CHM135 Chapter 8

London Dispersion Forces

-how attractive forces arise among nonpolar molecules or individual atoms of a noble gas -all atoms and molecules experience, result from motion of electrons -instantaneous dipole on one molecule can affect electron distributions in neighbouring molecules and induce temporary dipoles in neighbours (weak attractive forces develop)

Dipole-Dipole Forces

-neutral but polar molecules experience as result of electrical interactions among dipoles on neighbouring molecules -can be attractive or repulsive, depending on orientation -weaker than ion-dipole -strength depends on size of dipole moments involved -->more polar = greater strength -larger dipole moment = stronger intermolecular forces = greater amount of heat that must be added to overcome forces -->substances with larger dipole moments have higher boiling points

Ion-Dipole Forces

-result of electrical interactions between an ion and partial charges on a polar molecule -favoured orientation of polar molecule in presence of ions: positive end of molecular dipole near anion, negative end near cation -highly variable in strength, particularly important in aqueous solutions of ionic substances, in which polar water molecules surround the ions

Permanent Dipole

-2 things to think about: 1. polar bond: electron distribution 2. 3D shape

Induced Dipole

-London dispersion forces or van der Waals forces -all molecules have this force

Hydrogen Bonds

-attractive interaction between hydrogen atom bonded to very electronegative atom (O, N, F) and electron-rich region elsewhere in same molecule or in different molecule -higher boiling points -strong because hydrogen doesn't have any inner core electrons -end up with very positive charge on molecule because of nature of hydrogen bonded to very electronegative atom -arise because bonds are highly polar, with partial positive charge on H atom and partial negative charge on other electronegative atom -->H atom has no core electrons to shield nucleus and has small size, can be approached closely by other molecules -->dipole-dipole attraction between H and unshared electron pair on nearby atom is unusually strong, giving rise to H bond

Intermolecular Forces

-forces that occur between molecules -different than covalent bonds -influence macroscopic properties of matter (solubility, melting and boiling points) -play key role in stabilizing shapes and interactions of biomolecules -all are electrostatic in origin, result from mutual attraction of unlike charges or mutual repulsion of like charges

Strength of London Dispersion Forces

-generally small, exact magnitude depends on molecule's polarizability -->smaller molecule or lighter atom is less polarizable, has smaller dispersion forces because it only has a few, tightly held electrons -->larger molecule/heavier atom is more polarizable, has larger dispersion forces because it has many electrons, some of which are less tightly held and farther from nucleus -shape also important in determining magnitude of dispersion forces -->spread-out shapes (max surface area) allow greater contact between molecules, higher dispersion forces

Polarity

Driving force between interactions between molecules is electrostatic (positive attracted to negative) -influences both chemical and physical properties of molecules -molecules as a whole often polar because of net sum of individual bond polarities and lone pair contributions

Polarizability

Ease with which molecule's electron cloud can be distorted by a nearby electric field

van der Waals forces

Intermolecular forces as a whole, including dipole-dipole, London dispersion, hydrogen bonds (ion-dipole forces also exist between ions and molecules)

Electronegativity

Measure of how an atom pulls electrons within a bond

Dipole Moment

Measure of net molecular polarity, defined as magnitude of charge Q at either end of molecular dipole times distance r between charges (u = Q x r) -->expressed in debytes (D) -once known, can get idea of amount of charge separation in molecule -largest: NaCl -zero: CO2 and tetrachloramethane

Boiling Point

Temperature at which vapour pressure of liquid is equal to atmospheric pressure -->when we go down from F to I, boiling point increases -->higher boiling point means it takes more energy for molecules in liquid to escape gas (higher temperature = more kinetic energy) -->increased intermolecular interactions