Question

Compare and contrast the bonding found in the \(\mathrm{H}_{2}(g)\) and HF \((g)\) molecules with that found in NaF(s).

Short answer

In H2, there is a non-polar covalent bond with a bond strength of 436 kJ/mol and a molecular orbital configuration of \(\sigma_{1s}^2\). HF has a polar covalent bond with a bond strength of 565 kJ/mol, formed by the overlap of hydrogen's 1s orbital and fluorine's 2p orbital. NaF has an ionic bond, with a bond strength of 924 kJ/mol (lattice energy), formed by the transfer of an electron from sodium to fluorine and the electrostatic attraction between the resulting Na+ and F- ions.

Step-by-step solution

Identify the type of bonding in each compound

In H2, the bonding is a covalent bond between two non-metal hydrogen atoms. In HF, the bonding is a polar covalent bond between a non-metal hydrogen atom and a non-metal fluorine atom. In NaF, the bonding is an ionic bond between a metal sodium atom and a non-metal fluorine atom.

Compare bond strength and polarity

H2 has a non-polar covalent bond with a bond strength of about 436 kJ/mol. HF has a polar covalent bond with a bond strength of about 565 kJ/mol, which is greater than that of H2 due to the higher electronegativity of fluorine. NaF has ionic bonding with a bond strength of about 924 kJ/mol (lattice energy), which is greater than both H2 and HF due to the strong electrostatic attraction between the positive sodium ion and the negative fluoride ion.

Examine molecular orbital configurations

The molecular orbital configuration of H2 is represented by the following notation: \(\sigma_{1s}^2\). This means that there are two electrons in the \(\sigma_{1s}\) orbital, which is a bonding orbital that is formed by the overlap of the hydrogen 1s atomic orbitals. The molecular orbital configuration of HF is also represented by the overlap of atomic orbitals: the hydrogen 1s orbital with the fluorine 2p orbital. The resulting molecular orbitals consist of a bonding \(\sigma\) orbital and an antibonding \(\sigma^{*}\) orbital. The higher electronegativity of fluorine results in the electron density being more concentrated around the fluorine atom and gives rise to the polar nature of the bond. As for NaF, there are no molecular orbitals since it is an ionic compound. Instead, the sodium atom loses an electron from its 3s orbital to form a positively charged Na+ ion, and the fluorine atom gains an electron in its 2p orbital to form a negatively charged F- ion. The attraction between these oppositely charged ions is responsible for the strong ionic bond in NaF.

Summarize and contrast the bonding in H2, HF, and NaF

In summary, the bonding in H2 is a non-polar covalent bond with a molecular orbital configuration of \(\sigma_{1s}^2\) and has a bond strength of 436 kJ/mol. The bonding in HF is a polar covalent bond where the fluorine atom has a higher electronegativity, resulting in a stronger bond strength of 565 kJ/mol. The molecular orbitals in HF are formed by the overlap of the hydrogen 1s orbital and the fluorine 2p orbital. NaF has an ionic bond with a bond strength of 924 kJ/mol (lattice energy), and no molecular orbitals are present. Instead, NaF's bonding arises from the transfer of an electron from the sodium atom to the fluorine atom and the subsequent electrostatic attraction between the resulting Na+ and F- ions.

Key concepts

Covalent Bonding

Covalent bonding is a type of chemical bond where two atoms share electrons. This typically occurs between non-metal atoms with similar electronegativities. When atoms share electrons, they can achieve a stable electron configuration similar to noble gases.

In a covalent bond like the one in \(H_2\), each hydrogen atom shares one electron with the other, forming a strong bond. This type of bond is often found in molecules of the same element or in compounds where the atoms have similar electronegativities.

• Examples include non-polar molecules like \(H_2\).
• Covalent bonds can vary widely in strength and are influenced by the distances between nuclei and shared electron pairs.

An important trait of covalent bonding is that it does not result in the formation of charged particles. Covalent bonds can be single, double, or even triple, depending on how many electron pairs are shared between the atoms.

Ionic Bonding

Ionic bonding occurs when one atom donates one or more of its electrons to another, creating charged ions that attract each other. This usually happens between metal and non-metal atoms. Metal atoms typically lose electrons to become positively charged cations, while non-metal atoms gain electrons to become negatively charged anions.

An example of ionic bonding is seen in \(NaF\). The sodium atom transfers an electron to the fluorine atom, resulting in a \(Na^+\) ion and a \(F^-\) ion. The electrostatic attraction between these oppositely charged ions forms a strong ionic bond.

• Ionic bonds are typically stronger than covalent bonds due to the strong attraction between the ions.
• Ionic compounds form crystalline lattice structures which contribute to their high melting and boiling points.

Although ionic bonds are strong, the bonds can break when the compound is dissolved in water or melted, allowing the ions to move freely.

Bond Polarity

Bond polarity is a concept that explains the unequal sharing of electrons in covalent bonds due to differences in electronegativity between the atoms involved. When atoms with different electronegativities form a bond, the shared electrons are pulled closer to the more electronegative atom.

This results in a partial charge distribution across the molecule, making it polar. HF is a classic example, where the fluorine atom is more electronegative than the hydrogen atom.

• Polarity depends on the relative electronegativities of the atoms involved. Fluorine, being highly electronegative, attracts the shared electrons more strongly than hydrogen.
• Polar covalent bonds can give rise to dipole moments, which impacts the molecule's physical and chemical properties.

Polarity is crucial in determining how molecules interact with each other and with solvents, greatly affecting properties like solubility, boiling point, and interaction with light.

Molecular Orbital Theory

Molecular orbital theory (MOT) is a powerful method for understanding the electronic structure of molecules. Unlike simpler models like the Lewis Structure approach, MOT views electrons in molecules as occupying molecular orbitals that extend over all the atoms in the molecule.

For \(H_2\), the molecular orbital configuration is \(\sigma_{1s}^2\) with both electrons in a bonding orbital, stabilizing the molecule. This highlights how atomic orbitals combine to form molecular orbitals with lower energy, which helps to stabilize the molecule.

• MOT can explain phenomena like bond order, magnetism, and why certain molecules exist or not.
• It helps in understanding the behavior of electrons in more complex molecules.

In ionic compounds like \(NaF\), MOT is not commonly applied, as the bond arises from electron transfer rather than shared electron clouds. MOT is particularly useful for predicting the chemical and physical behavior of organic compounds and transition metal complexes.