intermolecular forces

intermolecular forces

intermolecular forces

The attractions between molecules are not nearly as strong as the intramolecular attractions (bonds) that hold compounds together

dispersion force

attractions between an instantaneous dipole and an induced dipole.

Factors Affecting strength of this force

• number of electrons in an atom (more electrons, more ______ force)

• size of atom or molecule/molecular weight

• shape of molecules with similar masses (more compact, less _______ force)

  • Molecules are not made up of ions only 

  • Molecules are not polar

dipole-dipole forces

• Polar molecules have a more positive and a more negative end–a dipole (two poles, δ⁺ and δ⁻).

• The oppositely charged ends attract each other.

  • Molecules are not made up of ions only

  • Molecules are polar

hydrogen bonding

• The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong.

• is an attraction between a hydrogen atom attached to a highly electronegative atom and a nearby small electronegative atom in another molecule or chemical group.

• arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.

• These atoms interact with a nearly bare nucleus (which contains one proton).

  • Molecules are not made up of ions only

  • Molecules are polar

  • A hydrogen atom is present and is bonded to either F, O, N

ion-dipole forces

• are found in solutions of ions.

• The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents.

  • Molecules are made up of only ions

  • Molecules are polar

boiling point

the temperature at which the liquid converts into a gas

viscosity

Resistance of a liquid to flow — it increases with stronger intermolecular forces

surface tension

Water acts as if it has a “skin” on it due to extra inward forces on its surface — it increases with stronger IMFA

vapor pressure

measure of the tendency of a material to change into the gaseous or vapor state

capillary action

The rise of liquids up narrow tubes

• The net result of 2 opposing sets of forces: cohesive and adhesive forces

cohesive forces

Intermolecular forces that bind similar molecules to one another

adhesive forces

Intermolecular forces that bind a substance to a surface

kinetic energy

• less when IMFA is strong

volume

• solids have strong IMFA so their ____ do not expand

vapor pressure

• not common among solids because of their strong IMFA; however, some solid particles have minimum energy to turn into a vapor without becoming liquid.

melting point

• the temperature at which a solid melts; the stronger the IMFA, the higher the ______

heat of fusion

• the amount of heat needed to melt a solid once it reached its melting point. 

crystalline

• particles are in highly ordered arrangement; regularly arranged; show regular shapes

  • Examples: metals, alloys, carbon, salts (e.g. NaCl and MgSO₄)

amorphous

• no particular order in the arrangement of particles.

  • Examples: glass, plastic, asphalt, and rubber

phase changes

• Happens when a substance either gain or lose heat

• Is the transition from one phase to another

• Generally occurs at a constant temperature when the pressure is constant

endothermic process

absorption of heat/energy

• Melting

• Vaporization    

• Sublimation 

exothermic process

loss / release of heat/energy

• Condensation

• Deposition

• Freezing

heating curve

a plot of temperature and heat added to the substance

cooling curve

a plot of temperature and heat removed from the substance

solid-liquid equilibrium

when solid and liquid phases of a substance coexist.

• When a solid is heated it starts melting at a certain fixed temperature (melting point). 

• At this stage even when the heating is continued, the temperature does not change until the whole of solid is converted into liquid.

liquid-vapor equilibrium

when liquid and gas phases of a substance coexist

• When a liquid is heated it starts evaporating at a certain fixed temperature . At this stage even when the heating is continued, the temperature does not change until the whole of liquid is converted into gas.

state of equilibrium

reached when two substances, one solid and one vapor, are in contact with each other in a closed system

• In this state, the concrete's vaporization rate is equal to the speed of condensation of the vapor

phase diagram

Shows the physical states of a substance under different conditions of temperature and pressure

melting (or freezing) curve

• represents the transition between liquid and solid states

vaporization (or condensation) curve

• represents the transition between gaseous and liquid states

sublimation (or deposition) curve

• represents the transition between gaseous and solid states.

triple point

• a point at which all states of matter coexist

critical point

the temperature and pressure at which the distinction between liquid and gas can no longer be made

supercritical fluid

a phase that occurs for a gas at a specific temperature and pressure such that the gas will no longer condense to a liquid regardless of how high the pressure is raised.

critical temperature

• the highest temperature at which the substance can remain in liquid state

critical pressure

the pressure required to liquefy a gas at its critical temperature