Chemical Bonding III: Intermolecular Forces and Hydrogen Bonding

Learning Outcomes

Describe intermolecular forces (van der Waals’ forces), based on permanent and induced dipoles, specifically in substances such as $CHCl_3(l)$ , $Br_2(l)$ , and the liquid noble gases.
Describe hydrogen bonding using ammonia ( $NH_3$ ) and water ( $H_2O$ ) as simple examples of molecules containing $N-H$ and $O-H$ groups.
Outline the importance of hydrogen bonding to the physical properties of substances, including ice and water.
Describe, interpret, and/or predict the effect of different types of bonding on the physical properties of substances. This includes:
- Ionic bonding
- Covalent bonding
- Hydrogen bonding
- Other intermolecular interactions
Deduce the type of bonding present from given information.

Overview of Chemical Bonds

Chemical bonds are categorized into two primary types based on whether they act within or between molecules:

Intramolecular Bonds: These are forces that hold the atoms together within a single molecule. They include:
- Ionic / Electrovalent Bond
- Covalent Bond
- Metallic Bond
- Coordinate / Dative Bond
Intermolecular Bonds: These are forces that exist between separate molecules. They are generally weaker than intramolecular bonds and include:
- Hydrogen bonding
- Van der Waals forces

Properties of Ionic Bonding

Electrical Conductivity: Ionic compounds are poor conductors in the solid state. They only conduct electricity when in aqueous solution or molten state.
Thermal Conductivity: They exhibit no thermal conductivity.
Mechanical Properties: They lack ductility and malleability.
Appearance: Most ionic compounds are colorless.
Solubility: They are generally soluble in polar solvents (like water) but insoluble in nonpolar solvents.
Physical State: They have high melting points and usually exist as solids at room temperature.

Properties of Covalent Bonds

Conductivity: Covalent compounds exhibit no electrical or thermal conductivity.
Mechanical Properties: They lack ductility and malleability.
Appearance:
- Water ( $H_2O$ ) is colorless.
- Halogens exhibit distinct colors.
Solubility:
- Nonpolar molecules are insoluble in water but soluble in organic solvents.
- Polar molecules are soluble in water.
Giant Structures: Covalent substances with giant molecular structures (like diamond or silica) are insoluble in all solvents.

Intermolecular Forces of Attraction

These are forces of attraction between molecules, distinct from the covalent bonds within the molecule.
Intermolecular forces are significantly weaker than covalent or ionic bonds.
There are two main types of intermolecular forces:
1. Van der Waals forces
2. Hydrogen bonds

Van der Waals Forces

Van der Waals forces are weak intermolecular forces comprising three distinct types of attractions:

Permanent dipole – permanent dipole attraction
Permanent dipole – induced dipole attraction
London dispersion force (Induced dipole-induced dipole)

General Rules for Van der Waals Forces Strength:

The larger the molecule size, or the larger the number of electrons in a molecule, the larger the Van der Waals forces.
Increased Van der Waals forces result in higher melting points and boiling points.

Three Types of Dipoles

Permanent Dipole – Permanent Dipole Attraction ( $pd/pd$ ):
- This is the force of attraction between the negative end of one polar molecule and the positive end of another polar molecule.
- In diagrams, these are always represented using dotted lines.
- This is generally the strongest type of Van der Waals interaction.
Dipole – Induced Dipole Attraction ( $pd/id$ ):
- This interaction occurs between a polar molecule and a nonpolar molecule.
- The nonpolar molecule is polarized by the proximity of the polar molecule.
Induced Dipole – Induced Dipole Attraction (London Dispersion Forces, $LDF$ ):
- This interaction occurs between nonpolar molecules (such as $O_2$ , $N_2$ , $CO_2$ , and noble gases).
- It is caused by the random movement of electrons within an atom or molecule.
- Example (Argon): Argon has $18$ electrons. While normally nonpolar with a symmetrical electron arrangement, the revolving electrons may occasionally create a higher density at one end, resulting in a temporary dipole.
- This temporary dipole distorts the electron cloud of neighboring atoms, creating induced dipoles that attract each other.
- $LDF$ exists in all molecules.

Relative Strength Hierarchy: $pd/pd > pd/id > LDF$

Factors Affecting Van der Waals Forces

Size of Atoms / Molecules:
- As size increases, the number of electrons and nuclei available to form attraction forces increases.
- This leads to stronger Van der Waals forces, higher boiling points ( $bp$ ), and more energy required to break the bonds.
Shape / Surface Area of Molecules:
- Higher Surface Area: Leads to more area available for interactions, increasing Van der Waals forces.
- Lower Total Surface Area ( $TSA$ ): A more compact structure reduces the surface area available, leading to weaker Van der Waals forces and less energy needed to break bonds.
Polarity of Molecules:
- As polarity increases, the partial charges ( $\delta+$ and $\delta-$ ) increase, strengthening the Van der Waals forces.
- Example: Methane ( $CH_4$ , Non-Polar) has a Boiling Point of $-161.6^{\circ}C$ . Fluoromethane ( $CH_3F$ , Polar) has a Boiling Point of $-78.4^{\circ}C$ because it experiences permanent dipole-permanent dipole attractions which are stronger than London Dispersion Forces.

Hydrogen Bonds

A hydrogen bond is a special, strong type of permanent dipole – permanent dipole attraction.
Conditions for formation:
- A Hydrogen ( $H$ ) atom must be bonded to a small, highly electronegative atom ( $O$ , $N$ , or $F$ ).
- The attracted partner must be a lone pair of electrons on another very electronegative atom ( $O$ , $N$ , or $F$ ).
These occur in polar molecules.

Hydrogen Bonding in Specific Molecules

Hydrogen Fluoride ( $HF$ ): The $H$ atom of one molecule is electrostatically attracted to the $F$ atom of another molecule ( $\delta+$ on $H$ to $\delta-$ on $F$ ).
Water ( $H_2O$ ):
- The $O$ atom has $2$ lone pairs of electrons, each forming one hydrogen bond.
- Each $H$ atom ( $2$ per molecule) also forms a hydrogen bond with neighboring molecules.
- Consequently, each water molecule can form a total of 4 hydrogen bonds.
Ammonia ( $NH_3$ ):
- The $N$ atom has $1$ lone pair, allowing for the formation of $1$ hydrogen bond via the nitrogen.
- Each $H$ atom can also participate in forming hydrogen bonds.

Effects of Hydrogen Bonding

Boiling Point ( $bp$ )

Relative Boiling Points: $NH_3 < HF < H_2O$ .
Reasoning:
- Nitrogen ( $N$ ) is less electronegative than Fluorine ( $F$ ) and Oxygen ( $O$ ).
- Even though $F$ is more electronegative than $O$ , $HF$ has a lower boiling point than $H_2O$ because water can form four hydrogen bonds per molecule ( $2$ lone pairs and $2$ hydrogens), whereas $HF$ forms fewer.
Comparative Boiling Point Data:

Compound	Molecular Weight ( $g\,mol^{-1}$ )	Boiling Point ( $^{\circ}C$ )
$H_2O$	$18.01$	$100$
$HF$	$20.01$	$19.5$
$NH_3$	$17.03$	$-33.3$
$CH_4$	$16.04$	$-161.6$

Comparison Example ( $HF$ vs $HBr$ ): $HF$ has a higher boiling point than $HBr$ because $HF$ possesses hydrogen bonding, while $HBr$ only possesses Van der Waals forces ( $pd/pd$ ). Hydrogen bonds are significantly stronger than standard Van der Waals forces.

Solubility

Hydrogen bonding is common in polar solvents containing $F$ , $O$ , or $N$ (like water).
Compounds capable of forming hydrogen bonds with water dissolve easily.
Solubility increases as the number of hydrogen bonds increases.
Example: Ethanol is more soluble in water than dimethyl ether due to its ability to form more effective hydrogen bonds.

Density (Why Ice Floats on Water)

Liquid State: Water molecules are held close together by weak, transient hydrogen bonds.
Freezing/Solid State: As water reaches its solid state, hydrogen bonds become stable and arrange the molecules farther apart in a tetrahedral cage-like structure.
Result: This arrangement creates empty space between molecules, causing ice to have a lower density than liquid water, allowing it to float. During melting, these bonds break, the cage structure collapses, and molecules move closer together again.

Miscellaneous Bonding Deductions

$HBr$ : Permanent dipole-dipole ( $pd/pd$ ) because it is polar.
$CH_4$ : London Dispersion Forces ( $LDF$ ) because it is non-polar.
$SO_2$ : Permanent dipole-dipole interaction ( $pd/pd$ ) because it is polar.
$CH_3OH$ (Methanol): Hydrogen bonding (due to the presence of an $-OH$ group) and London Dispersion Forces.