Question

Two different molecules have the formula \(\mathrm{C}_{2} \mathrm{H}_{6} \mathrm{O} .\) One of the molecules has the oxygen atom bonded to both carbon atoms. The other molecule has the oxygen atom bonded to only one carbon atom while both carbon atoms are bonded to each other. Write Lewis structures for both of these compounds.

Short answer

Question: Draw the Lewis structures for two compounds with the chemical formula \(C_{2}H_{6}O\). In one compound, the oxygen atom is bonded to both carbon atoms, while in the other compound, the oxygen atom is bonded to only one carbon atom, and the two carbon atoms are bonded to each other. Answer: Compound 1 Lewis structure: :O: // \\ C - H H - C Compound 2 Lewis structure: H : O : C // | H C C

Step-by-step solution

Determine the total number of valence electrons

For each element in the chemical formula \(\mathrm{C}_{2}\mathrm{H}_{6}\mathrm{O}\), we need to determine the number of valence electrons: - Carbon (C) has 4 valence electrons. - Hydrogen (H) has 1 valence electron. - Oxygen (O) has 6 valence electrons. Now, we can calculate the total number of valence electrons in the molecule: 2 Carbon atoms × 4 valence electrons + 6 Hydrogen atoms × 1 valence electron + 1 Oxygen atom × 6 valence electrons = 8 + 6 + 6 = 20 valence electrons

Draw the Lewis structure for the first compound

In the first compound, the oxygen atom is bonded to both carbon atoms. So we can start by connecting the two carbon atoms and the oxygen atom: O / \ C C Now, we need to add the hydrogen atoms. Each carbon atom needs to make a total of 4 bonds, and currently, they are each making 2 bonds (with oxygen and the other carbon atom). So we need to add 2 hydrogen atoms to each carbon atom: O / \ C - H H - C Finally, we need to add lone pairs to the oxygen atom. Oxygen needs 8 valence electrons, but it currently has only 4 (from the double bond to the carbon atoms). This means we need to add 2 more lone pairs to the oxygen atom: :O: // \\ C - H H - C The resulting Lewis structure for the first compound is drawn above.

Draw the Lewis structure for the second compound

In the second compound, the oxygen atom is bonded to only one carbon atom, while the two carbon atoms are bonded to each other. We can start by drawing a single bond between the two carbon atoms, and connecting the oxygen atom to one of the carbon atoms: O C // | C C To complete the carbon-oxygen bond, we need to add two lone pairs to the oxygen atom: :O: C // | C C Now, we need to add hydrogen atoms to the carbon atoms. Each carbon atom needs to make a total of 4 bonds. The left carbon atom has 3 bonds already, so we need to add 1 hydrogen atom to it. The right carbon atom has only 1 bond, so we need to add 3 hydrogen atoms to it: H : O : C // | H C C H We can now add in the remaining lone pairs for nonbonding electrons if desired. The resulting Lewis structure for the second compound is drawn above.

Key concepts

Valence Electrons

Valence electrons play a pivotal role in understanding the chemical behavior of atoms. These electrons are located in the outermost shell of an atom and are responsible for the bonding with other atoms to form molecules. Each atom has a characteristic number of valence electrons, which can be easily determined by referring to the periodic table: for instance, Carbon (C) has 4 valence electrons, Hydrogen (H) has 1, and Oxygen (O) has 6.

In the exercise, the calculation of the total number of valence electrons is essential as it provides the foundation for drawing the correct Lewis structures. For the compound \(\mathrm{C}_{2} \mathrm{H}_{6}\mathrm{O}\), there are 20 valence electrons to consider. Keep in mind, while Hydrogen can only form one bond, Carbon and Oxygen can form four and two bonds respectively, typically aiming to fulfill the octet rule - having eight electrons in their valence shell to gain stability, mimicking the noble gas configuration.

When creating the Lewis structures to represent molecules, the first step is always to count these electrons, ensuring the structure accounts for all of them, thus guaranteeing a stable configuration.

Chemical Bonding

Chemical bonding is how atoms connect to form molecules, and understanding how to display these connections is crucial in drawing Lewis structures. It's important to note that atoms bond in such a way that each achieves a full valence shell, which often corresponds to eight valence electrons, a rule of thumb known as the 'octet rule'. However, Hydrogen is an exception, as it fulfills its valence shell with two electrons.

In the provided solutions, the distinct ways in which Carbon and Oxygen can bond are demonstrated. The concept of single, double, and triple bonds comes into play, where a single bond involves two electrons (one pair), a double bond involves four electrons (two pairs), and a triple bond involves six electrons (three pairs). The Lewis structures depict these bonds as lines between atoms – single lines for single bonds, double lines for double bonds, and so forth.

The exercise depicts two isomers that vary based on their chemical bonding: the first with an Oxygen atom double-bonded to both Carbon atoms, and the second with the Oxygen atom single-bonded to just one Carbon atom. Understanding these bonding variations is key to grasping molecular diversity and reactivity.

Molecular Geometry

Molecular geometry is about the three-dimensional arrangement of atoms within a molecule. The geometry is determined by the number of bonds and lone pairs around a central atom. VSEPR (Valence Shell Electron Pair Repulsion) theory helps us understand these shapes by predicting how bond pairs and lone pairs will arrange themselves around a central atom to minimize repulsion.

In the case of the molecules \(\mathrm{C}_{2} \mathrm{H}_{6}\mathrm{O}\) that we've examined, the central atom differs for the two structures. In the first structure, the Oxygen has a linear arrangement due to the double bonds to Carbon, while in the second structure, the geometry around the Carbons to which Hydrogens are attached is tetrahedral. As Carbon is the central atom in the second structure, it forms a bent geometry with Oxygen due to the presence of lone pairs on Oxygen.

Molecular geometry is important not only for understanding the structure of a molecule but also for predicting its reactivity, polarity, phase of matter, color, magnetism, as well as its biological activity. Therefore, when drawing Lewis structures, it is crucial not only to get the connections and valence electrons right but also to remember that these structures translate into a three-dimensional world, where the shape of molecules matters immensely for their properties and functions.