Lecture 9 – Intermolecular Forces and Phase Transitions

Lecture Preview

• What are intermolecular forces?
• How can intramolecular interactions explain intermolecular forces?
• What types of IMFs are there and how are they different from each other?
• How can IMFs help us predict states of matter (phases)?

Lecture Content (Part 1)

Welcome to Chapter 3! In the previous chapter (Chapter 2 – Intramolecular Interactions and Representations), we talked about interactions within the molecule itself. Those intramolecular interactions mostly amount to the following concepts:

• Electron configurations and periodic table trends
• Different types of bonding (ionic, covalent, and metallic)
• Covalent bonding (Lewis structures, VSEPR, molecular geometry, and resonance)
• Visualization of different molecules using online resources

In this new chapter, we zoom out our perspective. Rather than looking at what happens within a molecule, we look at what happens between molecules. In other words, we will focus on intermolecular interactions.

Of course, intramolecular interactions usually affect intermolecular interactions. Consequently, we will still be discussing Chapter 2 material in this chapter.

So, we will present Chapter 3 materials using the following outline.

1. Firstly, we discuss different types of intermolecular interactions (intermolecular forces, or “IMF” for short). Depending on the IMFs, compounds can be in different states (gases, liquids, and solids).
2. Secondly, we will discuss how gases behave.
3. Thirdly, we will discuss how liquids and solids behave.
4. Lastly, we will discuss how solutions and colloids behave.

Using this outline, we will be exploring intermolecular forces this lecture.

The Source of Intermolecular Forces

Let’s first talk about what intermolecular forces are.

An intermolecular force (IMF) is the force that mediates interaction between molecules, including the electromagnetic forces of attraction or repulsion which act between atoms and other types of neighbouring particles (e.g. atoms or ions).

Intermolecular forces are weak relative to intramolecular forces – the forces which hold a molecule together. For example, the covalent bond, involving sharing electron pairs between atoms, is much stronger than the forces present between neighboring molecules.

So, here comes the question: What causes intermolecular forces? What causes molecules to interact with each other? The answer lies within the intramolecular compositions of a molecule.

Dipole Moments

Let’s focus back to what happens within a molecule. A molecule’s intramolecular properties can be described by electron movements, as discussed in Lectures 6 (Chemical Bonding and Electron Movements) and 7 (Lewis Structures and Molecular Geometry). Using this as a starting point, let’s considering the following logic.

1. Suppose an atom of a molecule is exceptionally good at retaining electrons around itself compared to the surrounding atoms. In other words, that atom has high electronegativity $\chi$.
2. As a result, that atom would have more electrons surrounding it compared to the surrounding atoms.
3. That atom, in the end, will obtain a partial negative charge $\delta^-$, while the surrounding atoms will obtain a partial positive charge $\delta^+$.
4. Now, let’s zoom out to look at the molecule – one part of the molecule will be slightly negative $\delta^-$, while another part of the molecule will be slightly positive $\delta^+$.
5. So, if we put a lot of the same molecules together, the molecules should align themselves such that the $\delta^-$ and $\delta^+$ are close together between 2 molecules.
6. That attraction between opposite partial charges ($\delta^-$ and $\delta^+$) is what we call an IMF.

Let’s look at an example visualization below.

In our logic above, we stumbled upon a new concept called the dipole moment.

The dipole moment is a measure of the separation of positive and negative electrical charges within a system.

We can actually express a dipole moment on a molecule by drawing an arrow pointing from the $\delta^+$ atom to the $\delta^-$ atom.

We can also add up individual dipole moments to get a net dipole moment, since a dipole moment is a vector (magnitude + direction).

Using the same molecule as the visualization example, we can draw out individual dipole moments and calculate the net dipole moment of this molecule.

As we can see, the water molecule has a net dipole moment pointing from the hydrogen atoms to the oxygen atom. This concept explains why molecules align in a certain way.

Polarity

Individual dipole moments and the net dipole moment both help us understand the concept of polarity – a property of a molecule.

Polarity is a separation of electric charge leading to a molecule or its chemical groups having a dipole moment, with a negatively charged end and a positively charged end.

• Bond polarity specifically refers to the polarity of the covalent bond of interest. It is caused by electronegativity differences $\Delta\chi$ between 2 atoms.
• Molecular polarity specifically refers to the polarity of the entire molecule. It can be determined by calculating the net dipole moment.

We can have a non-polar molecule (molecular polarity) connected by polar covalent bonds (bond polarity).

Just to clarify, polarity and dipole moments are different concepts.

• An individual dipole moment is a vector pointing from $\delta^+$ to $\delta^-$. It is caused by bond polarity.
• A net dipole moment is a vector that can be calculated by adding up all the individual dipole moments in the same molecule.
• Bond polarity is a property of a covalent bond that is caused by electronegativity differences between 2 atoms.
• Molecular polarity is a property of the entire molecule that can be determined by looking at the net dipole moment.

To summarize, $\Delta\chi$ between 2 atoms causes bond polarity, which causes individual dipole moments that can add up to a net dipole moment, which can help determine the molecular polarity.

Here, the concept of “bond polarity” should ring a faint bell in our minds. Recall in Lecture 6 (Chemical Bonding and Electron Movements), we talked about how electronegativity differences can help us predict bond formation.

We will look at some examples below.

If there are still confusions, don’t worry, there are plenty of examples and exercises coming up at the end of this lecture and in the worksheet.

A concept we introduced in the examples that would significantly help us in determining molecular polarity is the symmetry of a molecule. If the molecule is symmetrical, individual dipole moments can cancel each other out.

Types of Intermolecular Forces

We will introduce those forces from weakest to strongest. We can also predict IMFs by looking at the polarity of the molecules involved. Understandably, IMFs between nonpolar molecules would be weak, while IMFs between polar molecules would be strong.

So, the types of IMFs will be roughly categorized by molecules of different polarity.

• Nonpolar molecules and nonpolar molecules
• Nonpolar molecules and polar molecules
• Polar molecules and polar molecules

Dispersion Forces

This IMF is found between nonpolar molecules. Since nonpolar molecules don’t have net dipole moments, how can molecules attract each other? The answer lies in spontaneous dipoles.

Spontaneous individual dipole moments can spawn out of pure chance by electron movements – the electrons just happen to move into a position such that a separation of charge is present. Consequently, molecules can develop spontaneous net dipoles.

We can imagine spontaneous dipoles as “accidental” dipoles. Spontaneous dipoles also go by the term “instantaneous dipoles”.

Now, we have a dipole for one molecule, but IMFs need 2 molecules with dipoles to work. Here, we introduce another concept – induced dipoles.

Induced dipoles in a molecule (1) are net dipole moments that are induced by molecules (2) with existing dipole moments close to it (1) because existing dipole moments (2) have partial charges, aligning the original molecule’s (1) electrons to have favorable interactions with the other molecules (2).

This spontaneous dipole – induced dipole (molecule A has a spontaneous dipole, while molecule B has an induced dipole) IMF is called dispersion forces (also known as London dispersion forces in honor of the scientist Fritz London, not the city). We’ll abbreviate it as “LDF”. Let’s look at a visualization of LDFs.

As we can see, because this IMF relies on spontaneous movements, which implies that:

• Dispersion forces can spontaneously appear and spontaneously disappear.
• Dispersion forces are present in every type of molecule, regardless of polarity.
• Dispersion forces are very weak.

The following molecules have purely LDFs between themselves because they are all nonpolar.

$CH_4$

$C_2H_6$

$Ne$

$CF_4$

$SF_6$

$C_8H_{18}$

The following molecules experience predominantly LDFs between them because they are largely nonpolar. Those molecules do experience other IMFs but LDFs are the main stabilizing IMF because of sheer size. (Note: lone pairs are not included in the Lewis structures for the structures below)

Beta-carotene ($C_{40}H_{56}$)

Tristearin ($C_{57}H_{110}O_6$)

Vitamin E ($C_{29}H_{50}O_2$)

Corticosterone ($C_{21}H_{30}O_4$)

Again, these are biochemical molecules with important functions.

• Beta-carotene is a precursor to vitamin A, which has functions including growth during embryo development, maintaining the immune system, and healthy vision.

• Tristearin is a type of triacylglycerol – a main constituent of human body fat. It is heavily involved in efficient energy storage.

• Vitamin E is a group of compounds that help target reactive oxygen species (ROS). Recall that ROS contains radicals, which are highly reactive and can cause damage to our DNA.

• Corticosterone mainly exist as an intermediate for the synthesis of aldosterone – a compound playing a central role in the homeostatic regulation of blood pressure, plasma sodium (Na+), and potassium (K+) levels.

As we can see, despite having oxygen (a fairly electronegative atom) in the structure, possibly introducing dipole moments, the main IMF of those molecules are LDFs because of the sheer size of the nonpolar regions. Of course, other IMFs exist in those molecules, such as dipole-dipole interactions and hydrogen bonds, which we will discuss soon.

Dipole – Induced Dipole Interactions

This type of interaction is tailored to interactions between polar and nonpolar molecules. We will abbreviate it as “D-ID” for this class.

The “dipole” part of this IMF is easy to understand – it comes from the polar molecule, which itself has a dipole moment. The “induced dipole” part of this IMF should also be easy to understand. Since we have an existing dipole from the polar molecule, it can influence a nonpolar molecule next to it to have an electron distribution such that there is a dipole. In other words:

Existing dipoles can induce dipoles from nonpolar molecules.

Here are some examples where D-ID interactions are the predominant IMF. Note that in all of the examples below, LDFs are also present. However, LDFs are not the predominant IMF here.

D: Ethanol ($C_2H_6O$)
ID: Iodine ($I_2$)

D: Chloroform ($CHCl_3$)
ID: Methane ($CH_4$)

D: Acetone ($C_3H_6O$)
ID: Benzene ($C_6H_6$)

D: Water ($H_2O$)
ID: Oxygen ($O_2$)

For the examples above, try to answer the following questions:

1. Which molecule has existing dipoles? (i.e., which molecule is polar?)
2. Which molecule has induced dipoles? (i.e., which molecule is nonpolar?)
3. For the molecule with existing dipoles, what is the direction of its net dipole moment?
4. For the molecule with induced dipoles, how is it nonpolar?
   1. Is it because it has no obvious individual dipole moments, or…
   2. Is it because molecular symmetry canceled out the individual dipole moments?

Dipole – Dipole Interactions

This type of interaction is tailored to interactions between polar molecules. We will abbreviate it as “D-D” for this class. Here are some examples with D-D interactions being the predominant IMF. Note that for all of the examples below, they all contain LDF and D-ID interactions.

DMSO ($C_2H_6OS$)

Hydrochloric Acid ($HCl$)

Formaldehyde ($CH_2O$)

Propanal ($C_3H_6O$)

Similarly, for those molecules, try to answer the following questions.

1. Is the molecule polar? (The answer should be yes for all of them)
2. If the molecule is polar, where is the direction of the net dipole moment?

Hydrogen Bonding

Hydrogen bonding is important. Hydrogen bonding is very important. Hydrogen bonding is extremely important.

Hydrogen bonding is considered as a special case of dipole-dipole (D-D) interactions. We may abbreviate it as “H-bonding”.

Hydrogen bonding should look like the following:

$$R-Dn\text{-}H \cdots Ac-R$$

where:

• $Dn$ denotes hydrogen bond donator, because they contain the hydrogen
• $H$ denotes the hydrogen atom
• $\cdots$ denotes the hydrogen bond IMF
• $Ac$ denotes the hydrogen bond acceptor
• $R$ denotes unimportant parts of the molecule

$Dn$ and $Ac$ need to be the elements “N”, “O”, and “F” for this dipole-dipole interaction to be considered as hydrogen bonding.

The reason that it is important is because:

1. It is stronger than dipole-dipole interactions because it involves N, O, F – the most electronegative elements – and H.
2. It is everywhere in biochemistry, considering water $H_2O$ is the universal solvent.
3. Hydrogen bond is an IMF. Hydrogen bonds are different from ionic/covalent/metallic bonds.

Using this definition, we can look at the following examples with H-bonding as the predominant IMF. (As usual, they all contain LDFs and D-IDs)

Dn: Water ($H_2O$)
Ac: Water ($H_2O$)

Dn: Ammonia ($NH_3$)
Ac: Ammonia ($NH_3$)

Dn: Ethanol ($C_2H_6O$)
Ac: Ethanol ($C_2H_6O$)

Dn: Ethanol ($C_2H_6O$)
Ac: Formaldehyde ($CH_2O$)

Because formaldehyde does not have hydrogen atoms directly attached to N, O, or F, it can only be a hydrogen bond acceptor. As a result, it needs a hydrogen bond donator (Dn-H where Dn = N, O, or F).

Since hydrogen bonding is very important, we have to talk about medical/biochemical applications of hydrogen bonding. Also, the molecules get pretty big in those examples, so a molecule can have both hydrogen bond donators and acceptors. For example, we can look at carboxylic acid ($HCO_2H$).

For the oxygen at the top, since it does not have a hydrogen atom directly attached to it, it can only be a hydrogen bond acceptor. For the oxygen to the right, it can be both a donator and an acceptor.

Note: In the graphs below, H-bonds are in dashed lines, hydrogen bond acceptors are highlighted, and lone pairs are not included.

Adenine-Thymine (AT) base pair in DNA

Cytosine-Guanine (CG) base pair in DNA

Parallel beta-sheets in proteins

Antiparallel beta-sheets in proteins

• For the first 2 examples, we have AT and CG base pairings – the golden rule in DNA structure. Since AT has 2 hydrogen bonds while CG has 3 hydrogen bonds, CG pairings are typically stronger than AT pairings.

• For the next 2 examples, we have secondary protein structures called beta sheets, which are important in protein conformations (shapes). Similarly, they are stabilized by hydrogen bonds.

Other examples include things like structures dissolving in water, like acetone dissolving in water, aspirin dissolving in water, ammonia dissolving in water, etc.

Ion – Dipole Interactions

Ion-dipole interactions are for ions and polar molecules. They are even stronger than hydrogen bonds, because rather than having partial charges attracting each other, actual charges are involved. We’ll abbreviate this as “Ion-D”. The most textbook example of this is table salt ($NaCl$) dissolving in water.

$$NaCl (s) \rightarrow Na^+ (aq) + Cl^- (aq)$$

• Since $Na^+$ is positively charged, water will align in a way such that the oxygen (partial negative) will point towards $Na^+$.
• Since $Cl^-$ is negatively charged, water will align in a way such that the hydrogen (partial positive) will point towards $Cl^-$.

As long as we see ions and polar molecules, we can usually assume that ion-dipole interactions are predominant.

We can keep going on and have ion-ion interactions, but at that point, it’s the same thing as ionic bonds.

Summary

We introduced 5 types of IMFs – 4 main IMFs and 1 special IMF – listed in increasing strength below.

• Dispersion forces (LDF)
• Dipole – induced dipole (D-ID)
• Dipole – dipole (D-D)
  • Special case: Hydrogen bond (H-bond)
• Ion – dipole (Ion-D)

Ideally, we should memorize this by heart. Also, here are some terms we should clarify.

• The strongest and predominant IMF could be different in a sample. For example, let’s consider the molecule tristearin – a type of triacylglycerol.

Of course, the exact boundary of “overwhelmingness” is very blurry. So, as long as we can identify and acknowledge that “oh, there should be a strong LDF presence”, we should be fine. Also, if LDF is the predominant IMF in a molecule in a quiz/test problem, it should be very obvious.

Intermolecular Forces in Phase Transitions

We will talk more about the thermodynamics behind phase transitions in Unit 2. For now, we will be brief about it.

Review: Phases in Lecture 3

Recall back in lecture 3 (Phases, Properties, Particles, and the Periodic Table), we have some nice visualizations of what solids, liquids, and gases look like in a microscopic perspective.

Attribution: Julio Miguel A Enriquez and Monica Muñoz, CC BY-SA 4.0 https://creativecommons.org/licenses/by-sa/4.0, via Wikimedia Commons

• A solid has a stable, definite shape, and a definite volume. Solids can only change their shape by an outside force, as when broken or cut.
• A liquid is a nearly incompressible fluid that conforms to the shape of its container but retains a (nearly) constant volume independent of pressure.
• A gas is a compressible fluid. Not only will a gas conform to the shape of its container but it will also expand to fill the container.

Also, recall the graph regarding phase transitions.

• Melting: solid to liquid
• Freezing: liquid to solid
• Vaporization: liquid to gas
• Condensation: gas to liquid
• Sublimation: solid to gas
• Deposition: gas to solid

Melting point is the temperature where solid becomes liquid or liquid becomes solid. Boiling point is the temperature where liquid becomes gas or gas becomes liquid.

We will focus on the following concepts in this section.

• Comparatively, how do IMFs determine the state of matter? In other words, how do IMFs determine substances to be in solids, liquids, and gases?
• Comparatively, how do IMFs influence the melting and boiling points of substances?

Before we go predicting phases, here are the basics.

Compounds in gas phase usually have comparatively weak IMFs. Weak IMFs mean low levels of intermolecular attraction, allowing gas molecules to flow freely.

Compounds in liquid phase usually have comparatively moderate IMFs. Moderate IMFs mean moderate levels of intermolecular attraction, allow liquid molecules to flow freely with some restrictions (such as not being able to be too far away from another molecule).

Compounds in solid phase usually have comparatively strong IMFs. Strong IMFs mean high levels of intermolecular attraction, allowing solid molecules to stick together tightly and not move a lot.

Similarly, if a compound has strong IMFs, it usually has a high melting and boiling point. If a compound has weak IMFs, it usually has a low melting and boiling point.

Predicting Phases

It’s impossible to predict the absolute melting or boiling points just by looking at the molecular structure alone. However, it is possible to compare the molecule of interest with other known compounds and see if the molecule will have higher or lower melting or boiling points.

We have 2 factors we need to consider when comparing.

1. Given similar types of IMFs present between molecules, what is the amount of IMFs present in the molecule?
2. Given similar amounts of IMFs present between molecules, what are the strengths of the individual IMFs?

Let’s look at 2 examples – $CH_4 \text{vs.} C_2H_6$ and $CH_4 \text{vs.} NH_3$

• For $CH_4 \text{vs.} C_2H_6$, we have LDFs for both compounds. However, $C_2H_6$ has bigger surface area, allowing more LDFs present. Thus, $C_2H_6$ should have a higher melting and boiling point than $CH_4$.

• For $CH_4 \text{vs.} NH_3$, we have similar amounts of IMFs present because the molecules are of similar sizes. However, $NH_3$ has H-bonding while $CH_4$ only has LDFs. Thus, $NH_3$ should have a higher melting and boiling point than $CH_4$.

Comparing Substances

Let’s look at groups of compounds and compare them together.

Example 1 focuses on the amount of IMFs. Example 2, on the other hand, will focus on the types of IMFs, given similar amounts of IMFs present. We will, of course, encounter exceptions along the way.

Let’s also look at how hydrogen bonding can influence melting and boiling points. In the previous example, we move from left to right in the periodic table. In this example, we will move from up to down in the periodic table.

Lecture Content (Part 2)

Let’s practice. There will be 5 groups of molecules below. Each group contains 1 reference molecule and 2 other molecules. Our mission is to complete the following tasks.

• Task 1: Identify the predominant IMF in the molecules, including the reference molecule.

• Task 2: Identify the amount of IMF that can be present in the molecules compared to the reference molecule.
  • Use words like “more”, “less”, or “roughly equal”.

• Task 3: Predict if the molecules in question will have a higher or lower melting and boiling point compared to the reference molecule.
  • Use words like “higher”, “lower”, or “roughly equal”.

Similarly, feel free to visualize the molecules in 3D using MolView. This is also a great opportunity to get ourselves more familiar with Lewis structures, especially resonance structures.