Problem 122 - Chemistry Matter and Change

Question

Identify which molecule is polar. (Chapter 8) \(\begin{array}{ll}{\text { a. } \operatorname{SiH}_{4}} & {\text { c. } \mathrm{H}_{2} \mathrm{S}} \\ {\text { b. } \mathrm{NO}_{2}} & {\text { d. } \mathrm{NCl}_{3}}\end{array}\)

Step-by-Step Solution

Answer

Among the given molecules, only molecule b. NO2 is polar due to its bent geometry and significant difference in electronegativity between the central atom (N) and the surrounding atoms (O). The other three molecules have either symmetrical geometries or too small electronegativity differences to be considered polar.

Step 1: Identify the Molecular Geometry of each option

Using VSEPR theory, we identify the molecular geometry of each molecule. a. SiH4: Tetrahedral b. NO2: Bent c. H2S: Bent d. NCl3: Trigonal pyramidal

Step 2: Calculate Electronegativity Difference of Each Molecule

We calculate the electronegativity difference for each central atom and the surrounding atom. The electronegativity values can be found on the periodic table or using an electronegativity scale such as the Pauling scale. Below are the electronegativities for the particular elements involved: Si: 1.90, H: 2.20, N: 3.04, O: 3.44, S: 2.58, Cl: 3.16 a. SiH4: \(\Delta_{Si-H} = |1.90 - 2.20| = 0.30\) b. NO2: \(\Delta_{N-O} = |3.04 - 3.44| = 0.40\) c. H2S: \(\Delta_{H-S} = |2.20 - 2.58| = 0.38\) d. NCl3: \(\Delta_{N-Cl} = |3.04 - 3.16| = 0.12\)

Step 3: Determine Polar Molecules

For a molecule to be polar, it needs to have an asymmetrical molecular geometry with different electronegativities for the respective central and surrounding atoms. a. SiH4: Tetrahedral geometry, the small electronegativity difference, 0.30, is evenly distributed around the central atom due to symmetry. Therefore, it's non-polar. b. NO2: Bent geometry, 0.40 electronegativity difference, and asymmetrical distribution of the electron density - polar. c. H2S: Bent geometry and 0.38 electronegativity difference but symmetrical and therefore non-polar. d. NCl3: Trigonal pyramidal geometry but with small electronegativity difference and symmetrical distribution, non-polar.

Step 4: Conclusion

Among the given molecules, only molecule b. NO2 is polar due to its bent geometry and significant difference in electronegativity between the central atom (N) and the surrounding atoms (O). The other three molecules have either symmetrical geometries or too small electronegativity differences to be considered polar.

Key Concepts

VSEPR theory

VSEPR theory stands for Valence Shell Electron Pair Repulsion theory, a fundamental principle in chemistry for predicting the shape of molecules. It is based on the idea that electron pairs around a central atom will repel each other and thus will arrange themselves as far apart as possible to minimize repulsion.
This theory helps us determine the three-dimensional shape of molecules, which is crucial because a molecule's shape affects its physical and chemical properties.
For example, by applying VSEPR to the exercise's molecules, you can identify that

• SiH4 has a tetrahedral geometry because it has four bonding pairs of electrons.
• NO2 has a bent shape due to the presence of a lone pair of electrons.
• H2S is also bent due to two lone pairs of electrons.
• NCl3 forms a trigonal pyramidal shape due to one lone pair.

Identifying these shapes helps to further assess molecule polarity by observing how electron pairs are distributed.

Electronegativity

Electronegativity refers to an atom's ability to attract electrons towards itself. This concept plays a crucial role in determining molecule polarity.
Each element has a specific electronegativity value, typically measured on the Pauling scale. When you compute

• the difference in electronegativity between two bonded atoms, it hints at the bond's character such as nonpolar covalent (little or no difference), polar covalent (moderate difference), or ionic (large difference).

In our exercise, calculating these differences:

• SiH4 (0.30), NO2 (0.40), H2S (0.38), and NCl3 (0.12),

helps determine whether a molecule could be polar. A greater difference often indicates a polar bond, contributing significantly to a molecule's overall polarity.

Molecular Geometry

Molecular geometry influences a molecule's properties, including its polarity. Geometry is how atoms in a molecule are arranged in three dimensions, which determines how bonds influence each other.
The shape can be linear, bent, tetrahedral, etc., depending on the number of bonding and lone electron pairs around the central atom.
For instance:

• SiH4's tetrahedral structure leads to symmetry, balancing out any dipoles.
• NO2, with its bent shape, introduces an asymmetry that often results in a polar molecule despite having similar elements bonded.
• H2S also has a bent structure. However, its shape causes bond dipoles to cancel out, making it non-polar.
• Lastly, NCl3 forms a trigonal pyramidal shape that can lead to polarity if there's an unequal distribution.

So, a molecule's geometry plays a pivotal role in defining if a molecule is polar or non-polar.

Bent Molecule

Molecules with a bent shape often show unique properties due to their structure. A bent shape arises when there are lone pairs on the central atom that lead to an angular arrangement of bonds.
Take NO2 and H2S from our exercise as examples:

• NO2 has one unpaired electron that causes the molecule to take on a bent geometry due to electron pair repulsion, and thus, results in a polar molecule as the dipoles don't cancel out.
• In contrast, H2S has two lone pairs causing the bent shape. However, the dipoles in its bonds cancel compared to the structure, resulting in a non-polar molecule.

Bent molecules are fascinating, as the presence and placement of lone pairs significantly influence the molecule’s polarity.

Symmetrical and Asymmetrical Molecules

The symmetry in molecular geometry plays a pivotal role in determining whether a molecule is polar or non-polar.
Symmetrical molecules tend to be non-polar since any dipoles within the molecule cancel out due to their even distribution.
For instance:

• SiH4 is a symmetric tetrahedral molecule where bond dipoles cancel each other out, making it non-polar.
• NCl3 has a trigonal pyramidal shape. Its symmetry and low electronegativity difference contribute to its non-polar nature.

On the other hand, asymmetrical molecules often result in polar molecules due to uneven distribution, as seen in NO2:

• With its bent geometry, caused by the unpaired electron, NO2 results in polar characteristics since the dipoles don't cancel out.

Thus, symmetry can be an easy and quick indicator of molecular polarity.