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.
Ion-dipole forces describe the electrostatic attraction between an ion and a neutral molecule with a dipole. The ion will attract one side of the dipole and repel the other.
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:
Fig. 1 - CH4 is a nonpolar compound with no dipole, while NaCl is a polar compound with a dipole.
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 (μ) is the measurement of the magnitude of the dipole. The formula is:
$$ \mu = q * r $$
Where:
- q = partial charge of each end of the dipole
- r = separation between charges (i.e. bond length)
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
The anion (light-blue circle with a minus sign) attracts the positive end of the dipole while repelling the negative end. The colors represent electron density, with the cooler colors meaning less density, with the warmer colors meaning more density.
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.
The ion-dipole potential is the potential energy of an ion-dipole attraction. The formula is:
$$ E = \frac{ -k|q_1|\mu } {r_1^2} $$
Where:
- k = proportionality constant (Coulomb's constant)
- q1 = charge of the ion
- μ = dipole moment of the molecule
- r1 = the radius between the ion and the molecule
(the subscript is noted as 1 to differentiate between the q and r used to calculate μ)
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-dipole forces have an important role in proteins. They are typically utilized when a reaction requires high specificity or fixed geometry. For example, they act as gatekeepers in transporters and ion channels, making sure that only the appropriate ions are passing through the membrane. As another example, these interactions hold the enzymatic intermediate in a fixed position during an enzyme reaction so that it can proceed properly.
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.
An ion-induced dipole interaction is when an ion approaches a non-polar molecule and the electrons in the molecule "respond", creating a dipole.
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.
If you have ever used a magnet near a paperclip, then noticed that other paperclips are attracted to it, it's the same principle!
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.
Ion dipole Forces - Key takeaways
- Ion-dipole forces describe the electrostatic attraction between an ion and a neutral molecule with a dipole. The ion will attract one side of the dipole and repel the other.
- Dipoles in polar molecules are caused by a difference in electronegativities greater than 0.4.
- The formula for the ion-dipole potential is: \( E = \frac{ -k |q_1| \mu } {r_1^2} \)
- 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.
- An ion-induced dipole interaction is when an ion approaches a non-polar molecule and the electrons in the molecule "respond", creating a dipole. These interactions are much weaker than ion-dipole forces.
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