Question

For each statement, indicate whether it is true or false. (a) $\ln$ order to make a covalent bond, the orbitals on each atom in the bond must overlap. (b) A $p$ orbital on one atom cannot make a bond to an $s$ orbital on another atom. $(\mathbf{c})$ Lone pairs of electrons on an atom in a molecule influence the shape of a molecule. (d) The 1 s orbital has a nodal plane. $(\mathbf{e})$ The $2p$ orbital has a nodal plane.

Short answer

a) TRUE, b) FALSE, c) TRUE, d) FALSE, e) TRUE

Step-by-step solution

a) Orbital overlap for covalent bonds

To form a covalent bond, the orbitals of the bonded atoms must overlap, allowing electrons to be shared between the atoms. The statement is TRUE.

b) p orbital - s orbital bonding

It is possible for a p orbital on one atom to bond with an s orbital on another atom. For example, a carbon atom with three p orbitals and a hydrogen atom with one s orbital can form a C-H bond in a molecule like methane (CH4). The statement is FALSE.

c) Lone pairs' influence on molecular shape

Lone pairs of electrons on an atom can indeed influence the shape of a molecule, as they occupy space and create repulsion with surrounding electron regions. This can cause bond angles to change and result in different molecular geometries. The statement is TRUE.

d) 1s orbital nodal plane

A nodal plane is a region in space where electron probability density is zero. The 1s orbital is a spherical orbital surrounding the nucleus, and there is no region with zero electron probability density. Therefore, the statement is FALSE.

e) 2p orbital nodal plane

The 2p orbital has a dumbbell shape, with two lobes separated by a nodal plane running through the nucleus. In this nodal plane, electron probability density is zero. So, the statement is TRUE.

Key concepts

Orbital Overlap

Orbital overlap is a fundamental concept in the formation of covalent bonds. In a covalent bond, two atoms share pairs of electrons. For this sharing to occur efficiently, the atomic orbitals of the bonding atoms must overlap. This overlap brings the nuclei of the atoms closer together, allowing them to share electrons effectively.
An example of this process can be seen in the formation of the hydrogen molecule (H extsubscript{2}). When two hydrogen atoms approach each other, their 1s orbitals overlap, creating a region of electron density between the nuclei. This electron density acts as a glue that holds the atoms together, forming a stable bond.

Overlapping orbitals can vary, such as:

• s-s overlap, as seen in hydrogen molecules (H extsubscript{2})
• s-p overlap, such as in hydrogen fluoride (HF)
• p-p overlap, found in molecules like O extsubscript{2} and N extsubscript{2}

Understanding orbital overlap helps explain why certain atoms bond readily and the resulting molecular structure.

Molecular Geometry

Molecular geometry describes the 3D arrangement of atoms in a molecule and is influenced by the number of electron pairs surrounding a central atom, including both bonding pairs (shared electrons) and lone pairs (non-bonding electrons).
For example, the geometry of a water molecule (H extsubscript{2}O) is not linear despite having two hydrogen atoms bonded to oxygen. Instead, it adopts a bent shape due to the two lone pairs of electrons on the oxygen, causing repulsions that adjust the bond angles away from a linear configuration.

Molecular shapes can often be predicted using the VSEPR (Valence Shell Electron Pair Repulsion) theory, which considers:

• Total number of electron pairs (bonding and lone) around the central atom.
• Repulsions between electron pairs, which dictate the spatial arrangement.

This knowledge is crucial for predicting and understanding the properties and reactivity of molecules.

Electron Configuration

Electron configuration is a way of arranging electrons around the nucleus of an atom, following a set pattern to minimize energy. Electrons fill orbitals in a specific order, following rules like the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
Understanding electron configurations can help predict chemical properties and bonding capabilities of elements. Each atom has a unique electron configuration that influences its reactivity and interactions with other atoms.

For instance:

• The electron configuration of oxygen is 1s extsuperscript{2} 2s extsuperscript{2} 2p extsuperscript{4}, indicating 6 valence electrons ideal for forming two covalent bonds.
• Sodium's (Na) configuration ends in 3s extsuperscript{1}, showing a tendency to lose one electron and form a cation.

Mastering electron configuration is essential for understanding periodic trends and the chemical behavior of elements in reactions.