Unit 3 Part 1

Intermolecular and Interparticle Forces, Properties of Solids, Solids, Liquids, and Gases, Ideal Gas Law, Kinetic Molecular Theory, Deviation from Ideal Gas Law

intermolecular forces

forces of attraction between two molecules

london dispersion forces

temporary attractions found in all molecules but only force found in nonpolar molecules; weak attractions

types of LDFs

temporary dipole, induced dipole

temporary dipole

wobbly electrons shift to one atom, creating temporary partial charges

induced dipole

if something with permanent dipoles gets close to a nonpolar molecule, electrons will shift and create partial charges

how to compare strength of LDFs

the larger the molecule, the stronger the LDFs (nucleus has less pull on electrons so they move around more)

dipole-dipole

formed between permanent dipoles in polar molecules; stronger than LDFs in equivalent molecules

hydrogen bonding

found in molecules where H atom is bonded to N, O, or F; type of dipole-dipole but stronger attraction

ion dipole

cation attracts to partial negative charge, anion attracts to partial positive charge

ion induced dipole

the charge of the ion induces dipoles in a nonpolar molecule

types of solids

ionic, molecular, metallic, network

properties of ionic solids

ionic bonds, high MP, hard and brittle, poor conductors

properties of molecular solids

IMFs, low MP, poor conductors, soft

properties of metallic solids

metallic bonding, good conductors, ductile and malleable, variable MP, lustrous (shiny)

properties of network solids

covalent bonding, high MP, hard, poor conductors, very hard and brittle

examples of network solids

diamond (all C), quartz (SiO₂)

solids

definite shape and volume, strong IMFs, incompressible, molecules vibrate in place

liquids

indefinite shape, definite volume, weaker IMFs, molecules vibrate more & able to flow

gases

indefinite shape and volume, easily compressible, no IMFs, molecules freely moving

boyle’s law

when temperature constant, pressure and volume inversely related (P₁V₁ = P₂V₂)

charles’s law

when pressure constant, volume and temperature directly related (V₁/T₁ = V₂/T₂)

gay-lussac’s law

when volume constant, pressure and temperature directly related (P₁/T₁ = P₂/T₂)

avogadro's law

1 mol = 22.4L at STP

combined gas law

P₁V₁/T₁ = P₂V₂/T₂

daltons law of partial pressures

P_total = P₁ + P₂ + P₃ + …

partial pressure

pressure exerted by one gas in a mixture

mole ratio

moles of substance / total moles, denoted by chi

relationship between moles and partial pressure

mole ratio = partial pressure / total pressure

ideal gas law

PV = nRT

R constant

0.08206 L*atm / k*mol

conversion for pressure

1 atm = 101.3 kPa = 760 mmHg = 760 torr

molar mass derivation of ideal gas law

M = mRT/PV

density derivation of ideal gas law

D = MP/RT

kinetic molecular theory

1. gases are in continuous, random motion

2. gases have no volume

3. gases have no IMFs

4. when particles collide, energy is transferred—no net gain/loss of energy within the system

5. average KE is proportional to temperature

average KE formula

1/2mv²

average speed of smaller molecules when temp is constant

faster

KE is most dependent on

temperature

how to read distribution graph

x-axis is speed, y-axis is # of molecules; peak represents average speed of the molecules

on the distribution graph, peak of smaller gases is

further to the right (higher average speed)

on the distribution graph, peak of larger gases will be

further to the left (lower average speed)

assumptions for ideal gas

no volume and no IMFs

deviations from ideal gas law

strong IMFs, lower temperature, high pressure

what effect does the deviation from ideal gas law have

pressure is less than predicted

why is pressure less than predicted in real gases?

molecules form IMFs, therefore less number and force of collisions = less pressure

conditions for ideal gas

high temperature, low pressure

conditions for real gas

low temperature, high pressure