Chemistry (9701) | Intermolecular forces, electronegativity and bond properties

A Level Chemistry Study Notes: Intermolecular Forces, Electronegativity, and Bond Properties

Hello future Chemists! This chapter might seem theoretical, but it is one of the most important concepts in A Level Chemistry. Why? Because the forces we discuss here govern everything about a substance's physical properties: its melting point, boiling point, solubility, and even whether it's a solid, liquid, or gas at room temperature. Let's dive into the fascinating world of attraction!

1. Electronegativity: The Electron Tug-of-War (Syllabus 3.1)

Before we look at forces *between* molecules, we need to understand the forces *within* a molecule. This is all about how atoms share (or don't share) electrons, defined by a concept called electronegativity.

1.1 Definition of Electronegativity

Definition: Electronegativity is defined as the power of an atom to attract electrons to itself in a covalent bond.

Analogy: Imagine two atoms sharing a rope (the bond). Electronegativity is how strongly each atom pulls the rope toward itself. The atom with higher electronegativity wins the tug-of-war, pulling the electron density closer to its nucleus.

1.2 Factors Affecting Electronegativity

The strength of this attraction depends on three main factors:

1. Nuclear Charge (Protons): A higher positive charge in the nucleus leads to a stronger attraction for the bonding electrons.
2. Atomic Radius: The smaller the atom, the closer the outer shell electrons are to the nucleus, leading to a stronger pull.
3. Shielding by Inner Shells: Inner shell electrons "shield" the bonding electrons from the pull of the nucleus. More shells mean more shielding, reducing the electronegativity.

Did You Know? Fluorine (F) is the most electronegative element, followed closely by Oxygen (O) and Nitrogen (N). These three elements are crucial for understanding Hydrogen Bonding later!

1.3 Trends in Electronegativity

Understanding the factors above helps explain the trends in the Periodic Table:

• Across a Period (left to right): Electronegativity increases. Why? Nuclear charge increases, but shielding remains relatively constant (same principal quantum number, n). The nucleus pulls harder.
• Down a Group: Electronegativity decreases. Why? Atomic radius increases and shielding increases (more inner shells), making the nucleus's pull on the bonding electrons weaker.

2. Bond Polarity and Dipole Moments (Syllabus 3.6.2)

When two atoms form a covalent bond, the difference in their electronegativity determines the nature of that bond.

2.1 Polar vs. Non-Polar Bonds

When bonding electrons are shared:

• Equal Sharing (Non-Polar): If the electronegativity difference is zero or very small (e.g., \(\text{H}_2\), \(\text{Cl}_2\), or \(\text{C}\)-\(\text{H}\) bonds in methane), the electrons are shared equally.
• Unequal Sharing (Polar): If the electronegativity difference is significant (e.g., \(\text{H}\)-\(\text{Cl}\) or \(\text{H}\)-\(\text{F}\)), the electrons are pulled closer to the more electronegative atom.

In a polar bond, the more electronegative atom gains a partial negative charge (\(\delta-\)) and the less electronegative atom gains a partial positive charge (\(\delta+\)). This separation of charge creates a bond dipole.

2.2 Predicting Bond Type

We can use the difference in Pauling electronegativity values (which are provided if needed) to predict bonding type (Syllabus 3.1.4):

• Small Difference: Covalent bond (non-polar or slightly polar).
• Intermediate Difference: Polar covalent bond.
• Large Difference (typically > 1.7): Ionic bond (electrons are essentially transferred).

2.3 Molecular Polarity and Dipole Moments

Even if a molecule contains polar bonds, the molecule itself might be non-polar if the shape is symmetrical.

• Dipole Moment: This is the overall polarity of a molecule. It is the vector sum of all the individual bond dipoles.
• Example 1: Carbon Dioxide (\(\text{CO}_2\)) is linear (180°). The C=O bonds are polar, but they pull equally and oppositely, cancelling out the dipoles. The molecule is non-polar.
• Example 2: Water (\(\text{H}_2\text{O}\)) is non-linear (bent/V-shaped, 104.5°). The O-H bonds are polar, but because of the bent shape, the dipoles do not cancel. The molecule is polar and has a large overall dipole moment.

3. The Hierarchy of Forces: Strong vs. Weak

It's vital to distinguish between two categories of forces (Syllabus 3.6.4):

1. Intramolecular Forces: These are the strong chemical bonds *within* a molecule or giant structure (Ionic, Covalent, Metallic). Breaking these requires massive energy.
2. Intermolecular Forces (IMFs): These are the weak forces of attraction *between* separate molecules. Breaking these requires far less energy and is what happens when molecular substances melt or boil.

State Rule: In general, ionic, covalent (giant), and metallic bonding are stronger than intermolecular forces.

4. Van der Waals' Forces: The Molecular Glue (Syllabus 3.6.3)

The term Van der Waals' forces is used as a generic term to describe all intermolecular forces between molecular entities other than those due to bond formation (Syllabus 3.6.3a).

4.1 Instantaneous Dipole - Induced Dipole (id-id) Forces

These forces, also known as London Dispersion Forces (LDFs), are present in all molecules (polar and non-polar). They are the weakest forces.

How they work (Step-by-Step):

1. Instantaneous Dipole: At any moment, the electrons around an atom are moving randomly. This movement creates a temporary, instantaneous imbalance of charge—a momentary dipole.
2. Induced Dipole: This momentary dipole in one molecule repels the electrons in a neighboring molecule, causing the neighbor to form its own temporary, induced dipole.
3. Attraction: The instantaneous dipole and the induced dipole attract each other.

Factors Affecting id-id Strength:

• Number of Electrons (\(M_r\)): More electrons mean a larger electron cloud. This cloud is more "squishy" (or polarizable), making it easier to form instantaneous dipoles. Therefore, heavier molecules have stronger id-id forces.
• Surface Area/Shape: Molecules with large surface areas (long, straight chains) can pack closer together, leading to more points of contact and stronger id-id forces compared to branched molecules of similar mass.

4.2 Permanent Dipole - Permanent Dipole (pd-pd) Forces

These forces occur only between molecules that are permanently polar (i.e., they have a permanent dipole moment, like \(\text{HCl}\) or \(\text{SO}_2\)).

How they work: The molecules align themselves so that the partial positive end (\(\delta+\)) of one molecule is attracted to the partial negative end (\(\delta-\)) of a neighbor.

Strength: pd-pd forces are stronger than id-id forces (when comparing molecules of similar \(M_r\)), because the dipoles are permanent, not just momentary.

5. Hydrogen Bonding: The Strongest Intermolecular Force (Syllabus 3.6.1)

Hydrogen bonding is described as a special case of permanent dipole-permanent dipole forces where hydrogen is bonded to a highly electronegative atom (Syllabus 3.6.3c).

5.1 Conditions for Hydrogen Bonding

A hydrogen bond can only form if the molecule contains one of the following highly polar bonds (Syllabus 3.6.1a):

• \(\text{N}\)-\(\text{H}\) (e.g., Ammonia, \(\text{NH}_3\))
• \(\text{O}\)-\(\text{H}\) (e.g., Water, \(\text{H}_2\text{O}\), and Alcohols)
• \(\text{F}\)-\(\text{H}\) (e.g., Hydrogen Fluoride, \(\text{HF}\))

Memory Aid: Remember the abbreviation FON (Fluorine, Oxygen, Nitrogen).

5.2 How Hydrogen Bonds Form

When H is bonded to N, O, or F, the electron density is pulled so far away from the small hydrogen atom that the \(\text{H}\) nucleus (just a proton) is almost entirely exposed. This leaves the \(\text{H}\) with a very high \(\delta+\) charge, which is then strongly attracted to a lone pair of electrons on a neighboring N, O, or F atom.

Analogy: H-bonds are like miniature, super-strong magnets holding the molecules together. Since H is so tiny, the attraction is very close-range and extremely effective.

6. The Anomalous Properties of Water (Syllabus 3.6.1b)

The extensive network of strong hydrogen bonds between water molecules gives \(\text{H}_2\text{O}\) several unusual properties that are crucial for life on Earth. You must be able to explain these in terms of H-bonding.

6.1 Relatively High Melting and Boiling Points

If water only had weaker id-id forces (like its similar-sized neighbor, \(\text{H}_2\text{S}\)), it would be a gas at room temperature. However, to melt ice or boil liquid water, a large amount of energy must be supplied to break the extensive network of hydrogen bonds. This leads to anomalously high MP and BP.

6.2 Relatively High Surface Tension

Surface Tension is the energy required to increase the surface area of a liquid. Water molecules at the surface are pulled inward by the strong cohesive forces (H-bonds) from molecules below them, making the surface very stable and allowing light objects (or insects) to rest upon it.

6.3 Density of Ice Compared to Liquid Water

This is arguably the most famous anomaly:

1. In Liquid Water: H-bonds constantly form and break, allowing the molecules to pack relatively close together, maximizing density.
2. In Ice (Solid): As water freezes, the H-bonds lock the molecules into a highly ordered, tetrahedral crystalline lattice structure. This specific arrangement creates large amounts of empty space (cages) between the molecules.

Because of this open cage structure, ice is less dense than liquid water, causing it to float. This is critical as it prevents large bodies of water from freezing solid from the bottom up.

7. Applying Intermolecular Forces to Physical Properties (Syllabus 4.2.2)

The type and strength of the forces determine a substance's physical characteristics:

7.1 Melting Point and Boiling Point

• Simple Molecular Structures (e.g., \(\text{I}_2\), \(\text{H}_2\text{O}\), \(\text{CO}_2\)): Have only weak intermolecular forces between molecules. Only these weak forces need to be overcome to melt or boil the substance. Therefore, they have low MP and BP.
• Giant Structures (Ionic, Metallic, Giant Covalent): The particles are held together by strong intramolecular forces (ionic bonds, covalent bonds, or metallic bonds). Overcoming these strong bonds requires a huge amount of energy, resulting in very high MP and BP.

7.2 Solubility

A simple rule of thumb applies: "Like dissolves like."

• Polar Solutes (e.g., Ethanol) dissolve well in Polar Solvents (e.g., Water) because they can form strong pd-pd forces or H-bonds with the solvent molecules.
• Non-Polar Solutes (e.g., Oil, \(\text{I}_2\)) dissolve well in Non-Polar Solvents (e.g., Hexane) because they can form new id-id forces of comparable strength to those already present.