This episode recaps the five types of intermolecular forces, discusses factors influencing the strength of force and provides examples.
This episode recaps the five types of intermolecular forces, discusses factors influencing the strength of force and provides examples. It starts with the weakest IMF, London Dispersion Forces (2:05) and introduces polarizability (2:31). In order of increasing strength we recap dipole-induced dipole interaction (5:00), dipole-dipole interaction (6:55), hydrogen bonding (7:35) and ion-dipole interaction (8:38).
Question: (9:30) Which isomer of pentane has the higher boiling point - n-Pentane (C5H12) or neopentane (C(CH3)4)?
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap Intermolecular Forces.
Lets Zoom Out:
Unit 3 - Intermolecular Forces and Properties
Topic - 3.1
Big idea - Structure and Properties
Last week we started the episode pointing out that water is OH so special! Today, we continue along the same lines and will actually learn that water - or more specifically hydrogen - even has its own intermolecular force! Pf! So let’s recap Intermolecular Forces.
Let’s zoom in:
Intermolecular forces are the forces BETWEEN the molecules. They are in contrast to INTRAmolecular forces - the forces within a molecule, like with the attraction between the nuclei and shared valence electrons in the formation of covalent bonds.
We can look at boiling points to determine the strength of intermolecular forces. As we have discussed before, to transition from a liquid to a gas we have to overcome all intermolecular forces. And that takes different amounts of energy depending on the strength of intermolecular forces.
So let’s dive right in and compare the boiling points of nonpolar molecules, the halogens:
Fluorine has a boiling point of -188 degrees Celsius and Chlorine -35 degrees Celsius. They are both gases at room temperature. Bromine has a boiling point of 58 degrees Celsius and is therefore a liquid at room temperature. And Iodine has a boiling point of 184 degrees Celsius - and a melting point of 114 degrees Celsius and is solid at room temperature. But why the difference? The first, and weakest type of intermolecular forces, are London dispersion forces. And no, you don’t have to say that with a british accent - they are actually named after the German physicist Fritz London. London Dispersion Forces, or LDFs, are the result of the attraction between two nonpolar molecules with temporary dipoles which result from the polarizability of the electron cloud. WHAT? Let’s slow down.
Think about Helium with its two valence electrons. If we would take snapchats - uhm I mean - snapshots of helium, the two electrons wouldn’t always be symmetrically distributed. In some snapshots both electrons might be on one side. This leads to an instantaneous dipole. For a very, very, very short amount of time - the helium has a positive side and a negative side. It’s electron cloud is said to be distorted. The tendency to have this distorted electron cloud is called polarizability. This instantaneous dipole can now affect neighboring molecules and induce a dipole, leading to an attraction between two molecules, even when they are both nonpolar! Tada - a force between molecules! The “strength” - and I am saying this with quotations because it is a very weak force - is determined by different factors. First, the number of electrons: the more electrons in an atom, the easier it is to distort. Second, the size of the atom or molecule: the heavier, the more dispersion force. Third, the shape of the molecule: a greater surface area leads to more intermolecular forces.
So let’s go back to our halogens. Of the four, fluorine has the lowest number of electrons and is the lightest. Therefore it experiences the weakest LDFs and has the lowest boiling point. Iodine has the highest number of electrons and is the heaviest, which increases the London Dispersion forces and the boiling point.
Halogens are nonpolar molecules. What if we mix polar and nonpolar molecules? Similar to how the electrons induced a dipole on a neighboring atom or molecule, polar molecules can also induce a dipole and lead to - wait for it - dipole-induced dipole interactions. Quick recap about dipoles: Due to difference in electronegativity between bonded atoms, dipoles are molecules that have one or more areas that are partially negative and one or more areas that are partially positive - D- I - pole - TWO poles. Be careful, it is NOT the same as a polar molecule. Polar molecules have one more attribute: their molecular geometry doesn’t cancel out. So, all polar molecules have dipoles but not all dipoles are polar. Just like how all squares are rectangles but not all rectangles are squares… you get it.
Okay, back to the dipole interaction: let’s say we have a polar molecule and a nonpolar molecule. The polar molecule can induce a dipole in the nonpolar molecule by distorting its electron cloud. To simplify it, we can say the partially negative part of the polar molecule can repel the electrons of the nonpolar molecule, shifting them away and therefore creating an induced dipole with the positive side towards the negative charge of the polar molecule. Due to this mechanism, dipole-induced dipole interactions are always attractive. Their strength is determined by the magnitude of the dipole of the polar molecule and the polarizability of the nonpolar molecule. This intermolecular force is essential for survival. And no, I don’t mean for your survival of AP Chemistry, but for aquatic organisms. Dipole induced dipoles actually increase the solubility of the nonpolar oxygen molecule in polar water molecules.
The second type of dipole interaction are the dipole-dipole interactions between two polar molecules. The strength of the interaction between two polar molecules depends on the magnitude of the dipoles as well as the orientation.
Dipole-dipole interactions and LDFs are additive: substances that are experiencing dipole-dipole interactions are also experiencing London Dispersion Forces. Dipole-dipole interaction can be between molecules of the same chemical species - for example within NF3, or between molecules of two different species - for example between NF3 and H2S.
Now let’s take a closer look at water, because… it's special, you know. The water molecule is a polar molecule, with partially negative oxygen and a partially positive charge on hydrogen. Due to the large difference in electronegativity between hydrogen and oxygen, this partial charge has a relatively large magnitude. This makes the dipole-dipole interactions especially strong. So strong, that this type of intermolecular force has its own name: hydrogen bonding. The bonding part is a bit confusing, since it is NOT a chemical bond, but it is what it is. Sigh. Whenever small hydrogen shares its one valence electron with a highly electronegative atom - fluorine, oxygen, nitrogen - the partially positive hydrogen can interact with the lone pairs of the fluorine/oxygen/nitrogen atoms of the neighboring molecules - and have some FON! Get it? F , O, N, fun…. OK. Due to hydrogen bonding, water, ammonia and hydrogen fluoride have unusually high boiling points.
Last, but not least - actually quite the opposite: Ion-dipole forces. These are the strongest intermolecular forces between an ion and a dipole, for example sodium chloride in water. Because the ion has a full charge, and not just a partial charge, they tend to be stronger than dipole-dipole interactions. And, as usual, the greater the charge on the ion, the stronger the force between ion and dipole. Praise be to Coulomb!
To recap……
There are five types of intermolecular forces: London Dispersion forces, dipole-induced dipole interactions, dipole-dipole interactions, hydrogen bonding and ion-dipole interactions. These intermolecular forces determine physical properties of substances, like the boiling point.
Coming up next on the APsolute RecAP Chemistry Edition: Kinetic Molecular Theory
Today’s Question of the day is about the boiling point of isomers of pentane.
Question: Which isomer of pentane has the higher boiling point - n-Pentane (C5H12) or neopentane (C(CH3)4)?