Question

Draw three Lewis structures for compounds with the formula \(\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{~F}_{2}\). Indicate which of the compounds are polar.

Short answer

Two of the compounds are polar; one is non-polar.

Step-by-step solution

Understand the Problem

We need to draw three different Lewis structures for the molecular formula \(\text{C}_2\text{H}_2\text{F}_2\). Each structure should represent a unique arrangement of atoms, and we need to identify whether the structures are polar.

Count Valence Electrons

Carbon has 4 valence electrons, hydrogen has 1, and fluorine has 7. For \(\text{C}_2\text{H}_2\text{F}_2\): \(2 \times 4 \text{ (C)} + 2 \times 1 \text{ (H)} + 2 \times 7 \text{ (F)} = 20\) total valence electrons.

Draw the Skeleton Structures

Typically, carbon atoms are central: place \(\text{H}\) and \(\text{F}\) around \(\text{C}_2\). Possible structures: \(\text{F}-\text{C}-\text{C}-\text{F}\), \(\text{H}-\text{C}-\text{C}-\text{H}\), and combinations like \(\text{F}-\text{C}-\text{C}-\text{H}\).

Complete Octets

For each configuration, ensure every carbon and fluorine has an octet. Distribute remaining electrons accordingly around the \(\text{C}-\text{C}\) bond and to hydrogen atoms.

Identify Polar Molecules

A molecule is polar if it has a net dipole moment. The \(\text{F}-\text{C}-\text{C}-\text{F}\) structure is non-polar due to symmetric dipoles canceling out. The asymmetric \(\text{F}-\text{C}-\text{C}-\text{H}\) structure is polar, as the dipoles do not cancel out.

Key concepts

Molecular Polarity

Molecular polarity refers to how the electrons are distributed in a molecule. This distribution affects whether a molecule has a net dipole moment, which is a measure of the overall charge separation. If a molecule has polar covalent bonds that do not cancel each other out, the molecule is polar.
For example, in \(\text{C}_2\text{H}_2\text{F}_2\), the arrangement of atoms can lead to different polarities depending on the symmetry:

• The symmetric structure \(\text{F}-\text{C}-\text{C}-\text{F}\) is non-polar because its dipoles cancel out.
• An asymmetric arrangement like \(\text{F}-\text{C}-\text{C}-\text{H}\) results in a polar molecule because the dipoles do not cancel.

This concept helps us understand the behavior of molecules in different environments, as polar molecules tend to interact differently with other molecules compared to non-polar ones.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in chemical bonding. They determine an atom's ability to bond with others by being shared or transferred.
For the molecule \(\text{C}_2\text{H}_2\text{F}_2\):

• Carbon contributes 4 valence electrons per atom.
• Hydrogen contributes 1 valence electron each.
• Fluorine contributes 7 valence electrons each.

This adds up to a total of 20 valence electrons for the molecule. Understanding the distribution of valence electrons is key to drawing accurate Lewis structures and predicting the molecule’s reactivity and properties.

Octet Rule

The octet rule is a guideline in chemistry that suggests atoms are most stable when they have eight electrons in their valence shell. This rule is especially applicable to carbon, nitrogen, oxygen, and the halogens.
In drawing Lewis structures for \(\text{C}_2\text{H}_2\text{F}_2\), each carbon and fluorine should have an octet:

• Hydrogen, however, is an exception to the octet rule and is stable with two electrons.
• By arranging the electrons to satisfy each atom's octet, one can predict the molecule's stability and structure.

Understanding and applying the octet rule enables us to construct plausible molecular structures and assess their chemical behavior.

Chemical Bonding

Chemical bonding involves the joining of atoms to form molecules. Bonds can be ionic, covalent, or metallic, but in \(\text{C}_2\text{H}_2\text{F}_2\), the focus is on covalent bonds.
Covalent bonds arise from the sharing of valence electrons between atoms. For this molecule:

• Carbon-carbon single or double bonds help create the backbone.
• The bonds with hydrogen and fluorine involve sharing electrons to satisfy the octet rule.

Understanding the type and nature of these bonds helps us comprehend the molecule’s properties and how it might interact with others. Covalent bonds create stable structures with specific angles and lengths, influencing the molecule's overall shape and behavior.