## 3.9 Intermolecular Forces (Forces Between Molecules) Intermolecular forces (IMFs) are the attractive or repulsive forces that exist *between* molecules, particularly covalent molecules. These forces are significantly weaker than the intramolecular forces (like covalent or ionic bonds) that hold atoms together *within* a molecule. The physical properties of substances, such as melting point, boiling point, and density, are largely determined by the strength of their intermolecular forces. ### 3.9.1 Types of Intermolecular Forces There are three primary types of intermolecular forces: 1. **Permanent Dipole-Dipole Forces** 2. **Instantaneous Dipole-Induced Dipole Forces** (also known as London Dispersion Forces or van der Waals forces) 3. **Hydrogen Bonding** #### 1. Permanent Dipole-Dipole Forces These forces occur between **polar molecules**. A molecule is polar if it has a net dipole moment, which arises from an unequal sharing of electrons due to differences in electronegativity. This creates a permanent separation of charge, where one end of the molecule has a partial positive charge ($\delta^+$) and the other end has a partial negative charge ($\delta^-$). * **Example:** In hydrogen chloride ($HCl$), chlorine is more electronegative than hydrogen, creating a dipole: $H^{\delta+}-Cl^{\delta-}$. The positive end of one $HCl$ molecule is attracted to the negative end of a neighboring molecule. $$ H^{\delta+}-Cl^{\delta-} \cdot\cdot\cdot H^{\delta+}-Cl^{\delta-} $$ * **Non-Polar Molecules:** Some molecules contain polar bonds but are non-polar overall because their symmetrical shape causes the individual bond dipoles to cancel out. A key example is carbon tetrachloride ($CCl_4$). * **Effect on Properties:** The stronger the dipole-dipole forces, the more energy is required to separate the molecules. This leads to higher melting points, boiling points, heats of vaporization, and heats of sublimation. #### 2. Instantaneous Dipole-Induced Dipole Forces These are the weakest type of intermolecular forces and are present between **all molecules**, including non-polar ones. * **Mechanism:** Electrons in an atom or molecule are in constant, random motion. At any given moment, the electron distribution can be uneven, creating a temporary, *instantaneous dipole*. This temporary dipole can then *induce* a temporary dipole in a neighboring molecule, resulting in a weak, short-lived attraction between them. * **Factors Affecting Strength:** The strength of these forces (also known as London Dispersion Forces or LDF) increases with: * **Molecular size/mass:** Larger molecules have more electrons and a larger surface area. * **Number of electrons:** A greater number of electrons leads to a more polarizable electron cloud, making it easier to form temporary dipoles. **Example: Boiling Points of Alkanes** As the size of alkane molecules increases, the van der Waals forces become stronger, leading to higher boiling points. | Name | Molecular Formula | Boiling Point (°C) | | :------ | :-------------------------- | :-----------------: | | Methane | $CH_4$ | -164 | | Ethane | $C_2H_6$ | -89 | | Propane | $C_3H_8$ | -42 | | Butane | $C_4H_{10}$ | -1 | *Table 3.5: Boiling point of first four members of alkane.* For instance, propane ($C_3H_8$, $M_r = 44$) has a higher boiling point than ethane ($C_2H_6$, $M_r = 30$) because its larger size and greater number of electrons result in stronger instantaneous dipole-induced dipole forces. #### 3. Hydrogen Bonding This is a special, unusually strong type of dipole-dipole interaction. It is not a true chemical bond. * **Conditions:** Hydrogen bonding occurs when a hydrogen atom is covalently bonded to a highly electronegative atom with lone pairs of electrons, specifically **Nitrogen (N)**, **Oxygen (O)**, or **Fluorine (F)**. * **Mechanism:** The highly electronegative atom pulls the bonding electrons away from the hydrogen atom, leaving a highly exposed, partially positive hydrogen nucleus ($\delta^+$). This exposed proton is then strongly attracted to a lone pair of electrons on a neighboring N, O, or F atom. * **Effect on Properties:** * **High Boiling Points:** Substances with hydrogen bonding, like water ($H_2O$), have significantly higher boiling points than similar molecules without it. * **Solubility:** Covalent compounds that can form hydrogen bonds with water (e.g., ammonia, $NH_3$) tend to be soluble in water. ### 3.9.2 Peculiar Behaviour of Hydrogen Bonding #### Ice is less dense than water. Normally, the solid state of a substance is denser than its liquid state. Water is a notable exception due to hydrogen bonding. * In ice, water molecules arrange themselves into a rigid, three-dimensional crystalline lattice. * Each oxygen atom is tetrahedrally surrounded by hydrogen atoms, forming a network held together by relatively long hydrogen bonds. * This creates an "open" structure where molecules are held farther apart than they are in the more disordered liquid state. * Because the same mass of water occupies a larger volume as ice, the density of ice is less than that of liquid water, causing it to float. ### The Strength of Intermolecular Forces The general order of strength for intermolecular forces is: $$ \text{Hydrogen bonding} > \text{Permanent dipole-dipole} > \text{Instantaneous dipole-induced dipole} $$ This trend is reflected in the boiling points of different compounds: | Molecule | $CH_4$ | $HCl$ | $H_2O$ | | :--------------------- | :------------------------------------------------------- | :------------------------------------------------------- | :---------------------------- | | **Boiling Point (°C)** | -162 | -85 | 100 | | **Intermolecular Forces** | Instantaneous dipole
(van der Waals) | Permanent dipole-dipole
(and van der Waals) | Hydrogen bonding
(and others) | *Table 3.6: Comparison of boiling points of different compounds.* The boiling points of the hydrides of elements in Groups 14-17 also illustrate this trend. While boiling points generally increase down a group due to stronger van der Waals forces, the first members of Groups 15, 16, and 17 ($NH_3$, $H_2O$, $HF$) have anomalously high boiling points due to hydrogen bonding. For more on group trends, see . ### 3.9.3 Comparison of Bonding Strength It is crucial to distinguish between intermolecular forces and intramolecular forces (chemical bonds). * **Intramolecular forces** (Ionic, Covalent, Metallic bonds) are very strong forces *within* molecules or crystal lattices. * **Intermolecular forces** are much weaker forces *between* separate molecules. Breaking a substance down during melting or boiling involves overcoming the intermolecular forces, *not* the strong covalent bonds inside the molecules. | Substance | Bonding / Structure | Melting Point (°C) | Boiling Point (°C) | | :-------------- | :------------------------------------------------ | :----------------: | :----------------: | | Sodium Chloride | Ionic (strong electrostatic forces) | 801 | 1465 | | Diamond | Giant Covalent (strong covalent bonds) | 4000 | 4830 | | Iron | Metallic (strong electrostatic forces) | 1538 | 2862 | | Methane | Simple Covalent (weak intermolecular forces) | -182 | -161.6 | *Table 3.7: Comparison of melting and boiling points.* The extremely low melting and boiling points of methane ($CH_4$) highlight that only a small amount of energy is needed to overcome its weak intermolecular forces, whereas ionic, giant covalent, and metallic substances require immense energy to break their strong bonds. ### Possible Questions/Answers * **Q:** The boiling points of halogens are: Fluorine (-188°C), Chlorine (-35°C), Bromine (+59°C), and Iodine (+184°C). Explain this trend. **A:** All halogens are non-polar molecules ($F_2, Cl_2, Br_2, I_2$). The only intermolecular forces present are instantaneous dipole-induced dipole forces. As you go down the group, the number of electrons and the atomic size increase. This makes the electron cloud more polarizable, leading to stronger van der Waals forces and therefore higher boiling points. * **Q:** Bromine ($Br_2$) and iodine monochloride ($ICl$) have the same number of electrons, but the boiling point of $ICl$ is nearly 40°C higher than that of $Br_2$. Explain this difference. **A:** Although both have the same number of electrons (and thus similar strength of van der Waals forces), $Br_2$ is a non-polar molecule. In contrast, $ICl$ is a polar molecule due to the electronegativity difference between iodine and chlorine. Therefore, $ICl$ experiences both instantaneous dipole forces and stronger permanent dipole-dipole forces, leading to a higher boiling point. * **Q:** Describe the hydrogen bonding between two HF molecules. **A:** In an HF molecule, the fluorine atom is highly electronegative, creating a strong partial negative charge ($\delta^-$) on itself and a strong partial positive charge ($\delta^+$) on the hydrogen atom. The partially positive hydrogen of one HF molecule is strongly attracted to a lone pair of electrons on the partially negative fluorine atom of a neighboring molecule. $$ H^{\delta+}-F^{\delta-} \cdot\cdot\cdot H^{\delta+}-F^{\delta-} $$