Question

Write Lewis structures and predict the molecular structures of the following. (See Exercises 111 and \(112 .\) ) a. \(\mathrm{OCl}_{2}, \mathrm{KrF}_{2}, \mathrm{BeH}_{2}, \mathrm{SO}_{2}\) b. \(\mathrm{SO}_{3}, \mathrm{NF}_{3}, \mathrm{IF}_{3}\) c. \(\mathrm{CF}_{4}, \mathrm{SeF}_{4}, \mathrm{KrF}_{4}\) d. \(\mathrm{IF}_{5}, \mathrm{AsF}_{5}\) Which of these compounds are polar?

Short answer

The polar compounds among the given examples are OCl2, SO2, NF3, IF3, SeF4, and IF5. Their respective molecular structures are bent, bent, trigonal pyramidal, T-shaped, see-saw, and square pyramidal. Other compounds are non-polar due to their symmetric molecular structures.

Step-by-step solution

a. OCl2, KrF2, BeH2, SO2

1. OCl2: Lewis structure: O is in the center, with two Cl atoms bonded to it. O has two lone pairs, while each Cl atom has three lone pairs. Molecular structure: Bent Polarity: Polar 2. KrF2: Lewis structure: Kr is in the center, with two F atoms bonded to it. Kr has three lone pairs, while each F atom has six lone pairs. Molecular structure: Linear Polarity: Non-polar 3. BeH2: Lewis structure: Be is in the center, with two H atoms bonded to it. Be has no lone pairs, while each H atom has no lone pairs. Molecular structure: Linear Polarity: Non-polar 4. SO2: Lewis structure: S is in the center, with two O atoms bonded to it. S has one lone pair, while each O atom has two lone pairs. Molecular structure: Bent Polarity: Polar

b. SO3, NF3, IF3

1. SO3: Lewis structure: S is in the center, with three O atoms bonded to it. S has no lone pairs, while each O atom has two lone pairs. Molecular structure: Trigonal planar Polarity: Non-polar 2. NF3: Lewis structure: N is in the center, with three F atoms bonded to it. N has one lone pair, while each F atom has three lone pairs. Molecular structure: Trigonal pyramidal Polarity: Polar 3. IF3: Lewis structure: I is in the center, with three F atoms bonded to it. I has two lone pairs, while each F atom has three lone pairs. Molecular structure: T-shaped Polarity: Polar

c. CF4, SeF4, KrF4

1. CF4: Lewis structure: C is in the center, with four F atoms bonded to it. C has no lone pairs, while each F atom has three lone pairs. Molecular structure: Tetrahedral Polarity: Non-polar 2. SeF4: Lewis structure: Se is in the center, with four F atoms bonded to it. Se has one lone pair, while each F atom has three lone pairs. Molecular structure: See-saw Polarity: Polar 3. KrF4: Lewis structure: Kr is in the center, with four F atoms bonded to it. Kr has two lone pairs, while each F atom has three lone pairs. Molecular structure: Square planar Polarity: Non-polar

d. IF5, AsF5

1. IF5: Lewis structure: I is in the center, with five F atoms bonded to it. I has one lone pair, while each F atom has three lone pairs. Molecular structure: Square pyramidal Polarity: Polar 2. AsF5: Lewis structure: As is in the center, with five F atoms bonded to it. As has no lone pairs, while each F atom has three lone pairs. Molecular structure: Trigonal bipyramid Polarity: Non-polar In summary, the polar compounds are: OCl2, SO2, NF3, IF3, SeF4, and IF5.

Key concepts

Lewis Structure

The Lewis structure is a way to visualize the arrangement of electrons around atoms within a molecule. It involves drawing dots to represent valence electrons and lines to represent bonds between atoms. To create an accurate Lewis structure, one must follow a few steps: determine the total number of valence electrons available, place the least electronegative atom at the center (except hydrogen), draw single bonds between central and surrounding atoms, and then fill the octet of the surrounding atoms. Any remaining electrons are placed on the central atom.
If there are still not enough electrons to fill the octet of the central atom, double or triple bonds may be formed. For example, in the case of SO₂, sulfur is the central atom with oxygen atoms bonded to it. It's important to complete the octets of oxygen atoms first and then place any remaining electrons on the sulfur. This often leads to multiple bonding situations to satisfy the octet rule.

Molecular Geometry

The shape of a molecule, known as molecular geometry, is determined by the spatial arrangement of its bonds and can significantly affect its physical and chemical properties. The VSEPR (Valence Shell Electron Pair Repulsion) theory aids in predicting geometry by stating that electron pairs will arrange themselves to minimize repulsion.
For example, a molecule with two bond pairs such as BeH₂ will be linear. On the other hand, a molecule like SF₄ with five electron pairs (four bonding pairs and one lone pair) adopts a see-saw shape. It's the balance between lone pairs and bonding pairs that determines the actual geometry.

Molecular Polarity

Molecular polarity is a measure of how equally bonding electrons are shared between atoms in a molecule. This is influenced by both the molecule's geometry and the electronegativity of its atoms. Non-polar molecules have an even distribution of electrons, while polar molecules have a notable difference in electron distribution, giving rise to a dipole moment.
A molecule like CF₄ is non-polar because the symmetrical tetrahedral shape allows the individual bond polarities to cancel out. However, a molecule such as H₂O has a bent shape, and because of the higher electronegativity of oxygen compared to hydrogen, there's an uneven distribution of charge, making water a polar molecule.

Electronegativity

Electronegativity is a chemical property that describes the ability of an atom to attract electrons towards itself in a chemical bond. It's crucial for determining the polarity of molecules. The higher an atom's electronegativity, the more strongly it will pull shared electrons toward itself.
On the periodic table, electronegativity increases from left to right and decreases from top to bottom. This helps to assess which atoms in a molecule will attract electrons more. For instance, in the molecule NF₃, nitrogen is less electronegative than fluorine, which results in a polar bond. Electronegativity differences can also give rise to non-polar bonds if they are small, as in the case of KrF₂.

VSEPR Theory

The VSEPR theory, which stands for Valence Shell Electron Pair Repulsion, provides a systematic set of rules for predicting the shape of molecules based on the premise that electron pairs around a central atom will spread out as much as possible to minimize repulsion.
Applying VSEPR theory, molecules like IF₅ take on a square pyramidal shape due to its five bonding pairs and one lone pair. The lone pair's influence causes the molecule to deviate from a perfect octahedral shape. Understanding VSEPR theory is integral to predicting molecular shapes and consequently their reactivity and interactions.