Ion dipole Forces

Molecules are like people; they feel attraction and are drawn to one another. These forces of attraction, called intermolecular forces (IMF), are the attractive forces between molecules and/or ions. The strongest of these forces are called ion-dipole forces of attraction.

In this article, we will look at ion-dipole forces and see why these are considered the strongest IMF. We will also dive into ion-induced dipoles, which are a similar, but much weaker force.

• This article is about ion-dipole forces

• We will review the concept of dipoles.

• Then we will look at how ion-dipole interactions occur and their energies.

• Lastly, we will look at ion-induced dipole forces and see why they are weaker than ion-dipole forces.

Ion-Dipole forces meaning

Before we dive into ion-dipole interactions, we first need to dive further into the concept of the dipole. A dipole is present in a molecule when one side of the molecule is more electronegative than the other. This means that one side is more likely to accept an electron (more electronegative), so it has a slight negative charge, while the other side is more likely to lose an electron, so it has a slight positive charge.

The trends in electronegativity are:

• Elements that are closer to the top-right of the periodic table (like fluorine, F) are very electronegative.
• Elements that are less electronegative (like francium, Fr) are closer to the bottom left.

Dipoles in molecules

The first thing we need to remember is that not every molecule with a difference in electronegativities will have a dipole. Let's look at a comparison:

Methane, CH₄, does not have a dipole, and that's for two reasons. The first is that, while there is a difference in electronegativities, the difference isn't large enough.

For a compound to be considered polar, the difference in electronegativity has to be greater than 0.4. Carbon has an electronegativity of 2.5, while Hydrogen has one of 2.2. The difference in electronegativity is negligible enough that the molecule isn't considered polar. The second reason is that the molecule is symmetrical. Even if the bonds were polar, the symmetry negates that. Think of it this way, if 4 people are playing tug-of-war and all of them are pulling with the same force, the center isn't going to move.

For NaCl, this molecule is polar and has a dipole. The difference in electronegativity between Na and Cl is over 2, so the bond is very polar. Na has a lower electronegativity, so the "positive" side points towards it, while the "negative" side points towards Cl. The arrow is referred to as the dipole moment.

The dipole moment is important for calculating the total strength of the ion-dipole interaction. The dipole moment is directly proportional to the potential energy of the interaction

Ion-dipole forces strength

The basic interaction here is the attraction/repulsion between the ion and the dipole.

Fig. 2 - The anion attracts and repels different ends of the dipole

Ion-dipole forces are contactless, so there will always be a distance separating the ion and molecule. We measure the energy of these forces using the formula for ion-dipole potential.

We note that the ion-dipole potential is derived from the Coulomb's Law potential, thusly:

1. Given the Coulomb Law potential:

$$ E = \frac{ -k q_1 q_2 } {r_1} $$

Where k is Coulomb's constant, q₁ is the charge on the ion, q₂ is the charge on the dipole, and r₁ is the radius between the ion and the molecule.

2. Next, we solve for the charge on the dipole, q2, and insert this into the Coulomb Law potential:

\begin{align}& \mu = q_2 * r_1 \qquad (solving~ for~ q_2,~ we~ get) \\& q_2 = \mu / r_1\end{align}

Subtituting for q2, we obtain:

$$ E = \frac{ -k q_1 (\frac {\mu} {r_1}) } {r_1} = \frac { -kq_1 \mu} { r_1^2} $$

The Ion-dipole force strength is dependent on three things: The magnitude of the dipole moment, the distance between the ion and the molecule, and the size of the polar molecule. Based on the equation, we can see why these first two things are important. The size of the molecule affects the radius between the ion and the molecule, but it also affects how easily the ion will interact with it. If we have a large molecule with many bonds, it will be more difficult for an ion to approach the molecule.

Ion-dipole force examples

The Ion-dipole forces are commonly found in solutions where an ionic compound has been dissolved in a polar solvent. The most common example is salt in water.

Fig. 3 - Ion-dipole forces between salt (NaCl) and water

The sodium (Na⁺) cation attracts the partially negative oxygen (O), while the chlorine (Cl^-) anion attracts the partially positive hydrogen (H).

In these kinds of solutions, the ionized compound and the polar solvent form a "net". While this example only shows 1 Na⁺ and 2 Cl^- ions, in reality, there would be a lot, and each ion would be attracted to several water molecules.

Ion-Induced Dipole forces

There is another kind of ion-dipole force that is weaker than the one we have looked at previously. These are ion-induced dipole forces.

So, what do we mean by "respond"? Let's look at a diagram of the interaction.

Fig. 4 - A cation induces a dipole in a non-polar molecule

The electrons in the molecule are attracted to the cation. Since these electrons are being pulled, the electron density is shifting. Now that more electrons are on the left side than the right, a dipole is formed.

These interactions will be much weaker, since the charge of the induced dipole in a non-polar molecule is smaller than the charge of a dipole in a polar molecule.