Question

Consider two different organic compounds, each with the formula \(\mathrm{C}_{2} \mathrm{H}_{6} \mathrm{O}\). One of these compounds is a liquid at room conditions and the other is a gas. Write Lewis structures consistent with this observation, and explain your answer. (Hint: The oxygen atom in both structures satisfies the octet rule with two bonds and two lone pairs.)

Short answer

The two different organic compounds with the molecular formula \(C_2H_6O\) are ethanol and dimethyl ether. Ethanol has a structure like this: H H O H | | | | C-C-H | H Ethanol is a liquid at room temperature due to the presence of a hydroxyl (-OH) group, which allows for strong hydrogen bonding. Dimethyl ether has a structure like this: H H O H | | | | C C | | H H Dimethyl ether is a gas at room temperature because it has an ether (-O-) group, which can only form weak van der Waals forces. This difference in intermolecular forces results in different physical states for ethanol and dimethyl ether at room temperature.

Step-by-step solution

Applying the octet rule to the oxygen atom

According to the hint given in the problem, oxygen in both compounds has two bonds and two lone pairs. Therefore, we use this information in drafting the structures, as well as adhering to the molecular formula of \(C_2H_6O\).

Investigate different structure possibilities for organic compounds with the formula \(C_2H_6O\)

We start by considering different ways in which the atoms can bond with one another while maintaining the octet rule for oxygen and other atoms in the molecule. Note that the carbon atoms can either be one after the other (ethanol) or joined in such a fashion that one carbon connects three atoms (dimethyl ether).

Write Lewis structures for the possible compounds

Now we write the Lewis structures for both ethanol and dimethyl ether, as follows: Ethanol: H H O H | | | | C-C-H | H Dimethyl ether: H H O H | | | | C C | | H H

Identify which compound is a liquid and which is a gas at room temperature

The difference between these two compounds lies in the presence of a hydroxyl (-OH) group in ethanol and an ether (-O-) group in dimethyl ether. Due to the presence of the hydroxyl group, ethanol can form strong hydrogen bonds, which increase the intermolecular forces between the molecules. This makes ethanol a liquid at room temperature. In contrast, dimethyl ether molecules can only form weak van der Waals forces due to the presence of an ether group. As a result, dimethyl ether has a lower boiling point than ethanol and is a gas at room temperature. In conclusion, ethanol is the liquid compound with formula \(C_2H_6O\), while dimethyl ether is the gaseous compound with the same formula. The reasoning behind this observation is the different types of intermolecular forces present in these compounds, which are hydrogen bonding in ethanol, and van der Waals forces in dimethyl ether.

Key concepts

The Octet Rule and Its Application in Organic Compounds

Understanding the octet rule is fundamental in predicting and drawing the structures of organic compounds. This rule states that atoms tend to form bonds in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas. In the context of organic chemistry, carbon, nitrogen, oxygen, and the halogens frequently follow the octet rule. When drawing Lewis structures for organic compounds such as ethanol (\(C_2H_6O\)), we represent the valence electrons as dots around the elements and ensure that, except for hydrogen, each atom is surrounded by eight electrons.

Applying the octet rule, we can predict that the oxygen in the compounds will have two bonds and two lone pairs. This arrangement completes the oxygen's octet, where lone pairs represent non-bonding electrons. By following this guideline, we ensure stability in the drawn structures, as atoms lower their potential energy by achieving full valence shells. Knowing the valence of common atoms like carbon and hydrogen further directs our building of the molecular structure.

Exercise Improvement Advice

To enhance understanding, students should practice recognizing valence electrons and applying the octet rule to different organic molecules beyond the given example. It also helps to remember that there are exceptions to the rule, such as in the case of hydrogen (which follows the 'duet rule', seeking only two electrons), or in molecules with odd numbers of electrons, or those containing elements such as sulfur and phosphorus which can have expanded octets.

Intermolecular Forces in Organic Chemistry

Intermolecular forces, the attractions between molecules, play a crucial role in determining the physical properties of substances, including their boiling and melting points. In organic compounds, the primary intermolecular forces include London dispersion forces, dipole-dipole interactions, and hydrogen bonding. The strength and type of these forces depend on the nature of the molecules involved.

The London dispersion force is a temporary attractive force that results from the movement of electrons in molecules, creating instantaneous dipoles. These forces are generally weak and are significant in nonpolar compounds. Dipole-dipole interactions occur between polar molecules, where the positive end of one molecule attracts the negative end of another. These are stronger than dispersion forces but weaker than hydrogen bonds.

Exercise Improvement Advice

Students can benefit from visualizing these forces through molecular models and comparing how different structural features like molecular size and polarity affect the strength of the intermolecular forces. This can be linked to observable properties like volatility and solubility. Examples that contrast molecules with and without polar groups or hydrogen bonding capacity can reinforce the concept.

Hydrogen Bonding and Its Significance

Among intermolecular forces, hydrogen bonding is a particularly strong type of dipole-dipole interaction, but it's only present in molecules where hydrogen is covalently bonded to highly electronegative atoms like fluorine, oxygen, or nitrogen. The partially positive hydrogen atom of one molecule is attracted to the partially negative fluorine, oxygen, or nitrogen of another molecule, forming a hydrogen bond.

In the given organic compounds, ethanol exhibits this type of intermolecular force due to the presence of a hydroxyl (-OH) group. This hydrogen bonding significantly increases ethanol's boiling point, resulting in it being a liquid at room temperature. On the other hand, dimethyl ether lacks this -OH group, cannot engage in hydrogen bonding, and therefore, only displays weaker intermolecular forces like van der Waals forces, existing as a gas at room temperature.

Exercise Improvement Advice

To solidify their understanding, students should practice identifying potential sites for hydrogen bonding in molecular structures and relating this to the compound's physical properties. Comparing molecules with similar sizes but differing in hydrogen bonding capacity can illustrate how this type of intermolecular force affects the properties of a substance. Additionally, case studies of water's unique properties due to hydrogen bonding provide practical insights into the concept's importance.