Chemistry: The Central Science - Chapter 11: Liquids and Intermolecular Forces

Chemistry: The Central Science - Chapter 11: Liquids and Intermolecular Forces

Intermolecular Forces

  • Intermolecular forces are attractions between molecules that are significantly weaker than intramolecular attractions (bonds).

  • Physical properties of substances, such as boiling points, melting points, viscosity, surface tension, and capillary action, reflect the strength of intermolecular forces.

States of Matter

  • The primary distinction between various states of matter is attributed to the strength of intermolecular forces.

  • Stronger intermolecular forces pull molecules closer together, while kinetic energy tends to keep them apart and in motion.

  • Note: Average kinetic energy is directly related to temperature.

Properties and States of Matter

  • Table 11.1 Characteristic Properties of the States of Matter:

    • Gas:

      • Assumes both volume and shape of its container.

      • Expands to fill its container.

      • Is compressible.

      • Flows readily.

      • Rapid diffusion occurs.

    • Liquid:

      • Assumes shape of the portion of the container it occupies.

      • Does not expand to fill its container.

      • Is virtually incompressible.

      • Flows readily.

      • Slow diffusion occurs.

    • Solid:

      • Retains its own shape and volume.

      • Does not expand to fill its container.

      • Is virtually incompressible.

      • Does not flow.

      • Extremely slow diffusion occurs.

    • The atoms within a solid can vibrate in place, with increased vibrational motion at higher temperatures.

Relative Strength of Attractions

  • Table 11.2 Melting and Boiling Points of Representative Substances:

    • Chemical Bonds:

      • Ionic Bonds:

        • Lithium fluoride (LiF): Melting Point = 1118 K, Boiling Point = 1949 K.

      • Metallic Bonds:

        • Beryllium (Be): Melting Point = 1560 K, Boiling Point = 2742 K.

      • Covalent Bonds:

        • Diamond (C): Melting Point = 3800 K, Boiling Point = 4300 K.

    • Intermolecular Forces:

      • Dispersion Forces:

        • Nitrogen (N2): Melting Point = 63 K, Boiling Point = 77 K.

      • Dipole-Dipole Interactions:

        • Hydrogen Chloride (HCl): Melting Point = 158 K, Boiling Point = 188 K.

      • Hydrogen Bonding:

        • Hydrogen Fluoride (HF): Melting Point = 190 K, Boiling Point = 293 K.

  • Intermolecular attractions are generally weaker than chemical bonds.

  • Note: Hydrogen bonds are not considered chemical bonds.

Types of Intermolecular Forces between Neutral Molecules

  • Forces ranked from weakest to strongest:

    • Dispersion Forces (also known as London dispersion forces or induced dipole-induced dipole interactions).

    • Dipole-Dipole Forces.

    • Hydrogen Bonding (considered a special form of dipole-dipole interactions).

  • The first two types of forces are collectively referred to as van der Waals forces.

Dispersion Forces

  • Nonpolar particles can become temporarily polarized to create dispersion forces.

  • The ease of distortion in an electron cloud is referred to as polarizability.

Factors Affecting Dispersion Forces

  • Number of electrons in an atom: More electrons result in stronger dispersion forces.

  • Size of the atom or molecule/molecular weight: Larger size increases dispersion forces.

  • Molecular shape: Compacted shapes exhibit less dispersion force than more extended shapes of similar mass.

Polarizability and Boiling Point

  • Greater polarizability correlates with a lower boiling point, indicating weaker intermolecular forces; smaller or lighter molecules generally show this pattern due to fewer electrons and lower molecular weight.

Dipole-Dipole Interactions

  • Polar molecules possess both positive and negative poles, resulting in a dipole structure.

  • Oppositely charged ends attract, enhancing intermolecular forces.

  • For molecules of comparable mass and size, a more polar molecule will exhibit a higher boiling point due to stronger dipole-dipole interactions.

Comparative Effects of Dipole-Dipole and Dispersion Forces

  • In similarly sized and shaped molecules, dipole-dipole interactions dominate.

  • For larger molecules, dispersion forces often dictate physical properties.

Hydrogen Bonding

  • Occurs when hydrogen is bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine. The resultant attraction forms a hydrogen bond.

    • A hydrogen bond is defined as the interaction between a hydrogen atom bonded to a highly electronegative atom and another electronegative atom in a nearby molecule or chemical group.

  • The strength of hydrogen bonding is due to the high electronegativity of N, O, and F, enabling significant interaction with the hydrogen nucleus that has minimal electron shielding.

  • Comparison of Ice and Liquid Water:

    • Hydrogen bonding leads to increased space between molecules in ice than in liquid water, making ice less dense.

Ion-Dipole Interactions

  • These interactions are present when ions are mixed in polar solvents; the strength of these interactions facilitates the solubility of ionic substances in polar solvents like water.

Determining Intermolecular Forces

  • Type of Intermolecular Interaction:

    • Nonpolar molecules (e.g., Ne, Ar): Dispersion forces.

    • Polar molecules without OH, NH, or HF groups (e.g., BF3, CH4, HCl, CH3CN): Dipole-dipole interactions.

    • Polar molecules with OH, NH or HF groups (e.g., H2O, NH3): Hydrogen bonding.

    • Ionic solids dissolved in polar liquids (e.g., NaCl in H2O): Ion-dipole interactions.

  • General Energy Strengths:

    • Dispersion forces: 0.1–30 kJ/mol.

    • Dipole-dipole interactions: 2–15 kJ/mol.

    • Hydrogen bonding: 10–40 kJ/mol.

    • Ion-dipole interactions: >50 kJ/mol.

  • Key Point: All chemicals exhibit dispersion forces; the strongest force dictates the strength of molecular attractions.

Generalizations about Relative Strengths of Intermolecular Forces

  1. When comparing two molecules with similar molar masses and shapes, dispersion forces are approximately equal.

  2. For two molecules with significantly different molar masses and absent hydrogen bonding, dispersion force determines the stronger attractions.

Liquid Properties Affected by Intermolecular Forces

  • Key properties influenced by intermolecular forces include boiling point, melting point, viscosity, surface tension, and capillary action.

Viscosity

  • Viscosity refers to the resistance of a liquid to flow, linked to how easily molecules can move past one another.

  • Viscosity increases with stronger intermolecular forces and decreases with heightened temperature.

Surface Tension

  • Water appears to have a "skin" due to additional inward forces at its surface, leading to the phenomenon where water beads upon contact with nonpolar surfaces.

Cohesion and Adhesion

  • Cohesive Forces: Intermolecular forces that bond similar molecules.

  • Adhesive Forces: Intermolecular forces that draw a substance to a surface.

  • These forces are critical in capillary action.

Capillary Action

  • Defined as the rise of liquids in narrow tubes, driven by adhesive forces that attract liquid to tube walls and cohesive forces that allow the liquid to stick to itself.

  • Water demonstrates stronger adhesive forces with glass, while mercury exhibits stronger cohesive forces internally.

Phase Changes

  • A phase change is the conversion of one state of matter to another, involving energy gain or release. Common phase changes include melting/freezing, vaporizing/condensing, and subliming/depositing.

Energy Changes and Change of State

  • Heat of Fusion: The energy needed to convert a solid to a liquid at its melting point.

  • Heat of Vaporization: The energy needed for a liquid to become gas at its boiling point.

  • Heat of Sublimation: The energy required to change a solid directly to a gas.

Heating Curves

  • A heating curve is a graphical representation of temperature changes as heat is added. Within a single phase, heat relates to specific heat, mass, and temperature change. During phase changes, temperature remains constant.

Supercritical Fluids

  • Gases liquefy under pressure; the temperature at which a gas cannot be compressed is the critical temperature, with critical pressure defined as the pressure necessary for liquefaction at this temperature.

  • Above this temperature and pressure state, the substance is termed a supercritical fluid.

Vapor Pressure

  • At any given temperature, certain liquid molecules possess enough energy to transition into gas, which increases with a rise in temperature. The pressure created by these escaped molecules is the vapor pressure. Liquid and vapor achieve dynamic equilibrium when their evaporation and condensation rates equalize.

  • The boiling point of a liquid is defined as the temperature at which its vapor pressure equals atmospheric pressure, with the normal boiling point occurring at a vapor pressure of 760 torr.

Phase Diagram

  • A phase diagram visually represents states of matter across varying temperature and pressure. It details changes of state alongside Triples Points and Critical Points.

Phase Diagrams Summary

  • Phase diagrams delineate the three states of a substance, their boundary curves indicating phase change, with key points marking the conditions for phase coexistence:

    • Triple Point: The condition where solid, liquid, and gas phases coexist in equilibrium.

    • Critical Point: The temperature above which the liquid phase ceases to exist regardless of pressure.

Special Phase Features of Water and Carbon Dioxide

  • Water showcases a high critical temperature and pressure due to intense van der Waals forces, exhibiting an atypical negative slope of the melting curve.

  • Carbon dioxide presents features such as the inability to exist in liquid form at pressures below 5.11 atm and sublimation at standard pressures.