Comprehensive Study Notes: Bonds and Intermolecular Attractive Forces in Solids, Liquids, and Solutions

Comparison of Bonding Forces (Intramolecular)

• Intramolecular Attractive Forces (Bonds): These are the forces found within a molecule (strong bonds). They determine the chemical behavior of a substance. The chemical behavior of each phase of matter (solid, liquid, gas) is essentially consistent because the constituent particles (the molecules) remain the same.

• Intermolecular Attractive Forces (IMAFs): These are the nonbonding forces found between molecules (weak attractions). They determine the physical behavior of a substance, such as melting and boiling points, because the strength of these forces differs from state to state. Intermolecular forces are responsible for phase changes.

• Thermodynamics of Bonding:

  • Breaking Bonds/Attractions: Requires energy ($+E$), making it an endothermic process.

  • Forming Bonds/Attractions: Releases energy ($-E$), making it an exothermic process.

• Types of Bonding Forces:

  • Ionic:

    • Model: Cation-anion attraction ($M+NM$).

    • Energy Range: $400\text{--}4000 \text{ kJ mol}^{-1}$.

    • Basis of Attraction: Electrostatic attraction between positive and negative ions.

    • Example: $NaCl$.

  • Covalent:

    • Model: Nuclei shared electron pair ($NM+NM$).

    • Energy Range: $150\text{--}1100 \text{ kJ mol}^{-1}$.

    • Basis of Attraction: Mutual attraction of nuclei for shared electrons.

    • Example: $H-H$.

  • Metallic:

    • Model: Cations delocalized electrons.

    • Energy Range: $75\text{--}1000 \text{ kJ mol}^{-1}$.

    • Basis of Attraction: Attraction between metal cations and the "sea" of delocalized valence electrons.

    • Example: $Fe$.

• Note on Energy Units: $1 \text{ cal} = 4.18 \text{ J}$ .

Lewis Structures and Molecular Geometry (AXE Notation)

• Lewis Dot Structures: Used to visualize valence electrons and bonding. For example, water ($H_2O$) has $2(1) + 6 = 8 \text{ valence electrons}$.

• AXE Notation Guide:

  • A: Central atom.

  • X: Surrounding atoms.

  • E: Lone pairs on the central atom.

• Geometry Examples:

  • $AX_{2}$: Linear. Nonpolar (symmetric). Example: $BeCl_{2}$. ($2 + 2(7) = 16 \text{ e}^{-}$).

  • $AX_{2}E$: Bent. Polar (asymmetric). Example: $SO_{2}$.

  • $AX_{3}$: Trigonal planar. Nonpolar (symmetric). Example: $BF_{3}$. ($3 + 3(7) = 24 \text{ e}^{-}$).

  • $AX_{3}E$: Trigonal pyramidal. Polar (asymmetric). Example: $NH_{3}$. ($5 + 3(1) = 8 \text{ e}^{-}$).

  • $AX_{4}$: Tetrahedral. Nonpolar (symmetric). Example: $CH_{4}$. ($4 + 4(1) = 8 \text{ e}^{-}$).

  • $AX_{2}E_{2}$: Bent. Polar (asymmetric). Example: $H_{2}O$. ($2 + 6 = 8 \text{ e}^{-}$).

• Polarity Principles: A molecule is generally polar if it has an asymmetrical shape or if the central atom has one or more lone pairs (unsymmetrical electron distribution). Nonpolar molecules are generally symmetrical (e.g., $AX_{2}$, $AX_{3}$, $AX_{4}$).

Intermolecular Attractive Forces (IMAFs)

• London Dispersion Forces (LDFs):

  • Present in all molecules (both polar and nonpolar).

  • Mechanism: Electrons in orbitals are in constant motion. Occasionally, they wind up on the same side of an atom, creating an instantaneous dipole (momentary shortage or excess of electrons). This instantaneous dipole induces a dipole in a neighboring atom, creating an electrostatic attraction.

  • Polarizability: The tendency of an electron cloud to distort and become temporarily polar. Larger electron clouds are more polarizable.

  • Factors Affecting LDF Strength:

    1. Molecular Weight (MW): As $MW$ increases, the number of electrons increases, creating larger electron clouds that are more polarizable. This increases IMAF strength and boiling point ($BP$).

    2. Molecular Shape: Long, skinny (linear) molecules have stronger IMAFs than spherical isomers because they have more surface area to form attractions.

       • Example: $n\text{-Pentane}$ ($BP = 309.4 \text{ K}$) vs. $Neopentane$ ($BP = 282.7 \text{ K}$).

  • LDF Trends in Elements:

    • Halogens ($X_{2}$):

      • $F_{2}$ ($MW = 38.0 \text{ amu}$, $BP = 85.1 \text{ K}$).

      • $Cl_{2}$ ($MW = 71.0 \text{ amu}$, $BP = 238.6 \text{ K}$).

      • $Br_{2}$ ($MW = 159.8 \text{ amu}$, $BP = 332.0 \text{ K}$).

      • $I_{2}$ ($MW = 253.8 \text{ amu}$, $BP = 457.6 \text{ K}$).

    • Noble Gases:

      • $He$ ($MW = 4.0 \text{ amu}$, $BP = 4.6 \text{ K}$).

      • $Ne$ ($MW = 20.2 \text{ amu}$, $BP = 27.3 \text{ K}$).

      • $Ar$ ($MW = 39.9 \text{ amu}$, $BP = 87.5 \text{ K}$).

      • $Kr$ ($MW = 83.8 \text{ amu}$, $BP = 120.9 \text{ K}$).

      • $Xe$ ($MW = 131.3 \text{ amu}$, $BP = 166.1 \text{ K}$).

• Dipole-Dipole Forces:

  • Found only in polar molecules with permanent dipoles ($X\text{--}Y$).

  • The more polar the molecule (higher Dipole Moment), the higher the boiling point due to stronger IMAFs.

  • Data Comparison of Molecules with Similar Weights:

    • Propane ($CH_{3}CH_{2}CH_{3}$, $MW=44$, Dipole $= 0.1$, $BP = 231 \text{ K}$).

    • Dimethyl ether ($CH_{3}OCH_{3}$, $MW=46$, Dipole $= 1.3$, $BP = 248 \text{ K}$).

    • Methyl chloride ($CH_{3}Cl$, $MW=50$, Dipole $= 1.9$, $BP = 249 \text{ K}$).

    • Acetaldehyde ($CH_{3}CHO$, $MW=44$, Dipole $= 2.7$, $BP = 294 \text{ K}$).

    • Acetonitrile ($CH_{3}CN$, $MW=41$, Dipole $= 3.9$, $BP = 355 \text{ K}$).

• Hydrogen Bonding:

  • The strongest type of Dipole-Dipole interaction.

  • Occurs when an electron-deficient hydrogen atom (bonded to $N$, $O$, or $F$) is attracted to a small, very electronegative $N$, $O$, or $F$ atom on a nearby molecule.

  • Explains unusually high boiling points for group hydrides like $H_{2}O$, $NH_{3}$, and $HF$.

• Ion-Dipole Forces:

  • Observed when ionic solutes dissolve in polar solvents (e.g., $NaCl(aq)$).

  • Involves cation-dipole and anion-dipole attractions (solvation).

States of Matter and Phase Changes

• Physical States: Determined by the competition between Kinetic Energy ($KE$) and IMAFs.

  • Gas: High $KE$, low IMAF effect; particles are far apart. Increasing pressure or cooling promotes the liquid state.

  • Liquid: Particles are closer; IMAFs hold them together but allow Flow.

  • Solid: Low $KE$, high IMAF effect; particles are in fixed positions.

• Energy and Phase Changes:

  • Heat of Fusion ($H_{fus}$): Energy required to change a solid to a liquid at its melting point.

  • Heat of Vaporization ($H_{vap}$): Energy required to change a liquid to a gas at its boiling point. For most substances, H_{vap} > H_{fus}.

  • Heating Curve: During a phase change, the temperature does not change. Added kinetic energy is used solely to overcome IMAFs (increasing potential energy and separating particles).

• Types of Phase Transitions:

  • Melting/Sublimation/Vaporization (Endothermic, +E).

  • Freezing/Deposition/Condensation (Exothermic, -E).

Physical Properties Affected by IMAFs

• Boiling Point ($BP$) and Melting Point ($MP$): Increase as IMAF strength increases.

• Surface Tension: Results from the net inward force experienced by surface molecules of a liquid. Increases with IMAF strength.

• Viscosity: Resistance of a liquid to flow.

  • Increases with stronger IMAFs.

  • Decreases at higher temperatures.

  • Example: Viscosity of alkanes increases with chain length ($Hexane \rightarrow Decane$).

• Capillary Action: Result of cohesive forces (attraction to same molecules) and adhesive forces (attraction to different molecules like glass).

  • Water: Forms a concave meniscus (adhesion > cohesion).

  • Mercury: Forms a convex meniscus (cohesion > adhesion).

• Vapor Pressure ($VP$): The pressure exerted by a vapor in dynamic equilibrium with its liquid phase.

  • Higher IMAFs result in lower vapor pressure (less volatile).

  • As temperature increases, vapor pressure increases as more molecules have enough $KE$ to escape into the gas phase.

  • Normal Boiling Point: The temperature at which vapor pressure equals $1 \text{ atm}$ ($760 \text{ torr}$).

Bonding in Solids

• Molecular Solids: Held by weak IMAFs (LDFs, dipole-dipole, H-bonds). Soft, low $MP$. Examples: $I_{2}$, $H_{2}O$, $Ar$.

• Covalent-Network Solids: Hard solids where atoms are shared throughout via strong covalent bonds. Very high $MP$. Examples: Diamond ($C$), Quartz ($SiO_{2}$), Tungsten Carbide ($WC_{2}$).

• Ionic Solids: Crystal lattice of cations and anions held by Coulombic attraction ($E = k \frac{q_{1}q_{2}}{d}$). Hard, brittle, high $MP$, conduct only when molten or aqueous. Examples: $NaCl$, $CuSO_{4}$.

• Metallic Solids: Metal cations in a sea of delocalized valence electrons. Malleable, ductile, excellent conductors, variable hardness and $MP$. Includes pure metals ($Fe$, $Cu$) and alloys.

Alloys

• Definition: Homogeneous metallic mixtures (solutions) held by metallic bonding.

• Substitutional Alloys: Formed when atoms have similar atomic radii; density is usually between the components. Example: Brass ($Cu$, $Zn$).

• Interstitial Alloys: Formed when smaller atoms fit in the "holes" between larger atoms. These are denser and less malleable (harder) than pure components. Example: Steel ($Fe$, $C$).

Organic Functional Groups and IMAFs

• Hydrocarbons (Alkanes, Alkenes, Alkynes): Nonpolar. Only LDFs. $BP$ increases with chain length.

• Alcohols ($R\text{--}OH$): Polar and exhibit Hydrogen Bonding. Higher $BP$ than corresponding alkanes.

• Ethers ($R\text{--}O\text{--}R$): Polar (Dipole-Dipole).

• Amines ($R\text{--}NH_{2}$): Polar and exhibit Hydrogen Bonding.

• Aldehydes and Ketones ($C=O$ group): Polar (Dipole-Dipole).

• Carboxylic Acids ($COOH$ group): Polar and exhibit Hydrogen Bonding.

Biological Applications

• Protein Structure: Covalent peptide bonds form the backbone, while IMAFs (H-bonds, LDFs) determine the specific 3D folding.

• DNA Structure: Hydrogen bonds hold the double helix together ($T=A$, $C=G$) but are weak enough to allow the strands to separate for replication.