Question

Indicate the bond polarity (show the partial positive and partial negative ends) in the following bonds. a. \(C-O\) b. \(P-H\) \(\mathbf{c} . \quad \mathbf{H}-\mathbf{C l}\) d. \(\mathrm{Br}-\mathrm{Te}\) \(\mathbf{e} . \mathbf{S e}-\mathbf{S}\)

Short answer

The bond polarities are as follows: a. C(δ+)-O(δ-) b. P(δ+)-H(δ-) c. H(δ+)-Cl(δ-) d. Br(δ-)-Te(δ+) e. Se(δ)-S(δ) (nonpolar)

Step-by-step solution

Determine the electronegativity values of the atoms involved in each bond

First, let's list the electronegativity values of the atoms involved in each bond: C - 2.55 O - 3.44 P - 2.19 H - 2.20 Cl - 3.16 Br - 2.96 Te - 2.1 Se - 2.55 S - 2.58

Indicate bond polarity for C-O

Comparing the electronegativity values of the two atoms, we find that O is more electronegative than C (3.44 > 2.55). Hence, the oxygen atom will attract the shared electrons more towards itself, leading to a partial negative charge (denoted by δ-) on oxygen and a partial positive charge (denoted by δ+) on carbon. C(δ+)-O(δ-)

Indicate bond polarity for P-H

Comparing the electronegativity values of the two atoms, we find that H is slightly more electronegative than P (2.20 > 2.19). Hence, the hydrogen atom will attract the shared electrons more towards itself, leading to a partial negative charge (denoted by δ-) on hydrogen and a partial positive charge (denoted by δ+) on phosphorus. P(δ+)-H(δ-)

Indicate bond polarity for H-Cl

Comparing the electronegativity values of the two atoms, we find that Cl is more electronegative than H (3.16 > 2.20). Hence, the chlorine atom will attract the shared electrons more towards itself, leading to a partial negative charge (denoted by δ-) on chlorine and a partial positive charge (denoted by δ+) on hydrogen. H(δ+)-Cl(δ-)

Indicate bond polarity for Br-Te

Comparing the electronegativity values of the two atoms, we find that Br is more electronegative than Te (2.96 > 2.1). Hence, the bromine atom will attract the shared electrons more towards itself, leading to a partial negative charge (denoted by δ-) on bromine and a partial positive charge (denoted by δ+) on tellurium. Br(δ-)-Te(δ+)

Indicate bond polarity for Se-S

Comparing the electronegativity values of the two atoms, we find that the values for Se (2.55) and S (2.58) are very close. Therefore, the bond between Se and S is likely to be nonpolar. Since the electronegativity difference is so slight, we can consider them as having a similar charge. Se(δ)-S(δ)

Key concepts

Electronegativity

Electronegativity is a chemical property that describes the ability of an atom to attract a shared pair of electrons towards itself in a chemical bond. On the periodic table, electronegativity tends to increase from left to right and decrease from top to bottom. For instance, fluorine is the most electronegative element with a value of 3.98 on the Pauling scale, which is the most commonly used scale for electronegativity.

Understanding electronegativity is crucial in predicting the behavior of atoms in molecules. It helps in determining how the electrons in a bond are distributed between the bonded atoms. In a bond, the atom with higher electronegativity will pull the electron cloud closer to itself, causing the development of partial charges. These slight variations in charge are what lead to bond polarity.

Partial Charges

Partial charges arise in a molecule when there is an uneven distribution of electrons between bonded atoms. This is caused by differences in electronegativity. The more electronegative atom acquires a partial negative charge (denoted as \(\delta-\)), while the less electronegative atom obtains a partial positive charge (denoted as \(\delta+\)).

Particularly in chemistry, understanding partial charges is fundamental to grasping concepts such as molecular polarity, reactivity, and the physical properties of substances. For example, the presence of partial charges in water molecules \(H_2O\) is responsible for the unique properties of water, including its excellent solvent abilities and higher boiling point relative to other similar sized molecules.

Polar Bonds

Polar bonds are chemical bonds where there is a significant difference in electronegativity between the bonded atoms. This difference causes a dipole moment — a measure of charge separation in the bond. In a dipole moment, one end of the bond is slightly negative, while the opposite end is slightly positive.

For example, in the bond between hydrogen and chlorine (H-Cl), chlorine has a higher electronegativity and thus carries a partial negative charge, making the bond polar. The existence of polar bonds in a substance affects its physicochemical properties, like solubility in water, boiling and melting points, and interaction with electromagnetic fields.

Nonpolar Bonds

In contrast to polar bonds, nonpolar bonds occur when bonded atoms share their electrons equally, which generally happens in two scenarios: when the atoms are identical, as in the bonds found in \(O_2\) or \(N_2\), or when the difference in electronegativity between the atoms is negligible.

A perfect example of a nonpolar bond is the diatomic nitrogen molecule (N-N) where both nitrogen atoms have equal electronegativity, sharing the bonding electrons equally. Nonpolar substances do not mix well with water, often tend to have lower melting and boiling points, and can't conduct electricity when in a liquid state.