Question

State whether each of these statements is true or false. (a) An oxygen-oxygen double bond is shorter than an oxygenoxygen single bond. (b) There are three lone pair electrons in the \(\mathrm{NH}_{3}\) molecule. (c) The \(\mathrm{C}-\mathrm{C}\) bond in ethene is longer than the \(\mathrm{C}-\mathrm{C}\) bond in polyethene. (d) The \(\mathrm{C}-\mathrm{Cl}\) bond is shorter than the \(\mathrm{C}-\mathrm{Br}\) bond. \((\mathbf{e})\) The greater the difference in the electronegativity of atoms in a bond, the stronger the bond.

Short answer

(a) True (b) False (c) True (d) True (e) False

Step-by-step solution

Statement (a) - Oxygen-oxygen double bond and single bond

To determine whether an oxygen-oxygen double bond is shorter than an oxygen-oxygen single bond, we consider the fact that double bonds involve the sharing of more electrons between the bonded atoms. This leads to a higher electron density and stronger electrostatic attraction between the nuclei and the bonding electrons. As a result, double bonds are generally shorter and stronger than single bonds. Hence, the statement (a) is true.

Statement (b) - Lone pair electrons in NH3 molecule

The NH3 molecule has one nitrogen atom and three hydrogen atoms. Nitrogen has 5 valence electrons and each hydrogen atom contributes 1 electron. The nitrogen atom forms three single bonds with the three hydrogen atoms, using 3 of its 5 valence electrons. The remaining 2 electrons form a lone pair on the nitrogen atom. This means there is only one lone pair of electrons in the NH3 molecule. Thus, statement (b) is false.

Statement (c) - C-C bond in ethene and polyethene

Ethene has a carbon-carbon double bond, while polyethene has carbon-carbon single bonds. As previously discussed, double bonds are generally shorter and stronger than single bonds due to a higher electron density and stronger electrostatic attraction. Therefore, the C-C bond in ethene is shorter than the C-C bond in polyethene. Statement (c) is true.

Statement (d) - C-Cl and C-Br bond lengths

The bond length between two atoms is influenced by their atomic size. Generally, as you move down a group in the periodic table, the atomic size increases. Chlorine is above bromine in the periodic table, so it is smaller in size than bromine. As a result, the C-Cl bond is shorter than the C-Br bond. Statement (d) is true.

Statement (e) - Bond strength and electronegativity difference

The electronegativity of an atom is its ability to attract electrons in a bond. A bond formed between two atoms with a large electronegativity difference is polar and has a significant amount of ionic character. Such bonds are generally stronger than non-polar bonds. However, this statement is not universally true, as bond strength will also depend on other factors such as the size of the atoms and electron configuration. Therefore, statement (e) is false. In conclusion: - Statement (a) is true - Statement (b) is false - Statement (c) is true - Statement (d) is true - Statement (e) is false

Key concepts

Bond Length

Bond length refers to the average distance between the nuclei of two bonded atoms. The strength of this connection varies depending on several factors, primarily the type of bond: single, double, or triple.
A double bond involves more shared electrons than a single bond, leading to a higher electron density. This increased electron density results in a shorter and stronger bond because the attraction between the bonded nuclei and electrons is greater.
For instance, in the exercise's statement (a), an oxygen-oxygen double bond is indeed shorter than an oxygen-oxygen single bond. This is because double bonds share more electrons and thus, cause the atoms to stay closer together, which reduces the bond length.

• This principle applies to all covalent bonds: as the number of shared electron pairs increases, the bond length decreases.
• Besides, bond length is also influenced by the size of the atoms: the smaller the atomic radii, the shorter the bond length.

Electronegativity

Electronegativity is a measure of how strongly an atom can attract electrons towards itself in a chemical bond. Each element has a unique electronegativity value, which provides insight into the bond's nature.
Typically, bonds between atoms with a large difference in electronegativity exhibit a greater tendency towards ionic characteristics, resulting in a polar bond. In general, the more polar the bond, the stronger it is due to the electrostatic attractions between the charged atoms.
However, electronegativity alone does not determine bond strength, contrary to statement (e) in the exercise, which is false. Bond strength also depends on factors like:

• The bond type (single, double, or triple).
• Atomic size, since larger atoms generally form weaker bonds.

Understanding electronegativity is crucial when predicting molecular interactions and the behavior of different compounds.

Lone Pair Electrons

Lone pair electrons, also known as non-bonding pairs, are valence electrons not shared with another atom. They play a critical role in the shape and reactivity of molecules.
In the NH extsubscript{3} molecule example, nitrogen uses three of its five valence electrons to form single bonds with three hydrogen atoms. The remaining two electrons form a lone pair. So, statement (b) in the exercise is false, as there is only one lone pair.

• Lone pairs affect molecular geometry by repelling other electron pairs, leading to specific shapes according to VSEPR theory.
• They also influence physical properties and reactivity, as lone pairs can participate in chemical reactions or influence polar characteristics.

Lone pairs are pivotal for understanding advanced bonding concepts and analyzing chemical behaviors.