Problem 52 - General Chemistry: Principles and Modern Applications

Question

Draw Lewis structures for the following species, indicating formal charges and resonance where applicable: (a) \(\mathrm{HOSO}_{3}\) (b) \(\mathrm{H}_{2} \mathrm{NCN}\) (c) \(\mathrm{FCO}_{2}^{-}\) (d) \(\mathrm{S}_{2} \mathrm{N}_{2}\) (a cyclic structure with \(\mathrm{S}\) and \(\mathrm{N}\) alternating)

Step-by-Step Solution

Answer

The Lewis structures for the compounds are as follows: For HOSO3, a double bond is present between S and one O atom. For H2NCN, a triple bond is present between C and N atom. For FCO2-, F is single bonded to C and one O is double bonded to C. For S2N2, there is a double bond between each N-S pair.

Step 1: Lewis structure of HOSO3

Start by drawing a skeleton of the molecule, placing Hydrogen (H) on Sulphur (S), Sulphur (S) on Oxygen (O), and Oxygen (O) on Oxygen (O) forming the molecule H-S-O-O-O. Determine the total number of electrons available which is 24: 1 (from H) + 6 (from S) + 18 (from 3 O-atoms). Draw single bonds between the atoms and then fill the octets of Oxygen and Sulphur. A double bond is required between an Oxygen atom and the Sulphur atom to fill their octets.

Step 2: Lewis structure of H2NCN

Draw the skeleton of the molecule as H-N-C-N-H. The total number of electrons is 16: 2 (from 2 H-atoms) + 5 (from N) + 4 (from C) + 5 (from N). Start drawing single bonds and octets, and notice you need to form a triple bond between Carbon and one of the Nitrogen atoms to satisfy their octets.

Step 3: Lewis structure of FCO2-

Start by drawing a skeleton molecule F-C-O-O. The total number of electrons is 24: 7 (from F) + 4 (from C) + 12 (from 2 O-atoms) + 1 (from additional electron). Since Carbon is the central atom and it can form four bonds, draw single bonds between F-C and C-O, and a double bond between C-O. Fill the octets around F and O, finally check each atom to see if it obeys the octet rule.

Step 4: Lewis structure of S2N2

Draw the skeleton molecule as S-N-S-N forming a cyclic structure. The total number of electrons is 20: 12 (from 2 S-atoms) + 10 (from 2 N-atoms). Start with single bonds and octets, notice that a double bond is needed between each N-S pair to satisfy the octet rule.

Key Concepts

Formal Charges

Formal charges help us determine the most stable Lewis structure for a molecule or ion. This is important for understanding where electrons are more likely to be found in a molecule. The formal charge can be calculated using the formula:

• Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (Bonding Electrons/2)

For each atom, count how many electrons are associated with it based on this formula. When drawing the Lewis structures for species like \(\mathrm{HOSO}_{3}, \mathrm{H}_{2} \mathrm{NCN}, \mathrm{FCO}_{2}^{-}, \mathrm{and} \mathrm{S}_{2} \mathrm{N}_{2}\), you assign electrons in each bond and those not shared directly with other atoms.
By minimizing the number of formal charges, we can infer which structure has the least energy and is most stable. Making sure the formal charge is as close to zero as possible enhances molecular stability, especially in neutral molecules.
Formal charges are useful in predicting the behavior of molecules in reactions.

Resonance Structures

Resonance structures occur when more than one valid Lewis structure exists for a molecule. This is often visible in molecules with double bonds or lone electron pairs. In such cases, resonance structures average the positions of electrons, leading to resonance stabilization.
For example, when looking at \(\mathrm{HOSO}_{3}\) and \(\mathrm{FCO}_{2}^{-}\), resonance structures allow us to represent delocalized electrons across the molecule.

• When drawing resonance structures, we do not move atoms, only the position of electrons.
• Different resonance forms do not occur separately but contribute to the "true" form.

This concept is crucial for predicting reaction sites and molecule stability as certain forms are more reactive than others.
Keep in mind, the resonance hybrid—which combines all resonance structures—gives a more accurate depiction of the electron distribution.

Octet Rule

The octet rule is a fundamental principle stating that atoms tend to form bonds until they surround themselves with eight electrons, achieving a noble gas configuration. This rule primarily applies to atoms in the second period of the periodic table.

• The octet rule explains the bonding preference and stability of molecules.
• Most atoms don't naturally have eight valence electrons, so they form bonds with other atoms, either by sharing or transferring electrons.

While drawing Lewis structures, as demonstrated in \(\mathrm{HOSO}_{3}, \mathrm{H}_{2} \mathrm{NCN}, \mathrm{FCO}_{2}^{-}, \mathrm{and} \mathrm{S}_{2} \mathrm{N}_{2}\), fulfilling the octet rule often requires creating double or even triple bonds.
In some cases, exceptions occur, such as in molecules with atoms having more than eight electrons, or less, but these are less common.
Be aware that while the octet rule is a guiding principle, it is not an absolute rule and should be considered alongside other bonding theories.

Molecular Geometry

Molecular geometry is the 3D arrangement of atoms within a molecule, determined by the positions of the electron pairs around the central atom. This concept is essential for understanding the shape and properties of a molecule.
Using VSEPR (Valence Shell Electron Pair Repulsion) theory, the geometry around a central atom can be predicted by examining how electron pairs repel each other.

• Common geometries include linear, bent, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.
• The geometry affects polarity, reactivity, phase of matter, color, magnetism, biological activity, and more.

When evaluating the given molecules such as \(\mathrm{S}_{2} \mathrm{N}_{2}\) and \(\mathrm{H}_{2} \mathrm{NCN}\), determining the molecular geometry helps in understanding their chemical behavior and interactions.
Remember that lone pairs of electrons also play a significant role in shaping the geometry and should not be overlooked when mapping molecular structure.