Problem 48 - Chemistry The Central Science In Si Units

Question

Write Lewis structures for the following: \((\mathbf{a}) \mathrm{H}_{2} \mathrm{CO}\) (both \(\mathrm{H}\) atoms are bonded to \(\mathrm{C}),(\mathbf{b}) \mathrm{H}_{2} \mathrm{O}_{2},\) (c) \(\mathrm{C}_{2} \mathrm{~F}_{6}\) (contains a \(\mathrm{C}-\mathrm{C}\) bond \(),(\mathbf{d}) \mathrm{AsO}_{3}{\underline{\phantom{xx}}}^{3-} (\mathbf{e}) \mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{H}\) is bonded to \(\mathrm{O})\) (f) \(\mathrm{NH}_{2} \mathrm{Cl}\).

Step-by-Step Solution

Answer

Draw Lewis structures by arranging atoms, forming bonds, and completing octets with lone pairs based on valence electrons.

Step 1: Draw Lewis Structure for H₂CO

1. Count total valence electrons: H (1) x 2 + C (4) + O (6) = 12 electrons. 2. Arrange atoms: Central atom C, with two H atoms and one O atom attached. 3. Distribute electrons starting with bonds: Use 8 electrons forming two C-H bonds and a C-O double bond. 4. Place remaining electrons as lone pairs on O to complete the octet.

Step 2: Draw Lewis Structure for H₂O₂

1. Calculate total valence electrons: H (1) x 2 + O (6) x 2 = 14 electrons. 2. Arrange atoms: H-O-O-H. 3. Create single bonds between each O and H, and a single bond between the two O atoms. 4. Use remaining electrons to form lone pairs on the O atoms to complete their octets.

Step 3: Draw Lewis Structure for C₂F₆

1. Total valence electrons: C (4) x 2 + F (7) x 6 = 50 electrons. 2. Arrange structure: C-C and each C bonded to three F atoms. 3. Use single bonds to attach the F atoms to each C. 4. Each C and F should complete its octet with lone pairs if needed.

Step 4: Draw Lewis Structure for AsO₃³⁻

1. Total valence electrons: As (5) + O (6) x 3 + 3 extra electrons from charge = 26 electrons. 2. Center As with three O atoms around it. 3. Form single bonds with each O to As. 4. Distribute remaining electrons to complete octets on O atoms, accounting for the charge with extra electrons forming double bonds if necessary.

Step 5: Draw Lewis Structure for H₂SO₃

1. Total valence electrons: H (1) x 2 + S (6) + O (6) x 3 = 26 electrons. 2. Arrange atoms: Central S bonded to three O atoms, with one O bonded to two H atoms. 3. Create single bonds, adjusting to form a double bond for one S=O if necessary for octet. 4. Complete octets for O atoms and distribute remaining electrons as lone pairs on S.

Step 6: Draw Lewis Structure for NH₂Cl

1. Total valence electrons: N (5) + H (1) x 2 + Cl (7) = 14 electrons. 2. Arrange atoms: N central with two H and one Cl attached. 3. Form single bonds between N and each H and Cl. 4. Distribute remaining electrons as lone pairs on N to complete its octet.

Key Concepts

Valence Electrons

Before drawing a Lewis structure, it is important to count the total number of valence electrons available for bonding in a molecule. Valence electrons are the outermost electrons of an atom that participate in bond formation. Knowing the number of valence electrons helps in predicting how atoms combine to form molecules. For instance:

• Hydrogen () has 1 valence electron.
• Carbon () has 4 valence electrons.
• Oxygen () has 6 valence electrons.
• Fluorine () has 7 valence electrons.

To calculate the total number of valence electrons in a molecule, sum up the valence electrons of each atom present in the molecule. For charged species like ions, add or subtract electrons based on the charge. Consider  containing a C-C bond. Carbon contributes 4 valence electrons each, and each fluorine contributes 7, totaling 50 valence electrons for . Adding these correctly is crucial, as it determines how atoms are linked and the stability of the molecule.

Octet Rule

A cornerstone of understanding molecular bonding is the octet rule. This rule states that atoms tend to form bonds until they are surrounded by eight valence electrons, resembling the electron configuration of a noble gas. This preference for eight electrons guides how atoms bond in molecules like . Typically, single bonds contribute two electrons to each bonded atom. For example, in , the carbon atoms form single bonds with fluorine and each other.

• Single bonds: Each bond involves a pair of electrons.
• Double bonds: Atoms may form double bonds (4 electrons) when single bonds alone do not satisfy the octet. This can be seen in the  molecule, where oxygen forms a double bond with carbon.
• Lone pairs: Additional unused valence electrons are placed as lone pairs, counted to reach the octet for elements like oxygen in .

Understanding the octet rule helps in constructing a correct and stable Lewis structure, ensuring that each atom (except hydrogen, which follows the duet rule) fills its outer shell appropriately.

Molecular Geometry

The spatial arrangement of atoms in a molecule is referred to as its molecular geometry. Lewis structures provide an initial step by showing the connections between atoms, which helps to predict the 3D shape of the molecule by applying VSEPR theory (Valence Shell Electron Pair Repulsion). VSEPR theory suggests that molecules take on a shape that minimizes the repulsion between valence electron pairs around the central atom.
Different molecules adopt different shapes:

• Linear: Atoms are arranged in a straight line.
• Tetrahedral: Atoms form a 3D shape similar to a pyramid with a base.
• Trigonal planar: Atoms are arranged in a flat, triangular pattern around a central atom.

For example, in  (water), the molecule adopts a bent geometry due to the two lone pairs on oxygen, which push the hydrogen-oxygen bonds closer together. Understanding molecular geometry is crucial for predicting the physical and chemical properties of compounds, such as how they interact with other molecules and their reactivity.