Question

Write the Levis structure for each molecule. MISSED THIST Read Section 10.7; Watch KCV 10.7, IWE 10.4 a. \(\mathrm{SF}_{2}\) b. \(\mathrm{SiH}_{4}\) c. HCOOH (both O bonded to C) d. \(\mathrm{CH}_{3} \mathrm{SH}\) (C and \(\mathrm{S}\) central)

Short answer

The Lewis structure for \textrm{SF}_2 shows S with two bonds to F and two lone pairs. \textrm{SiH}_4 has Si in the center with 4 H atoms attached. HCOOH has C double-bonded to one O and single-bonded to another O, which is bonded to H; C also has a bond to H. \textrm{CH}_3\textrm{SH} has a central C attached to three Hs and one S, which is bonded to another H and has two lone pairs.

Step-by-step solution

Drawing the Lewis Structure for \textrm{SF}_2

First, determine the total number of valence electrons. Sulfur (S) has 6 and each Fluorine (F) has 7, for a total of 20. Next, place Sulfur in the center and arrange the 2 Fluorines around it. Connect the S to each F with a single bond. Sulfur will have two lone pairs to complete its octet, while Fluorine atoms will each have three lone pairs.

Drawing the Lewis Structure for \textrm{SiH}_4

Silicon (Si) is in Group 14 and has 4 valence electrons, and each Hydrogen (H) has 1, for a total of 8. Silicon will be the central atom with no lone pairs, as it follows the octet rule with 4 bonds. Connect each of the 4 Hydrogen atoms to Silicon with a single bond. Each H atom has no lone pairs as they are satisfied with 2 electrons.

Drawing the Lewis Structure for HCOOH (Formic Acid)

Carbon (C) is the central atom with 4 valence electrons, Hydrogen (H) has 1, each Oxygen (O) has 6, for a total of 18. Connect the C to one O by a double bond and another O by a single bond. Then, connect the O with a single bond to an H atom, and the C to another H atom. The O with a single bond will have three lone pairs and the one with the double bond will have two lone pairs. Carbon has no lone pairs.

Drawing the Lewis Structure for \textrm{CH}_3\textrm{SH}

The central atoms are Carbon (C) and Sulfur (S). C has 4 valence electrons, S has 6, Hydrogen (H) has 1, for a total of 14. Connect C to three H atoms and S with single bonds. Then, connect C to S with a single bond. Sulfur connects to an H atom and has two lone pairs to fulfill its octet. Carbon has no lone pairs.

Key concepts

Chemical Bonding

Chemical bonding is at the heart of understanding molecules and compounds. It explains how atoms attach themselves to each other to form more complex structures. In essence, atoms bond together to become more stable, often by achieving a complete outer shell of electrons similar to the noble gases, which are inherently stable due to their full valence electron shells.

There are three primary types of chemical bonds:

• \(Ionic bonds\) occur when one atom donates an electron to another, creating ions that attract each other due to opposite charges.
• \(Covalent bonds\) involve the sharing of electron pairs between atoms. These can be single, double, or triple bonds, depending on how many electron pairs are shared.
• \(Metallic bonds\) are found in metals where electrons are free to move throughout the structure, giving rise to properties like conductivity.

In the given exercise, we observe covalent bonding where nonmetals share electrons to satisfy the octet rule, which says that atoms are most stable with eight electrons in their valence shell (with the exception of hydrogen, which aims for two). For instance, in the molecule \(SF_2\), sulfur and fluorine share electrons, forming single covalent bonds.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a pivotal role in chemical bonding and the physical properties of substances. Since they are the farthest from the nucleus, they are the electrons most involved in chemical reactions. In the Lewis structures shown in the exercise, valence electrons are represented as dots around the chemical symbols of the elements.

The number of valence electrons an atom has determines its chemical reactivity and bonding behavior. For example, in the exercise step for \(SiH_4\), we find that silicon has four valence electrons and thus can form four single covalent bonds to satisfy the octet rule. This step-by-step approach - first determining the number of valence electrons and then using them to form bonds until all atoms satisfy the octet rule (or duet rule for hydrogen) - allows us to draw the Lewis structure accurately.

Molecular Geometry

Molecular geometry, or molecular shape, is crucial because it influences the molecule's properties and how it interacts with other molecules. The arrangement of atoms around a central atom is determined by the number of bonding electron pairs (shared between atoms in a molecule) and lone electron pairs (remaining on the central atom).

VSEPR theory (Valence Shell Electron Pair Repulsion theory) allows us to predict the molecular geometry of a molecule. Basically, it states that electron pairs around a central atom will arrange themselves as far as possible to minimize repulsion. In the exercise, for molecules like \(SF_2\), the two bonding pairs and two lone pairs on the sulfur create a 'bent' shape. For \(SiH_4\), the four bonding pairs result in a 'tetrahedral' shape around the silicon atom with no lone pairs. Understanding molecular geometry helps us grasp how molecules fit together, react, and donate or accept electrons, aiding in visualizing the three-dimensional form from the two-dimensional Lewis structures.