Intermolecular Forces Liquids and Solids CHAPTER 11 CHEM

Intermolecular Forces, Liquids and Solids CHAPTER 11 CHEM 160

A Molecular Comparison of Liquids and Solids • Physical properties of substances understood in terms of kinetic molecular theory: – Gases are highly compressible, assumes shape and volume of container: • Gas molecules are far apart and do not interact much with each other. – Liquids are almost incompressible, assume the shape but not the volume of container: • Liquids molecules are held closer together than gas molecules, but not so rigidly that the molecules cannot slide past each other.

A Molecular Comparison of Liquids and Solids – Solids are incompressible and have a definite shape and volume: • Solid molecules are packed closely together. The molecules are so rigidly packed that they cannot easily slide past each other.

A Molecular Comparison of Liquids and Solids

A Molecular Comparison of Liquids and Solids

A Molecular Comparison of Liquids and Solids • Converting a gas into a liquid or solid requires the molecules to get closer to each other: – cool or compress. • Converting a solid into a liquid or gas requires the molecules to move further apart: – heat or reduce pressure. • The forces holding solids and liquids together are called intermolecular forces.

Intermolecular Forces • The covalent bond holding a molecule together is an intramolecular forces. • The attraction between molecules is an intermolecular force. • Intermolecular forces are much weaker than intramolecular forces (e. g. 16 k. J/mol vs. 431 k. J/mol for HCl). • When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).

Intermolecular Forces

Intermolecular Forces Ion-Dipole Forces • Interaction between an ion and a dipole (e. g. water). • Strongest of all intermolecular forces.

Intermolecular Forces • • • Dipole-Dipole Forces Dipole-dipole forces exist between neutral polar molecules. Polar molecules need to be close together. Weaker than ion-dipole forces. There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

Intermolecular Forces Dipole-Dipole Forces

Intermolecular Forces Dipole-Dipole Forces

Intermolecular Forces • • • London Dispersion Forces Weakest of all intermolecular forces. It is possible for two adjacent neutral molecules to affect each other. The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). For an instant, the electron clouds become distorted. In that instant a dipole is formed (called an instantaneous dipole).

Intermolecular Forces London Dispersion Forces • One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom). • The forces between instantaneous dipoles are called London dispersion forces.

Intermolecular Forces • • • London Dispersion Forces Polarizability is the ease with which an electron cloud can be deformed. The larger the molecule (the greater the number of electrons) the more polarizable. London dispersion forces increase as molecular weight increases. London dispersion forces exist between all molecules. London dispersion forces depend on the shape of the molecule.

Intermolecular Forces London Dispersion Forces • The greater the surface area available for contact, the greater the dispersion forces. • London dispersion forces between spherical molecules are lower than between sausage-like molecules.

Intermolecular Forces London Dispersion Forces

Intermolecular Forces London Dispersion Forces

Intermolecular Forces Hydrogen Bonding • Special case of dipole-dipole forces. • By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. • Intermolecular forces are abnormally strong.

Intermolecular Forces Hydrogen Bonding • H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N). – Electrons in the H-X (X = electronegative element) lie much closer to X than H. – H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X. – Therefore, H-bonds are strong.

Hydrogen Bonding

Hydrogen Bonding

Intermolecular Forces Hydrogen Bonding • Hydrogen bonds are responsible for: – Ice Floating • • Solids are usually more closely packed than liquids; Therefore, solids are more dense than liquids. Ice is ordered with an open structure to optimize H-bonding. Therefore, ice is less dense than water. In water the H-O bond length is 1. 0 Å. The O…H hydrogen bond length is 1. 8 Å. Ice has waters arranged in an open, regular hexagon. Each + H points towards a lone pair on O.

Intermolecular Forces Hydrogen Bonding

Intermolecular Forces

Some Properties of Liquids Viscosity • Viscosity is the resistance of a liquid to flow. • A liquid flows by sliding molecules over each other. • The stronger the intermolecular forces, the higher the viscosity. Surface Tension • Bulk molecules (those in the liquid) are equally attracted to their neighbors.

Some Properties of Liquids Viscosity

Surface Tension

Some Properties of Liquids Surface Tension • Surface molecules are only attracted inwards towards the bulk molecules. – Therefore, surface molecules are packed more closely than bulk molecules. • Surface tension is the amount of energy required to increase the surface area of a liquid. • Cohesive forces bind molecules to each other. • Adhesive forces bind molecules to a surface.

Some Properties of Liquids Surface Tension • Meniscus is the shape of the liquid surface. – If adhesive forces are greater than cohesive forces, the liquid surface is attracted to its container more than the bulk molecules. Therefore, the meniscus is U-shaped (e. g. water in glass). – If cohesive forces are greater than adhesive forces, the meniscus is curved downwards. • Capillary Action: When a narrow glass tube is placed in water, the meniscus pulls the water up the tube.

Phase Changes • Surface molecules are only attracted inwards towards the bulk molecules. • Sublimation: solid gas. • Vaporization: liquid gas. • Melting or fusion: solid liquid. • Deposition: gas solid. • Condensation: gas liquid. • Freezing: liquid solid.

Phase Changes

Phase Changes • • • Energy Changes Accompanying Phase Changes Sublimation: Hsub > 0 (endothermic). Vaporization: Hvap > 0 (endothermic). Melting or Fusion: Hfus > 0 (endothermic). Deposition: Hdep < 0 (exothermic). Condensation: Hcon < 0 (exothermic). Freezing: Hfre < 0 (exothermic).

Phase Changes Energy Changes Accompanying Phase Changes • Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization: – it takes more energy to completely separate molecules, than partially separate them.

Phase Changes

Phase Changes Energy Changes Accompanying Phase Changes • All phase changes are possible under the right conditions. • The sequence heat solid melt heat liquid boil heat gas is endothermic. • The sequence cool gas condense cool liquid freeze cool solid is exothermic.

Phase Changes Heating Curves • Plot of temperature change versus heat added is a heating curve. • During a phase change, adding heat causes no temperature change. – These points are used to calculate Hfus and Hvap. • Supercooling: When a liquid is cooled below its melting point and it still remains a liquid. • Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.

Phase Changes Critical Temperature and Pressure • Gases liquefied by increasing pressure at some temperature. • Critical temperature: the minimum temperature for liquefaction of a gas using pressure. • Critical pressure: pressure required for liquefaction.

Phase Changes Critical Temperature and Pressure