Question

Draw Lewis structures for the following molecules and ions: (a) \(\mathrm{OF}_{2},(\mathrm{~b}) \mathrm{N}_{2} \mathrm{~F}_{2} (\mathrm{c}) \mathrm{Si}_{2} \mathrm{H}_{6},\) (d) \(\mathrm{OH}^{-},\) (e) \(\mathrm{CH}_{2} \mathrm{ClCOO}^{-},\) (f) \(\mathrm{CH}_{3} \mathrm{NH}_{3}^{+}\).

Short answer

Lewis structures show electrons in molecule/ions; identify valence electrons, arrange atoms, add bonds/lone pairs, ensure octet/charge balance.

Step-by-step solution

Draw Lewis Structure of OF₂

For the molecule OF₂, we start by counting the total valence electrons. Oxygen (O) has 6 valence electrons, and each fluorine (F) atom has 7 valence electrons. Total electrons: 6 + 7*2 = 20. Place O in the center, and connect with single bonds to each F. Distribute remaining electrons to satisfy the octet for F first, then O. The structure is O with 2 lone pairs, each F with 3 lone pairs.

Draw Lewis Structure of N₂F₂

For N₂F₂, count the valence electrons. N (Nitrogen) has 5, F has 7. Total electrons: 5*2 + 7*2 = 24. Place the N atoms in the center with a single bond between them, then single bond each to an F atom. Distribute electrons to satisfy octets first around F, then N, using lone pairs.

Draw Lewis Structure of Si₂H₆

For Si₂H₆, Si (Silicon) has 4 valence electrons, and H (Hydrogen) has 1. Total electrons: 4*2 + 1*6 = 14. Si atoms go in the center, connected by a single bond. Each Si is bonded to three H atoms. All bonds are single, with no lone pairs remaining.

Draw Lewis Structure of OH⁻

OH⁻ has 1 oxygen and 1 hydrogen. O has 6 valence electrons, H has 1. Add an electron for the negative charge: 6 + 1 + 1 = 8. Place O and H with a single bond, with O holding 3 lone pairs. Distribute the electrons to complete the octet on O, then show the negative charge.

Draw Lewis Structure of CH₂ClCOO⁻

This is an organic ion. Total valence electrons: C (4) * 2 + H (1) * 2 + Cl (7) + O (6) * 2 + 1 (negative charge) = 42. Arrange carbons in backbone, place Cl on the first C, COO⁻ group on the second C. Single bonds throughout, with necessary lone pairs added to complete octets, noting the negative charge on one of the oxygens in COO⁻.

Draw Lewis Structure of CH₃NH₃⁺

For CH₃NH₃⁺, total electrons: C (4) + H (1) * 3 + N (5) + H (1)*3 - 1 (positive charge) = 13. Arranged with C bonded to N, both bonded to hydrogens. Display C with 3 H's, and N with 3 H's, N is the central atom after C, forming single bonds, with the positive charge indicated on N.

Key concepts

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom. These are the electrons involved in chemical bonding.
For example, in the organic ion \( ext{CH}_2 ext{ClCOO}^-\), you need to tally up all the valence electrons to draw an accurate Lewis structure.

• Carbon (C) typically has 4 valence electrons.
• Hydrogen (H) has 1 valence electron.
• Chlorine (Cl) has 7 valence electrons.
• Oxygen (O) has 6 valence electrons.

Electrons equal to the negative charge of ions are added, while electrons are subtracted for positive charges.
Understanding how valence electrons contribute to the structure is crucial for predicting how atoms bond and form complex molecules.

Octet Rule

The octet rule is a chemical rule of thumb that reflects the idea that atoms tend to combine in a way that allows them to have eight electrons in their valence shell.
This rule is especially applicable to main group elements, which strive to reach the same electron configuration as noble gases.
For instance, the \( ext{OH}^-\) ion must fulfill the octet rule for oxygen, meaning after sharing electrons with hydrogen and considering the extra electron from its negative charge, oxygen should have eight electrons surrounding it.
Exceptions to this rule exist, such as with hydrogen and helium, which are stable with only two valence electrons. In our example exercises, following the octet rule ensures each atom (except H) has a complete set of valence electrons, leading to stable molecular structures.

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms within a molecule.
Understanding this concept helps predict the shape and bond angles around a central atom.
Using \( ext{Si}_2 ext{H}_6\), knowing the shape is helpful. The molecule's geometry is determined by the number of bonds and lone pairs around the silicon atoms.

• No lone pairs mean the geometric arrangement is based solely on the connected atoms.
• In \( ext{Si}_2 ext{H}_6\), each silicon atom is tetrahedrally coordinated with hydrogen atoms making a stable structure.

Considering molecular geometry aids in understanding the physical and chemical properties of molecules.

Bonding

Bonding refers to the interaction between atoms that leads to the formation of molecules.
In Lewis structures, bonds are often represented by lines linking atoms.
There are primarily three types of chemical bonds:

• Ionic bonds are due to the attraction between oppositely charged ions, as seen in ionic structures.
• Covalent bonds involve the sharing of electron pairs between atoms, prevalent in molecules like \( ext{N}_2 ext{F}_2\).
• Metallic bonds, which are not covered in these examples, occur among metal atoms.

In the provided molecules, covalent bonding plays a critical role. Identifying whether atoms share one or multiple pairs of electrons helps visualize and predict molecule behavior.

Ion Structures

Ion structures involve charged atomic or molecular assemblies. Ions can be anions (negatively charged) or cations (positively charged).
The exercise showcases ion structures like \( ext{CH}_3 ext{NH}_3^+\) and \( ext{CH}_2 ext{ClCOO}^-\).

• In \( ext{CH}_3 ext{NH}_3^+\), the amine group carries a positive charge, illustrating a cation structure.
• For \( ext{CH}_2 ext{ClCOO}^-\), additional electrons contributing to a negative charge on oxygen illustrate an anion.

Understanding these ion structures is crucial as they significantly affect the molecule's reactivity and interactions. When constructing Lewis structures, accounting for the correct charge is necessary to ensure accuracy and predict the compound's potential chemical interactions.