Dipole Moments and Dipoles

Dipole Moments

After reading this article and doing the quizzes at the bottom, you should be able to

• Apply electronegativity to identify covalent and polar covalent bonds
• Rank similar molecules according to magnitude of dipole moment
• Understand how geometry affects dipole moment in bent, trigonal planar, trigonal pyramidal and tetrahedral molecules
• Predict the effect of dipole moments on boiling points

Table of Contents

1. 1. Dipole Moment In Molecules With One Polar Covalent Bond
   2. In Molecules With One Bond, The Dipole of the Bond is The Dipole Moment
   3. Dipoles In Bonds To Carbon
   4. How Geometry Affects Dipole Moment: BeCl2 And Linear Geometry 
   5. H2O And “Bent” Geometry
   6. BF3 and Trigonal Planar Geometry
   7. NH3 and Trigonal Pyramidal Geometry
   8. Tetrahedral Geometry and Dipole Moment
   9. Dipoles and Boiling Point
   10. Notes
   11. Quiz Yourself!
   12. (Advanced) References and Further Reading

1. Dipole Moment In Molecules With One Polar Covalent Bond

In a polar covalent bond, two atoms with vastly different electronegativities (>0.4) share a pair of bonding electrons.

Probably the simplest examples are the hydrogen halides HF, HCl, and HBr. In these molecules, we have a highly electronegative element (F, Cl, Br) sharing a pair of electrons with a much less electronegative element (H).

In theory, the bonding electrons are “shared” equally, since we count one electron from each covalent bond as “belonging” to that atom for the purposes of calculating formal charge.

In practice, the more electronegative element pulls an unequal share of the shared electron pair towards it.  (It’s like when Mom tells two brothers to “share” a bag of chips; the older, bigger brother will inevitably get more than his fair share.)

Since electrons bear a negative charge, the result is that the more electronegative element will bear a partial negative charge, and the less electronegative element will bear a partial positive charge since it has a deficit of electron density.

These adjacent, opposite partial charges are known as dipoles. 

The strength of a molecule’s dipole can be measured experimentally and is known as a dipole moment. (The unit of dipole moment is the DeBye, D )

For example, the dipole moments of the hydrogen halides have all been measured experimentally.

In simple diatomic molecules like HF, HCl, and HBr, the strength of the dipole in the molecule is equal to the strength of dipole in the bond.

Click to Flip

In larger molecules, geometry can result in some dipoles being cancelled out, as we shall see shortly.

2. In Molecules With One Bond, The Dipole of the Bond is The Dipole Moment

At the other end of the spectrum are purely covalent bonds such as those found in  Cl₂, Br₂, and N₂.

In these cases both atoms have equal electronegativities. Like two completely even teams in a tug-o’-war, the forces are completely balanced.

These diatomic molecules have an overall dipole moment of zero.

3. Dipoles In Bonds To Carbon

Since this is an organic chemistry course, you probably could have guessed where this is going. We’re going to talk a lot about bonds to carbon!

There are no diatomic molecules of carbon to talk about, so we’re just going to zoom out for a second and talk about the bonds.

In a C-H bond there is a small dipole. Being more electronegative, the carbon (2.5) is slightly electron rich (δ-) and the hydrogen (2.2)  is slightly electron poor (δ+).

It’s good to file away in your brain the fact that carbon is somewhat electron-rich when connected to H because we will use that later (See Article: 3 Factors That Stabilize Carbocations).

But for now, we typically classify C-H bonds as covalent (rather than polar covalent) because hydrocarbons (i.e. molecules just composed of carbon and hydrogen) are very non-polar, in that they do not mix very well with polar solvents like water.

If we replace a C-H bond with a C-N or C-O bond, a dipole also results. But notice in this instance that the polarity is flipped: due to the greater electronegativity of oxygen and nitrogen relative to carbon, the carbon is electron-poor (δ +) and the other atom is electron-rich (δ – ). This will also have important consequences, as we shall see shortly!

Bonds between carbon and highly electronegative elements tend to be classified as polar covalent bonds.

results in a large enough dipole that molecules start to have some polar properties. For that reason bonds between carbon and highly electronegative elements fall into the realm of polar covalent bonds.

Note in these cases that carbon goes from being partially electron-rich (δ-)  to partially electron-poor (δ+)

4. How Geometry Affects Dipole Moment: BeCl2 And Linear Geometry

Individual bonds have dipoles. But whether a molecule has a dipole moment is a function of its geometry.

When a molecule has just one bond, it’s simple to visualize its dipole moment.

Once we move beyond diatomic molecules, we need to start considering how the sum of the individual dipoles (which are vectors, by the way) is affected by their arrangement in the molecule, i.e. their geometry.

Let’s take BeCl₂, for instance. Linear molecule. Cl is very electronegative [3.2], and Be is very… non-electronegative [1.7].  So each Be–Cl bond should have a large dipole where Cl sucks electrons away from beryllium.

The overall dipole moment of BeCl₂ has been measured, and it is zero.

Why?

What happens here is that the two dipoles act in equal and opposite directions and therefore cancel each other out.

This is a good lesson. A molecule can have large individual dipoles and have no overall dipole moment if the forces act in equal and opposite directions. We will be seeing much more of this!

What about a molecule like BeFCl ?  Will it have a dipole moment? What do you think?

5. H2O and “Bent” Geometry

What about water? As soon as we learn the formula of water is H₂O, the first assumption every kid makes is that it’s linear.

But it ain’t.  How do we know? It’s not like we can go in with a little microscope or something and pull out a tiny protractor to measure the angle. 

Because if it was linear, it would have a dipole moment of zero, like BeCl₂. But we can measure the dipole moment and it’s 2.9 D (in the liquid phase, 1.85 in the gas – reference). That’s a pretty big number! So clearly it ain’t linear. Something else is going on.

We now know that the “something else” is that there are two lone pairs on oxygen in addition to the two bonding pairs in the O-H bonds, and all four electron pairs repel each other. Repulsion is minimized in a tetrahedral geometry, but since the lone pairs take up a little more space the H-O-H angle is about 104.5° instead of 109.5°.

6. BF3 And Trigonal Planar Geometry

Moving up to three bonds around a central atom, let’s start with the trigonal planar geometry, where all three groups are arranged around a central atom with bond angles of 120°.

Let’s take BF₃, for example. What do you think? Dipole moment or no?

Fluorine is very electronegative [4.0], boron is very electropositive [2.0]. So we have three bond dipoles that point from B to F.

Just like with BeCl₂, the vector sum results in cancelation. No dipole moment.

At the bottom of the article are some quizzes on the dipole moments of alkenes (olefins).  

7. NH3 Trigonal Pyramidal Geometry

Next up is trigonal pyramidal geometry, such as that found in NH₃.

Hydrogen has an electronegativity of 2.2, nitrogen has an electronegativity of 3.0, so we can immediately see that the bonds should be polarized toward nitrogen (δ-) and away from hydrogen (δ+).

Nitrogen has four electron pair domains (3 in the N-H bonds, and one lone pair) that are arranged in a roughly tetrahedral geometry to minimize electron pair repulsion.

The three N-H dipole vectors all point up toward nitrogen. When you add these three vectors up you get a molecular dipole moment that points straight up along the z-axis.

8. Tetrahedral Geometry and Dipole Moment

Here’s where things get real: dipole moments in tetrahedral carbon atoms.

Now first off: don’t get stressed out about it. Nobody’s going to ask  you to calculate a dipole moment (unless you happen to be in a computational chemistry class, in which case, why are you even reading this)

However, you could reasonably be expected to identify if one is present or not. That’s important. We’ll walk through this.

The principles are exactly the same as those we just discussed: First, look for bond dipoles, and second, see if they’re balanced.

Let’s first look at methane and carbon tetrachloride.

C-H bonds have a weak dipole polarized towards carbon (carbon is δ-) and C-Cl bonds have a relatively strong dipole polarized toward chlorine (chlorine is δ-)

In both cases the overall dipole moment is zero because all the vectors cancel.

What happens when all bonds *aren’t* identical?

For example, let’s take CCl4 and swap out one of the C-Cl bonds for a C-H bond. (This molecule is called chloroform).

In doing so, we’ve replaced a bond dipole pointing towards chlorine with a bond dipole that points in the opposite direction.

This means that the dipoles will no longer be in balance, and the molecule will have a net dipole moment.

In CHCl₃ the experimentally observed dipole moment is 1.15 D.

9. Dipole Moment And Boiling Point

Dipoles can also have a large effect on the physical properties of a molecule, particularly its liquid state.

In a liquid, oppositely charged partial charges attract each other. These attractions between partial charges are a type of intermolecular force. Generally speaking, the higher the intermolecular forces, the higher the boiling point, since it will take more energy to separate the molecules.

When one of the dipoles is an O-H, N-H, or F-H bond, we call these dipole-dipole interactions hydrogen bonding. Hydrogen bonds are particularly strong intermolecular forces, and are typically about 5-10% of the strength of normal covalent bonds (the hydrogen bond in water has been measured to have a strength of about 5 kcal/mol) . 

Attractions between dipoles that don’t involve hydrogen are typically called, “dipole-dipole” interactions, (or Van Der Waals dipole-dipole interactions) and are weaker than hydrogen bonds.

In the absence of a strong dipole, the only attractive forces present in a molecule are typically London or dispersion forces, which are temporary dipoles that result from momentary imbalances in electron deficiency in a molecule. These weaker forces result in considerably lower boiling points.

Notes

[notes]

Quiz Yourself!

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

Become a  MOC member to see the clickable quiz with answers on the back.

(Advanced) References and Further Reading

1. An Interpretation of the Enhancement of the Water Dipole Moment Due to the Presence of Other Water Molecules
   Daniel D. Kemp and Mark S. Gordon
   The Journal of Physical Chemistry A 2008 112 (22), 4885-4894
   DOI: 10.1021/jp801921f