CH 10: Liquids and Solids

Liquids and Solids Overview

  • Course: CHEM 1500

  • Focus: Polar Molecules, Intermolecular Forces, Properties of Liquids and Solids, Phase Changes, and Heating Curves

Polar Molecules

Definitions

  • Polar Molecules: Molecules with an uneven distribution of charge due to polar covalent bonds, leading to molecular dipoles.

  • Polar Bonds: A bond between two atoms with different electronegativities, resulting in a bond dipole.

Electronegativity Differences

  • Polar Covalent Bond: Difference in electronegativity (∆EN) greater than 0.5 but less than 2.0.

  • Ionic Compound: Difference in electronegativity greater than 2.0.

  • Labeling Polarity: Use of symbols δ+ (partial positive) and δ- (partial negative) to indicate charge distribution.

Dipole Moments

  • Definition: A measurement of polarity within a molecule.

    • Formula: extDipoleMoment(µ)=Qimesrext{Dipole Moment (µ)} = Q imes r

    • Where Q = charge (Coulombs) and r = distance (meters).

    • Unit: Debyes (D)

    • Example: HCl has a dipole moment of 1.11 D.

Determining Molecular Polarity

  1. Draw Complete Lewis Dot Structure.

  2. Use VSEPR to Determine Geometry.

  3. Assess Bond Polarity.

  4. Determine Molecular Polarity Based on Bond Dipoles and Geometry.

  • Example molecules:

    • Ammonia (NH3) : Polar

    • Carbon Dioxide (CO2) : Nonpolar

Intermolecular Forces

Definition

  • Interaction forces that exist between molecules, responsible for determining physical properties like boiling/melting points, vapor pressure, and viscosity.

Types of Intermolecular Forces

  1. London Dispersion Forces:

    • Also known as dispersion forces; arise from instantaneous dipoles that induce dipoles in neighboring molecules.

    • Present in all molecules, regardless of polarity; strength depends on polarizability.

    • Effect of Shape: Long, skinny molecules (n-pentane) have stronger forces than short, fat molecules (neopentane).

      • Boiling Points Example:

      • n-Pentane: 309.4 K

      • Neopentane: 292.7 K

  2. Dipole-Dipole Interactions:

    • Occur between molecules with permanent dipoles; the positive end of one dipole is attracted to the negative end of another.

    • Importance: Significant only when molecules are in close proximity.

    • Example Data: (Molecular Weight in amu, Dipole Moment in D, Boiling Point in K)

      • Propane (C3H8): 44, 0.1, 231

      • Dimethyl Ether: 46, 1.3, 248

      • Acetaldehyde: 44, 2.7, 294

  3. Hydrogen Bonding:

    • A strong type of dipole-dipole interaction when H is bonded to highly electronegative elements (N, O, F).

  4. Ion-Dipole Interactions:

    • Interactions between an ion and a polar molecule; crucial in ionic substances dissolving in polar solvents.

Properties of Liquids

Viscosity

  • Resistance of a liquid to flow; affected by intermolecular forces and temperature.

  • Viscosity Data (Substance, Formula, Viscosity in kg m$^{-1}$ s$^{-1}$):

    • Hexane (C5H12): 3.26imes10−43.26 imes 10^{-4}

    • Heptane (C7H16): 4.09imes10−44.09 imes 10^{-4}

    • Decane (C10H22): 1.42imes10−31.42 imes 10^{-3}

Surface Tension

  • Results from the net inward force experienced by surface molecules in a liquid.

Phase Changes

Types of Phase Changes

  1. Fusion/Freezing: Solid to liquid / Liquid to solid

  2. Vaporization/Condensation: Liquid to gas / Gas to liquid

  3. Sublimation/Deposition: Solid to gas / Gas to solid

Energetics of Phase Changes

  • The sign of ΔH (enthalpy) is positive for vaporization and negative for condensation.

  • The sign of ΔS (entropy) is positive for vaporization and negative for condensation.

Heating Curves

  • Illustrate temperature changes and phase transitions as heat is added to or removed from a substance.

    • Example of propane in cooling systems, showcasing the endothermic nature of phase transitions.

    • Calculate energy needed to convert liquid propane at specific temperatures.

Vapor Pressure

Definition

  • The pressure exerted by vapor in equilibrium with its liquid/solid phase.

Vapor Pressure Curves

  • Plots representing the relationship between vapor pressure and temperature.

  • Understanding the Clausius-Clapeyron equation:

    • Form 1: extlnP<em>vap=−racriangleH</em>vapRrac1T+Cext{ln } P<em>{vap} = - rac{ riangle H</em>{vap}}{R} rac{1}{T} + C

    • Linear relationship with pressure.

    • Form 2: ext{ln } P2 = ext{ln } P1 + rac{ riangle H{vap}}{R} igg( rac{1}{T1} - rac{1}{T_2} igg)

Boiling Point

  • The temperature at which the vapor pressure equals atmospheric pressure, vital for understanding phase transitions.

Types of Solid Materials

Classification

  1. Molecular Solids: Discrete molecules (e.g., I2).

  2. Metallic Solids: Atoms held by metallic bonds (e.g., Fe).

  3. Network Covalent Solids: Atoms connected by covalent bonds (e.g., diamond).

  4. Ionic Solids: Composed of positive and negative ions (e.g., NaCl).

Bonding Properties in Crystalline Solids

  • Variations in bonding governed by types of forces; examples listed showing melting points, hardness, and conductivity properties.

  • Unit Cell Definition: The smallest repeating unit in a crystal lattice that can depict the entire crystal structure.

Phase Diagrams

  • Graphical representation of the physical states of a substance under varying temperature and pressure conditions.

    • Examples: Phase diagrams for water, CO2, and nitrogen illustrating phase changes under different environmental conditions.

    • Critical Temperature and Pressure definitions provided: critical temperature is the highest temperature at which a liquid phase can exist, while critical pressure is the minimum pressure needed to maintain this phase at the critical temperature.