General Chemistry 004 Summer 2025 Lecture Notes

Chapter 12 Learning Outcomes

• Understand and describe intermolecular attractive interactions between molecules and ions.

• Compare relative strengths of intermolecular attractions based on molecular structure or physical properties.

• Understand concepts of polarizability, viscosity, and surface tension in liquids.

• Know phase changes for a pure substance.

• Interpret heating curves and calculate quantities related to temperature and enthalpies of phase changes.

• Define critical pressure, critical temperature, vapor pressure, normal boiling point, normal melting point, critical point, triple point.

• Classify solids based on bonding/intermolecular forces; understand how differences in bonding relate to physical properties.

• Recognize some properties and characteristics of metals.

• Explain the electron-sea model of metallic bonding.

• Interpret melting and boiling point trends for metals.

States of Matter

• Gases:

  • Compressible fluids, flow readily, and show high diffusion

  • Assume volume and shape of the container

• Liquids:

  • Nearly incompressible, flow readily, diffuse slowly

  • Assume shape, but not volume of the container

• Solids:

  • Virtually incompressible, does not flow, nearly no diffusion

  • Do not assume shape or volume of the container

• Plasma:

  • Highly charged gases with extremely high kinetic energy

States of Matter: Interparticle Distance

• The fundamental difference between states of matter is the distance between atoms or particles.

• Condensed phase.

• Temperature.

Real Gases

• Real gases do not behave like ideal gases at high pressure or low temperature.

• Ideal gas law assumes no intermolecular attraction and gases do not take up space.

Intramolecular Bonding

• All substances containing multiple atoms are “bonded.”

• Electrostatic interactions:

  • Attractive and repulsive forces

• Ionic bonds:

  • Positive and negative electrostatic interactions

• Covalent bonds:

  • When atoms share a pair of electrons

  • Positive charge of nuclei interacting with the negative electrons

Covalent Bonds

• Non-polar covalent bond: when two atoms share electrons equally.

• Polar covalent bond: when two atoms share electrons unequally.

• Electrons have a higher probability of being found near the more electronegative atom.

• Not a complete transfer of electrons (ionic bond) and this results in partial negative and positive charges.

Unequal Covalent Bonds

• Polar covalent bonds:

  • Electrons in a covalent bond are not always shared equally

  • Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does

  • Therefore, the fluorine end of the molecule has more electron density than the hydrogen end

Unequal Covalent Bonds: Electronegativity

• Electronegativity is the ability of an atom in a molecule to attract electrons to itself.

• On the periodic table, electronegativity generally:

  • Increases from left to right across a period

  • Increases from the bottom to the top of a group

Intermolecular Bonding

• Intermolecular – forces between two molecules

• Dictate physical properties including:

  • Boiling point

  • Melting Point

  • Vapor pressure

  • Viscosity

• Intermolecular Bonding

Types of Bonding

Dipole-Dipole

• Molecules with permanent dipoles are attracted to each other.

• Molecules need to be close.

• H2O, HF, NH3

Ion-Dipole

• Solvation

• Polar liquids dissolve salts

• Full charge of ions attracted to opposite partial charge of solvent

London Dispersion Forces

• All substances display dispersion forces

• Described by Fritz London for non-polar molecules being liquids

• Attractive forces between an instantaneous and induced dipole

• Present in all atoms and molecules

• Tendency of an electron cloud to distort is called polarizability

Influences on London Dispersion Forces

• Shape of the molecule

  • Cylindrical vs spherical

• Size of the electron cloud

  • Correlates with mass

Dipole vs Dispersion Forces

• What has a greater effect

  • If two molecules are similar in size, dipole forces will be dominant influence for differences in physical properties

  • If two molecules are dissimilar in size, dispersion forces likely will be dominant in affecting physical properties

  • But…what about water?

Hydrogen Bonding

• High electronegativity of N, O, F exposes the nucleus of bound hydrogen

Viscosity

• Viscosity: Resistance to flow

• How readily molecules move past each other

• Increases with stronger intermolecular forces

• Decreases with temperature

• Dependent on shape and size

Surface Tension

• Surface Tension: inward force experienced by molecules on the surface

• Correlates to strength of intermolecular forces

• Cohesive vs adhesive forces

Surface Tension: Description

Surface Tension: inward force experienced by molecules on the surface.
Influence of gravity

Capillary Action

• Capillary action: intermolecular forces between a substance and surface

• Adhesive forces and cohesive forces and the meniscus

Phase Changes

• Solid → Liquid = Melting *endothermic

• Liquid → Solid = Freezing *exothermic

• Liquid → Gas = Vaporization *endothermic

• Gas → Liquid = Condensation *exothermic

• Solid → Gas = Sublimination *endothermic

• Gas → Solid = Deposition *exothermic

Energy and Phase Change

• Heat of fusion: energy required to change a solid to a liquid

  • \Delta H_{fus} = 6.01 \frac{kJ}{mol}

• Heat of vaporization: energy required to change a liquid to a gas

  • \Delta H_{vap} = 40.7 \frac{kJ}{mol}

Evaporation and Vaporization

• Some molecules have enough kinetic energy to escape

• Everything has a vapor pressure

• As temperature increases, kinetic energy and partial pressure increases

• Dynamic equilibrium is when condensation and vaporization are equal

Energy and Phase Change: Clausius-Clapeyron Equation

• Boiling point of a liquid occurs when the vapor pressure equals atmospheric pressure

• Standard pressure is 1 atm or 760 torr

• Clausius-Clapeyron Equation

• ln\frac{P1}{P2} = \frac{\Delta H{vap}}{R} \left( \frac{1}{T2} - \frac{1}{T_1} \right)

Phase Diagrams

• Substances exist as a solid, liquid or gas depending on external temperature and pressure

• Lines of equilibrium

• Critical point

• Triple point

Structures of Solids

• Crystalline – solids with a regular repeating pattern

• Amorphous – solids with a distinct lack of order

Liquid Crystals

• Friedrich Reinitzer and cholesteryl benzoate

• Intermediate waxy phase between solid and liquid

• Alkyl chain, polar group, aromatic group

• Ordering affected by temperature, pressure and electric fields

Liquid Crystals Applications

• Nematic liquid crystals can become translucent when subjected to an electric field

• Liquid crystals can be influenced by temperature and photons

Applications of Solids

• Integrated circuits

• Purification and remediation

• Medical devices

• Nanoparticles

• Catalysts

• Insulators

• Sensors

• Fabrication and consumer products

• Jewelry

Bonding in Solids

• Metallic – sea of electrons

  • Conducts electricity; strong; malleable

• Ionic – electrostatic interactions

  • Poor conductors; brittle

• Covalent-network – covalent bonds

  • Extremely hard, often semi-conductors

• Molecular – intermolecular bonds

  • Soft; low melting points

• Nanomaterials

Structures of Solids (Revisited)

• Crystalline – solids with a regular repeating pattern

• Amorphous – solids with a distinct lack of order

Ionic Solids

• Electrostatic interactions (Coulombic forces)

• Charge and size dictate lattice energy

• Insulators

• Brittle

• High melting and boiling points

• Close-packed motifs

  • Maximize cation-anion interaction

  • Alternating charged ions

Ionic Solids Influenced by Size

• Comparison of CsCl, NaCl and ZnS crystal lattices.

• Each with different Cation radius, Anion radius, Crystal structure, r+/r_ ratio, Cation coordination number, and Anion coordination number.

Ionic Solids Influenced by Charge

• Comparison of NaF, MgF2, and ScF3.

• Each with different Cation coordination number, Cation coordination geometry, Anion coordination number, and Anion coordination geometry.

Simplified Metallic Bonding

• Covalent vs metallic bonding diagram, contrasting bonding in Group 3A-7A elements.

Orbital Review

• MO Diagram and Band Molecular Orbitals.

Band Structure: a MO Approach

• Energy plot showing various atomic orbitals of a single nickel atom vs the band structure of a nickel crystal

Types of Conductors

• Table of various materials with their band gap energy values and conductivities

Transition Metal Trends

• Melting point and boiling point trends for transition metals displayed on periodic tables

Alloys

• Contains two or more elements

• Similar properties to elemental metals

• TABLE 12.2⚫Some Common Alloys

Alloys - Types

• Substitutional – substitute atoms of similar size

• Interstitial – fill void spaces with smaller atoms

• Heterogeneous alloys – creates grain boundaries

• Intermetallic compounds – discrete compounds through bonding

• Eutectic – homogeneous mixture with low melting point (InGa)

Superconductors

• At very low temperatures, a substance that has no resistance to the flow of electricity and creates its own magnetic field

Covalent Networks

• Vast and continuous bonding of atoms

• Silicon, silica, carbon (graphite and diamond)

• High melting points and hardness

Molecular Solids

• Intramolecular forces holding atoms or neutral molecules together

• Low melting points, soft solids

• Individual units are discrete molecules

• Mass, functional groups and symmetry influence physical properties

Polymers

• Important for fabrication, plastics, clothing

• Only carbon and boron are capable of polymeric systems

• Thermoplastics – can be remolded with heat

• Thermosetting – cannot be reshaped

• Addition – catalyzation of ethylene

• Condensation – copolymerization of NH2 with COOH

Polymerization

• Addition

  • Converting electrons within monomeric π-bonds to σ-bonds

• Condensation

  • Reaction of two subunits (co-polymers) to a polymer with a byproduct

Properties of Polymers

• Macromolecules

• Chain length – longer chains have higher molecular weight

• Branching inhibits crystallinity

• HDPE chains are longer with less branching leading to higher crystallinity compared to LDPE

Crosslinking Polymers

• Chemical bonding of polymeric chains stiffen and strengthen the overall polymer

• Vulcanization

• Charles Goodyear

Interdisciplinary Solids

• Polymers – often molecular solids connected by long chains

  • Stronger and higher melting points than molecular solids

  • More flexible than other solids

• Nanomaterials – generally 1-100 nm individual systems

  • Can display unique quantum or physical properties not observed in bulk

Scaling Nanoparticles

• We have two nanoparticles, one measuring 10 nm and the other 25 nm

• At 10 nm:

  • There are approx. 32,000 atoms and 8,000 exist on the surface

• At 25 nm:

  • There are approx. 480,000 atoms, and 48,000 exist on the surface

Nanoparticles

• Particles between 1 and 1000 nm (debatable)

• Anything from proteins to crushed sand to quantum dots

• In chemistry, we’re interested in those that display unique quantum properties which are between 1 and 100 nm

Carbon Nanoparticles

• 1985 - Richard Smalley and the buckyball (buckminsterfullerene)