General Chemistry 2 - Intermolecular Forces and Phase Transitions

Kinetic Molecular Theory

• Explains how gases behave
• "Kinetic" comes from the Greek word kinein, meaning "to move"
• Gas comprises randomly moving, widely separated molecules
• Negligible attractive forces between ideal gas molecules due to wide separation
• Increased absolute temperature raises the average kinetic energy of gas molecules
• This increase leads to more collisions and energy transfer

Phases of Matter

• Kinetic Molecular Theory explains the existence of solid, liquid, and gas phases
• Explains physical characteristics and phase changes

Key Concepts Explained by Kinetic Molecular Theory

• Gas pressure
• Compressibility
• Diffusion
• Mixing
• Reaction rates
• Equilibrium

Particle Arrangement

• Solid: Closely packed together
• Liquid: Slightly far from one another
• Gas: Freely moving, very far from each other

Volume and Shape

• Solid: Definite volume and shape
• Liquid: Definite volume, assumes the shape of part of a container
• Gas: Indefinite volume and shape

Density

• Solid: High
• Liquid: High
• Gas: Low

Motion of molecules

• Solid: Vibration in place
• Liquid: Random
• Gas: Fast, random

Phase Transition

• Describes the temperature variation of melting, sublimation, evaporation, freezing, condensation or deposition

Intermolecular Forces

• Forces holding individual particles (atoms, molecules, ions)
• Strength depends on arrangement, proximity, and nature of interacting particles
• Influence resulting properties of solids, liquids, and gases
• Stronger forces lead to higher melting points and heats of fusion in solids

Intramolecular forces of attraction

• Responsible for interactions within a molecule, like covalent and ionic bonds
• Stronger than intermolecular attractions

London Dispersion Force

• Weakest intermolecular force
• Caused by polarization or distortion of electron cloud due to highly charged particle
• One atom's electron cloud attracted to another's positively charged nucleus
• Occurs when a molecule exhibits temporary polarity and weakly attracts another molecule

Strength dependence

• Strength depends on the number of electrons present, due to polarization

Dipole-Dipole Interaction

• Attraction between partially positive and partially negative ends
• Occurs in polar covalent molecules (e.g., amino acids) with shared electrons
• Effective over short distances only -Weak, about 1% of ionic bond strength
• Increased temperature diminishes strength

Ion-Dipole Interaction

• Interaction between an ion and a polar molecule
• Anion attracted to partially positive end
• Cation attracted to partially negative end
• Responsible for cation formation in solutions
• Elements from groups I and II easily form cations
• Elements lose electrons easily, which makes them cations

Hydrogen Bond

• Special dipole-dipole interaction
• Forms when hydrogen bonds with fluorine, oxygen, or nitrogen
• Weaker than ionic or covalent bonds
• The strongest intermolecular force when hydrogen bonds is between two atoms of two different molecule

Properties of Liquids due to Intermolecular Forces

• Viscosity
• Surface tension
• Vapor pressure
• Boiling point

Viscosity

• Liquid's resistance to flow
• Strongly dependent on intermolecular forces
• Highly viscous liquids exhibit strong intermolecular forces
• Example: glycerol has a viscosity of 1.7 pascal-seconds at 293K (20°C)

Surface Tension

• Resistance needed to increase liquid's surface area
• Surface free energy: work required to increase surface area by a unit area
• Waters surface tension is 72.8 mN/m, with a surface free energy of 72.8 mJ/m²
• Surface tension decreases with decreasing polarity and increasing temperature
• Important in aerosol preparation for drug delivery

Vapor Pressure

• Vaporization: transformation of liquid to gas
• Endothermic, requires energy to break intermolecular forces
• Evaporation: liquid escapes surface to gas above the liquid
• Condensation: gaseous molecules return to liquid phase
• The point at which these interactions occur equally is considered equilibrium
• Vapor pressure: pressure of vapor present at equilibrium
• High vapor pressure indicates weak intermolecular forces
• Volatility: ability to readily evaporate from an open vessel
• Stronger intermolecular forces lead to lower vapor pressure, require more energy leading to a higher boiling point

Distillation

• Method to separate mixture components or remove solvent impurities
• Performed by heating mixture at known temperature

Boiling Point

• Liquid boils when vapor pressure equals atmospheric pressure
• Normal boiling point is at 101 325 Pa
• Influenced by intermolecular forces
• Stronger forces causes the need of more energy, directly resulting in a higher boiling point

Importance of Water

• Participates in chemical reactions, especially aqueous processes
• Photosynthesis oxidizes water to yield oxygen
• Converts sun's energy into chemical energy
• Transports molecules, nutrients, and reaction products
• Influences biomolecule properties

Pure Water

• Odorless, tasteless, and colorless
• Molecule can form hydrogen bonds with many other molecules
• Polar due to lone pair of electrons in oxygen atom
• Considered the universal solvent
• Each molecule surrounded by four others through hydrogen bonding

Acid and Base in water based solutions

• Acid contains a hydrogen proton
• Base contains hydroxide (Arrhenius definition)
• Water acts as both acid and base, auto-protolyzing into hydronium and OH- ions

Buffer

• Ability to resist extreme pH changes
• Important environmental perspective
• Accidental spills of extreme pH substances won't severely alter natural bodies of water

Heat capacity

• High heat capacity prevents oceans or river systems from generating large temperature changes
• Solids have definite shape and volume

Crystalline Solids

• Highly ordered arrangement

Amorphous Solids

• Random and disordered arrangement
• Gradually soften when heated, melting at a wide range of temperatures

Examples of crystalline solids

• amethyst, fluorite, and pyrite

Examples of Amorphous solids

• glass, charcoal, and plastic containers

Space lattice

• Three-dimensional pattern formed representing particle locations
• Defines crystal structure

Unit Cell

• Smallest unit of the lattice
• Repeatedly stacked to resemble the whole

Coordination Number

• Number of particles present in a unit cell
• To quantify physical properties

Guidelines

• Particle inside belongs entirely, counted as one
• Particle on face is shared, counted as one-half
• Particle occupying edge, counted as one-fourth
• Particle at vertex, counted as one-eighth

X-ray Diffraction

• Solid's crystal structure determined via X-ray Diffraction
• Provides bond lengths and angles
• Pioneered by William Henry Bragg and son William Lawrence Bragg, who were awarded the 1915 Nobel Prize in Physics

Crystal Systems

• Cubic
• Tetragonal
• Orthorhombic
• Monoclinic
• Triclinic
• Hexagonal
• Rhombohedral

Isomorphous

• Different substances crystallizing in same lattice

Polymorphous

• Substances crystallizing in several arrangements

Polytypes

• Crystals same in two dimensions but different in the third
• Hexagonal or cubic close-packed

Crystalline Solid Types

• Classified by bonding nature or interactions among particles
• Ionic
• Molecular
• Network
• Metallic

Ionic Solids

• Hard, brittle, and poor conductors of heat and electricity
• Strong electrolytes when molten due to free mobile ions

Molecular solids

• Weak intermolecular forces, low melting points (1 to 673K)
• Soft and poor conductors of heat and electricity
• Examples include CH4, P4, O2, CO2, and fullerenes

Network Solids

• Large, giant molecules with covalently bonded atoms in a highly cross-linked, rigid network
• Polymeric materials, diamond, and quartz classified as network covalent solids
• Very hard with high melting points (1473 to 4273 K)
• Poor thermal and electrical conductors due to localized electrons in covalent bonds

Allotropes

• Allotropes are different forms of a substance that are localized
• Diamond and graphite forms of carbon

Metallic Solids

• Bounded by metallic bonding
• Intermolecular attraction between metal atom nucleus and negatively charged electrons
• Excellent thermal and electrical conductors
• Possess malleability, ductility, luster, and hardness
• Melting points range from 234 to 3673K
• Copper, Nickel, and Chromium are some examples of metals

Phase Changes involve transition to another phase

• Transformation from one phase to another
• Solid changes to liquid during melting
• Includes Vaporization, Condensation, Sublimation

Melting

• Solid to liquid transition
• Energy increases random motion of water molecule
• The temperature at which the rate is the same as the freezing rate at a given pressure is at equilibrium, called the melting point

Vaporization

• Liquid boils and turns to gas with continued heating
• Boiling point occurs
• Liquids and solids have the same vapor pressures at melting point

Sublimation

• Solid is vaporized at atmospheric pressure without going through a liquid phase
• Solids with high vapor pressure easily sublime
• Household mothballs and deodorizers undergo sublimation

Solutions Classified by Physical State

• Solutions that, on a microscopic level, are homogenous

Homogeneous Liquid

• Liquid is the primary component with small amounts of additional components
• Solutions are also classified by physical state once solutes are dissolved in solvents

Types of Solutions

• Solid: jewelry, coins, bronze, brass, etc.
• Liquid: soft drinks, coffee, juices, mineral water, etc.
• Gaseous: air

Components

• Solvent: dissolves another substance
• Solute: being dissolved
• The component present in a larger quantity is considered the solvent if they are in the same state

Types Based on State

• Gas in Gas
• Liquid in Liquid
• Solid in Liquid
• Gas in Liquid
• Solid in Solid
• Gas in Solid

Concentration

• Classifies the concentration or ratio of the solute and solvent
• Dilute: small amount of solute in a solvent
• Concentrated: large amount of solute in a solvent

Saturation, Solubility & Equilibrium in Solutions

• Solubility is the amount of solute that can be dissolved in a quantity of solvent to be come saturated at a particular temperature
• Saturated: solution contains the maximum amount of solute
• Supersaturated: contains more dissolved solute
• Unsaturated: contains less than the maximum amount of solute

Solution Process

• Process begins as solute is dissolved in solvent
• Molecule/ion is surrounded by new molecules
• A solute is considered "solvated" when this occurs with water, it is "hydrated"

Factors in Solution formation processes

• Solute-solute, solvent-solvent, and solute-solvent intermolecular forces

Solution Composition

• Concentration is a general term describing the amount of solute in a solution.
• Concentration can be expressed as Parts-per concentration
• Solution parts may be expressed in either mass, volume, or both

Mass/Mass Concentration

• Percentage refers to amount of solute (in grams) out of 100 grams of solution
• Equation to find % (w/w) = (Mass of solute)/(Mass of solution) * 100

Volume/Volume Percent

• Percentage refers preparation of a liquid solutions measured in liquids
• Solute concentration amount measured in volume units

Explore intermolecular forces, their impact on melting and boiling points, and phase transitions like vaporization. Understand the relationship between intermolecular forces and physical properties of compounds. Learn about London dispersion forces, dipole-dipole interactions & ion-dipole interactions.