Problem 7 - Chemistry and Chemical Reactivity

Question

Draw a Lewis structure for each of the following molecules: (a) chlorodifluoromethane, CHCIF \(_{2}\) (C is the central atom) (b) acetic acid, \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) (basic structure pictured below) (IMAGE NOT COPY) (c) acetonitrile, \(\mathrm{CH}_{3} \mathrm{CN}\) (the framework is \(\mathrm{H}_{3} \mathrm{C}-\mathrm{C}-\mathrm{N}\) ) (d) allene, \(\mathrm{H}_{2} \mathrm{CCCH}_{2}\)

Step-by-Step Solution

Answer

Draw the structures by forming bonds and distributing electrons to satisfy the octet rule for all atoms.

Step 1: Understand the Concept of Lewis Structures

Lewis structures represent molecules showing all atoms, bonds, and lone pairs of electrons. They help predict the arrangement of atoms in a molecule.

Step 2: Determine Total Valence Electrons

For each molecule, calculate the total number of valence electrons. This is the sum of the valence electrons for all the atoms present in the molecule.

Step 3: Write the Skeletal Structure

Arrange the elements with the least electronegative atom in the center (except for hydrogen). Connect the atoms with single bonds to represent the basic framework of the molecule.

Step 4: Distribute Electrons to Outer Atoms

Distribute the remaining electrons, first assigning them to satisfy the octet rule for the outer atoms (except hydrogen, which needs only 2 electrons).

Step 5: Assign Remaining Electrons to the Central Atom

If any electrons are left, assign them to the central atom. If the central atom does not have a complete octet, try forming double or triple bonds to satisfy the octet rule.

Step 6: Draw Lewis Structure for CHCIF2

For CHCIF2, carbon in the center is bonded with H, Cl, and two F atoms. Distribute the remaining electrons to complete the octets for Cl and F atoms.

Step 7: Draw Lewis Structure for Acetic Acid

For CH3CO2H, the basic framework involves C bonded to three H on one end and a COOH group on the other. Ensure the C and O atoms satisfy the octet, with O in OH having lone pairs.

Step 8: Draw Lewis Structure for Acetonitrile

For CH3CN, connect C with three H atoms and the other C, then C to N with a triple bond to satisfy the octet, assigning remaining electrons as lone pairs on N.

Step 9: Draw Lewis Structure for Allene

For H2CCCH2, two terminal C are bonded to 2 H each and centrally to a middle C with double bonds, making sure all C atoms have a complete octet.

Key Concepts

Chemical Bonds

In chemistry, understanding chemical bonds is crucial to studying molecules. Chemical bonds are the connections between atoms in a molecule. They hold atoms together and determine the molecule's structure and properties. There are different types of chemical bonds, such as:

• Covalent Bonds: where atoms share pairs of electrons. This bond type is common in organic molecules, like those made up of carbon, hydrogen, and oxygen.
• Ionic Bonds: formed when one atom gives up one or more electrons to another atom. This typically occurs between metals and non-metals.
• Metallic Bonds: where electrons are shared over many nuclei, allowing them to conduct electricity.

In molecules, single, double, and triple bonds can form depending on how many electron pairs are shared. For instance, a single bond shares one pair, a double bond shares two pairs, and a triple bond shares three pairs. In the exercise above, examples like acetonitrile and allene involve double and triple bonds to maintain the proper structure. Integrating these bonds into Lewis structures helps us visualize these connections.

Valence Electrons

Valence electrons are the outermost electrons in an atom. They play a key role in forming chemical bonds. The number of valence electrons determines an atom's chemical reactivity and bonding potential.
For instance:

• Carbon, with 4 valence electrons, can form four bonds.
• Oxygen, having 6 valence electrons, usually forms two bonds.
• Nitrogen, possessing 5 valence electrons, typically forms three bonds.

In Lewis structures, valence electrons are depicted as dots around the atoms. When constructing Lewis structures, first sum up the total valence electrons. Then distribute these electrons to form pairs and to ensure the atoms in a molecule reach a stable state.
This method aids in predicting how atoms bond in molecules. For example, in the exercise, acetic acid requires us to consider how many valence electrons are needed for each atom to complete its bonding and attain stability.

Octet Rule

The octet rule is a guiding principle in chemistry that suggests atoms tend to form compounds in ways that provide them with eight electrons in their outer shell, leading to increased stability. This rule is followed largely by main group elements.
Some key points include:

• Atoms like carbon, oxygen, and nitrogen achieve a full octet through bond formation.
• Exceptions include hydrogen, which is stable with just two electrons, and certain elements that may have expanded octets.

While drawing Lewis structures, the octet rule guides how electrons are distributed among atoms. For instance, when constructing the Lewis structure for acetonitrile, the triple bond between carbon and nitrogen helps them both reach a complete octet.
If any central atom lacks a full octet after distributing available electrons, forming double or triple bonds is necessary. This ensures that stability is achieved for the molecule. Understanding the octet rule is crucial for interpreting and predicting the structure of chemical compounds.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. Understanding this concept helps predict not just the shape but also the reactivity and properties of the molecule.
VSEPR Theory (Valence Shell Electron Pair Repulsion theory) helps predict the geometry by assuming electron pairs around the central atom repel each other to maximize distances.

• Linear: like CO2, where atoms are in a straight line.
• Tetrahedral: as in CH4, with atoms forming a shape like a pyramid with a triangular base.
• Trigonal Planar: where three atoms are spread out in one plane around a central atom.

Understanding molecular geometry helps in visualizing the spatial arrangement of atoms, as seen with the molecules in the exercise. For example, the central carbon in chlorodifluoromethane bonds with other atoms in a way that respects geometric constraints and keeps electron pairs and bonds at an optimal angle.
This concept is vital for comprehending how molecules function in chemical reactions and their interactions with other molecules.