Question

Indicate which orbitals overlap to form the \(\sigma\) bonds in each compound. $$\quad \text { a. }\mathrm{BeBr}_{2} \quad \text { b. } \mathrm{HgCl}_{2} \quad \text { c. } \mathrm {ICN}$$

Short answer

For BeBr2, sigma bonds form through the overlap of Be's 2s and Br's 4p orbitals. In HgCl2, Hg's 6s and Cl's 3p orbitals overlap to form sigma bonds. In ICN, I's 5p with C's 2p and C's 2p with N's 2p orbitals overlap to form sigma bonds.

Step-by-step solution

Analyzing the Sigma Bonds in BeBr2

To determine the orbitals overlapping in BeBr2, examine the electron configuration of beryllium (Be) and bromine (Br). Beryllium is in the 2nd period and has an s orbital available for bonding, while bromine is in the 4th period and utilizes its p orbital for bonding. A sigma bond forms between the s orbital of Be and the p orbital of each Br.

Analyzing the Sigma Bonds in HgCl2

Mercury (Hg) can use its 6s orbital for bonding, while chlorine (Cl) uses its 3p orbitals. In HgCl2, each Cl atom will form a sigma bond with Hg through the overlap of Hg's 6s orbital and Cl's 3p orbital.

Analyzing the Sigma Bonds in ICN

In ICN, iodine (I) uses its 5p orbital, carbon (C) uses its 2p orbitals, and nitrogen (N) also uses its 2p orbitals for bonding. The sigma bond between I and C involves overlap of I's 5p orbital and C's 2p orbital. For the C—N sigma bond, the overlap occurs between C's 2p orbital and N's 2p orbital.

Key concepts

Molecular Orbitals

When we talk about molecular orbitals, we refer to the regions in a molecule where electrons are most likely to be found. They are formed when atomic orbitals, which are the regions in an atom where electrons reside around the nucleus, combine as a result of chemical bonding. This combination can result in various types of molecular orbitals, such as bonding orbitals, which lower the energy between two atoms and help them stick together, or antibonding orbitals, which do the opposite. For example, in BeBr₂, the molecular orbitals that form the sigma bonds are a result of the overlap between the s orbital of Be and the p orbitals of Br.

Understanding molecular orbitals is crucial in predicting and explaining the properties of molecules, such as their shape, magnetism, and reactivity. The concept of molecular orbitals also helps explain why certain molecules exist and how electrons are distributed within a molecule.

Chemical Bonding

Chemical bonding is the process that enables atoms to associate and form compounds. There are several types of chemical bonds, including ionic, covalent, and metallic bonds. In our exercise, the focus is on covalent bonds, particularly sigma bonds. These bonds are formed by the sharing of electron pairs between atoms. Covalent bonding occurs when atoms strive to reach a more stable electronic configuration.

This is accomplished through the overlap of atomic orbitals, where each contributing atom provides one electron to the shared pair. The strength and number of these bonds play a crucial role in determining a molecule's physical properties and chemical reactivity. For instance, BeBr₂ and HgCl₂ are formed by the overlap of different types of atomic orbitals creating covalent bonds that hold the atoms together.

Covalent Bonds

Zooming in on covalent bonds, we can see they are characterized by the sharing of electrons between atoms. One common type of covalent bond is the sigma bond. A sigma bond, denoted with the Greek letter σ, is the strongest type of covalent bond and is formed by the head-on overlap of orbitals. They allow for the free rotation of atoms and contribute to the stability and shape of molecules.

In our textbook case, the Be in BeBr₂ is bonded to the Br atoms via sigma bonds through the overlap of an s orbital from Be with a p orbital from each Br. Sigma bonds also appear in HgCl₂ and ICN where the s orbitals of Hg bond with the p orbitals of Cl, and the p orbitals of I overlap with the p orbitals of C, respectively.

Electron Configuration

Finally, let's dive into electron configuration. It refers to the arrangement of electrons in an atom's orbitals, and follows a set of rules dictated by quantum mechanics. The configuration is often depicted in shorthand using the principal quantum number and letter designations (s, p, d, f) to indicate the type of orbital and the number of electrons in them. This configuration determines an atom's potential to bond and the types of orbitals available for bonding.

In BeBr₂, the Be atom has its valence electrons in a 2s orbital, ready to bond. In HgCl₂, the Hg atom uses a 6s orbital, while the Cl atoms use 3p orbitals for forming bonds. Correctly understanding electron configuration is essential for predicting how atoms will interact and what types of molecular orbitals will result from these interactions.