry involving electron-rich regions.
• A molecule's dipole moment and overall polarity can be drastically changed by these non-bonding electrons.

In \(\mathrm{SF_2}\), sulfur has lone pairs on it in addition to bonded electrons. These lone pairs push the fluorine atoms closer together, resulting in a bent shape for the molecule. This, in turn, affects the molecule's properties and how it interacts with other substances.