IMFs and State Changes- Chapter 10

Nov 17 and 19 Lectures

Ideal Gas Law Assumptions

• There are no (or entirely negligible) intermolecular forces between the gas molecules

• The volume occupied by the molecules themselves is negligible

Intramolecular Forces

• covalent and ionic bonds

• provide structure and shape

Intermolecular Forces

• Forces between molecules

• Responsible for condensed phases (aka liquids and solids)

are covalent or ionic forces stronger when intramolecular?

ionic

are inter or intramolecular forces stronger?

intramolecular are stronger

types of intermolecular forces

• ion-dipole

• dipole-dipole

• hydrogen bonding

• London dispersion forces

ion-dipole forces

ionic solid dissolves in water (or any polar substance)

dipole-dipole forces

between two polar molecules (such as water)

• regions of partial positive charge are attracted to regions of partial negative charge on neighboring polar molecules

Hydrogen bonding

Extreme dipole-dipole

• H-N

• H-O

• H-F

Why is water special, with a high BP?

The O in H2O has two lone pairs, so can H-bond with two other water molecules

volatility

the tendency of a substance to vaporize or evaporate

vapor pressure

in a closed system, the equilibrium pressure where there is a dynamic equilibrium between liquid and vapor

polarizability

how easily a molecule’s electron cloud can be distorted

• larger e clouds= more polarizable= stronger IMFs

• larger e clouds are squishy, easier to distort

Dipole-Induced

Polar molecule inducing a temporary dipole charge in a nonpolar molecule

London Dispersion Forces (weakest force)

One atom’s positive nucleus is attracted to the other’s electrons, and vice versa

• Competing forces cause uneven electron distribution, causing temporary induced dipoles

• Occur between all molecules!!

the weaker the attractive forces, the ___ the vapor pressure

higher 

• aka more volatile

vapor pressure increases with an increase in ____

temperature

clausius clapeyron equation

Pvap= mmHg

C= constant characteristic of the liquid you’re using

vaporization (evaporation)

gas in equilibrium with liquid (of same compound)

bolzmann distribution main idea

at higher temperatures, more molecules tend to be higher in energy to the point where their energy overcomes that of the intermolecular forces (vaporization)

• Average energy depends on temperature

-∆H of condensation is the same as

∆H of vaporization

• measured in kJ/mol

condensation is ____ thermic

condensation is exothermic

• being a gas requires more energy (heat) and you give that up when you condensate

vaporization is ____thermic

endo

• this is why we sweat!! “evaporative cooling

• think of it as swallowing the heat E around in order to “fuel” evaporation

boiling point definition

when vapor pressure = external pressure

• therefore, boiling point increases with increased pressure

critical point

above the critical point, the interface between liquid and vapor disappears

• supercritical fluid- density like a liquid, viscosity like a gas

• “green solvent” (little footprint) used to extract caffeine from coffee beans

melting point definition

temperature at which solid is converted to liquid

• ∆H of fusion is equal to the opposite of ∆H of melting

∆H fusion

enthalpy change that occurs at the melting point when a solid melts

• magnitude of ∆H fusion depends on what is holding it together

sublimation

solid to gas  

general solubility rule

“like dissolves like”

you need a solubility table for determining which ___ compounds are soluble

ionic

SO IMPORTANT!!! a polar bond has a minimum electronegativity difference of 

0.4

are larger or smaller molecules more soluble

smaller

what makes something most soluble in water

• polar

• can form H-bonds with water

• smaller molecule

Henry’s law

gas solubility increases with increasing pressure

S∝P

gas solubility factors

• solubility increases when pressure increases (Henry’s Law)

• solubility decreases with increasing temperature (the kinetic energy of molecules increases. Therefore, more gas can escape from solution)

Capillary Action

the ability of a liquid to flow up a narrow tube unassisted against gravity, is the result of cohesive and adhesive forces. When the adhesive forces between the liquid and the narrow tube are greater than the cohesive forces between the liquid molecules, the liquid molecules in contact with the wall of the tube are drawn up the side of the tube. The cohesive forces in the liquid cause the liquid molecules not in contact with the tube walls to be pulled up the tube as well. The liquid rises in the tube until the capillary action is balanced by the force of gravity. In liquids in which the cohesive forces are greater than the adhesive forces, capillary action does not occur.

Surface Tension

Surface tension is the result of a liquid attempting to minimize its surface area. The molecules at the surface of a liquid have fewer neighboring molecules to interact with than molecules in the interior of the liquid which are surrounded by molecules on all sides. Because there are only molecules to the side and below the surface molecules, the intermolecular forces pull the molecules at the surface downward into the bulk of the liquid. This downward pull minimizes the surface area and results in an elastic‑like surface. The surface tension decreases as the intermolecular forces decrease.