Question

Dipole-dipole forces become as the distance between the dipoles decreases (gets closer together).

Short answer

The strength of dipole-dipole forces between polar molecules depends on the distance between the dipoles, with the force being inversely proportional to the cube of the distance between them (\( F \propto \frac{p_1 p_2}{r^3} \)). As the dipoles get closer together (distance \(r\) decreases), the attractive force between them becomes stronger.

Step-by-step solution

Understanding Dipole-Dipole Forces

Dipole-dipole forces are one of the most common types of intermolecular forces that exist between polar molecules. A polar molecule has a net dipole moment due to the presence of polar bonds, which are bonds between atoms with different electronegativities. The positive and negative ends of a dipole are termed as the two poles of the dipole. These poles will orient themselves, such that the positive end of one dipole will align with the negative end of another dipole to create an attractive force.

The importance of distance

The strength of dipole-dipole forces depends on the distance between the interacting dipoles. When the dipoles are close together, the attractive forces between them are stronger. Conversely, when they are farther apart, the forces weaken. This is because the force acting between two dipoles is inversely proportional to the cube of the distance between them, as per Coulomb's Law.

Coulomb's Law for dipoles

To understand the relationship between dipole-dipole forces and the distance between dipoles, we can use Coulomb's Law for dipoles. In this context, Coulomb's Law states that the force acting between two dipoles is proportional to the product of their magnitudes and inversely proportional to the cube of the distance between them: \( F \propto \frac{p_1 p_2}{r^3} \) Here, \(F\) represents the force between the dipoles, \(p_1\) and \(p_2\) are the magnitudes of the dipoles, and \(r\) is the distance between the dipoles.

Analyzing the relationship

From the formula derived in step 3, we can see that as the distance between the dipoles, \(r\), decreases, the force \(F\) increases. As dipoles get closer together, the strength of the attractive force between them becomes stronger, due to the inverse relationship based on the cube of the distance \(r^3\). This explains the statement given in the exercise, "Dipole-dipole forces become stronger as the distance between the dipoles decreases (gets closer together)."

Key concepts

Intermolecular Forces

Intermolecular forces are the forces of attraction or repulsion that act between neighboring particles, such as atoms, molecules, or ions. They are weaker than intramolecular forces, the forces which hold a molecule together. Among the various types of intermolecular forces, there are van der Waals forces (including dipole-dipole forces), hydrogen bonds, and London dispersion forces.

These forces are crucial in determining the physical properties of a substance, including its boiling point, melting point, and solubility. Generally, stronger intermolecular forces result in higher melting and boiling points because more energy is required to break the interactions between the particles.

Polar Molecules

Polar molecules are chemical species with an uneven distribution of charge due to the presence of polar bonds. These bonds form between atoms with different electronegativities, where one atom has a stronger tendency to attract the shared electrons.

As a result, the more electronegative atom acquires a partial negative charge (\t\([-]\)), and the less electronegative atom acquires a partial positive charge (\t\([+]\)). This separation of charge creates a dipole moment, with one end of the molecule being slightly negative and the other end being slightly positive. Water (\t\(H_2O\)) is one of the most common examples of a polar molecule. The polarity of a molecule significantly influences its interactions with other molecules and its physical and chemical properties.

Coulomb's Law

Coulomb's Law describes the force of attraction or repulsion between two charged particles. In its classical form, it states that the force (\t\(F\)) is directly proportional to the product of the magnitudes of the charges (\t\(q_1\) and \t\(q_2\)) and inversely proportional to the square of the distance (\t\(r\)) between the charges:

\t\[ F = k \frac{q_1 q_2}{r^2} \]
\t
where \t\(k\) is Coulomb's constant. The law applies to point charges and becomes a useful analogy when describing the forces between dipoles in polar molecules, with the force being inversely proportional to the cube of the distance, as explained in the step by step solution.

Electronegativity

Electronegativity is a measure of how strongly an atom attracts bonding electrons to itself when it is part of a compound. The concept was introduced by Linus Pauling, and it's important for determining the type of bond—ionic, polar covalent, or non-polar covalent—that will form between atoms.

On the periodic table, electronegativity values typically increase from left to right and bottom to top. The most electronegative element is fluorine. Differences in electronegativity between bonded atoms create polarity in the resulting molecule. For instance, in water, oxygen is more electronegative than hydrogen, resulting in a polar molecule with distinct positive and negative regions.