What is the relationship between the dipole moment and the bond moment? How is it possible for a molecule to have bond moments and yet be nonpolar?

When discussing the intricacies of chemical bonds, the term 'bond moment', or sometimes referred to as 'bond dipole moment', is crucial for understanding the polarity of individual bonds within a molecular structure. A bond moment is a vector quantity that indicates both the magnitude and the direction of a bond's polarity. It arises from the uneven distribution of electrons between two bonded atoms.

For example, in a bond between hydrogen and chlorine, chlorine has higher electronegativity and thus, pulls the shared electrons closer to itself, resulting in a charge imbalance and a bond moment pointing towards chlorine. The magnitude of this bond moment is measured in Debye units, which quantitatively express the strength of the bond's polarity.

However, to obtain a complete picture, one must consider all bond moments in a molecule, as their orientations and magnitudes will determine if the molecule itself exhibits overall polarity or remains nonpolar.

Delving deeper into the realm of molecular structures, we encounter the fundamental concept of molecule polarity. It's a trait that dictates many of a substance's physical properties, such as boiling point, solubility, and reactivity. Molecule polarity refers to the distribution of electrical charge over the atoms comprising the molecule.

The crux of the concept lies in the summation of individual bond moments, considering both their magnitude and geometric arrangement. When bond moments in a molecule do not cancel out due to asymmetrical geometry, the molecule is deemed polar. This gives rise to a net dipole moment, with distinguishable positive and negative poles.

Understanding this is vital as it explains why molecules like carbon dioxide, despite having polar bonds, are nonpolar. The linear shape of CO2 ensures that the bond moments point in opposite directions and cancel each other out, leading to no net dipole moment across the molecule.

Electronegativity is a term coined to gauge an atom's tendency to attract and bond with electrons. This property plays a starring role in determining the bond type between atoms and the resultant molecular characteristics.

An atom's electronegativity is influenced by both its atomic number and distance from the valence electrons to the nucleus. Atoms like fluorine have high electronegativity because they possess a strong ability to attract electrons, while elements such as cesium have low electronegativity. It's electronegativity differences that drive the formation of bond moments — the larger the difference, the more polar the bond.

In the periodic table, electronegativity values generally increase as one moves from left to right across a period and decrease down a group. Familiarity with electronegativity patterns is fundamental for predicting the polarity of molecules and bonding behavior in chemical reactions.