3.1: Intermolecular Forces

Intermolecular forces(IMF): the attractions between atoms, ions, or molecules

Intramolecular forces: chemical bonds holding the atoms in a molecule together

London dispersion forces(LDFs):

• occurs in all substances

• as electrons move around a temporary dipole is formed

  • area of excess negative charge w/ lots of electrons & area of positive charge

• negative and positive charges attract nearby molecules/atoms

  • forms an induced dipole

• usually temporary

  • weak

• polarizability: ability of an atom to form a temporary or induced dipole

• stronger when atom/molecule has more electrons & when atom/molecule is bigger

  • halogens

Dipole Dipole:

• occurs between molecules that have a permanent dipole from it being polar

  • polar when electron distribution is asymmetrical

• similar to LDFs, but dipoles are permanent and forces are stronger

• stronger when substance is more polar

Hydrogen bonding:

• occurs in molecules that contain H-N, H-F, H-O bonds

• difference in electronegativity is large

• hydrogen is partially positive and F,O, or N is partially negative

Ion dipole:

• occurs between an ion and a neutral dipole

Ionic bonding:

• occur between metal and nonmetal atoms

• lose/gain electrons to form ions

• stronger when charges are larger and ions are smaller

  • due to coulomb’s law

Metallic bonding:

• occurs between metal atoms

• attractions due to metallic cations being attracted to a delocalized sea of valence electrons

• stronger with smaller cations and more valence electrons

3.2: properties of solids

4 basic types of solids

1. ionic solids

2. covalent network solids

3. molecular solids

4. metallic solids

classified by what type of component occupies lattice points

Ionic solids:

• have ions at lattice points

• strong interactions between ions

  • low vapor pressures

  • high melting/boiling points

• electrostatic attraction: attraction between positive and negative ions

• smaller ions & ions with higher charges will have stronger attractions

  • results in higher lattice energy values

• brittle

• conduct electricity when ions are mobile

  • when solid is melted or dissolved

• typically between metal cation and non metal anion

Covalent Network Solids:

• atoms at lattice points with strong covalent bonds

• only formed from non metals

• high melting points

  • due to having to break covalent bonds

• rigid and brittle

Molecular solids:

• composed of distinct, individual units of covalently bonded molecules attracted through weak IMFs

  • have molecules at lattice points

• molecules at lattice points composed of non polars

• low melting point

  • due to weak IMFs

• do not conduct electricity

  • valence electrons rightly held within covalent bonds

Metallic solids:

• consist of metallic crystals w/ spherical metal atoms packed together

• closed packed lattice of ions surrounded by sea of moving electrons

• movement of electrons allow good conductivity

• malleable & ductile

• Metal alloys: mixture of metals

  • keep a sea of electrons so they can conduct

  • Substitutional alloy

    • atoms of similar sizes

    • density falls between 2 metals

  • Interstitial alloy

    • smaller atoms fill space between larger atoms

    • more rigid

strengths:

molecular solids<ionic solids~=metallic solids<covalent solids

3.3: solids, liquids & gases

particulate diagrams can be used to show the different properties between solids, liquids, and gases

solids:

• particles don’t have enough energy to move freely

  • vibrating

• little space between particles

  • IMFs hold particles in place

• rigid, fixed shape & volume

• cannot be compressed

• location can be crystalline or amorphous

  • crystalline solids have repeating 3D structure

  • amorphous have disordered particles

liquids:

• particles close together

  • move around a tiny bit

• cannot be compressed

• temp. range determined by IMF

gases:

• gained enough energy to overcome the IMFs holding particles

• particles far apart and moving quickly

• easily compressed

• takes shape of container

• no regular arrangement of particles

3.4: ideal gas law

ideal gas law: PV = nRT

• P = pressure in atm

• V = volume in L

• n = moles of gas

• R = universal gas constant(0.08206 L atm/mol K)

• T = temperature in kelvin(celsius + 273)

molar mass(g/mol) = (density(g/L)*R*temp(kelvin))/pressure

dalton’s law of partial pressure: the sum of all the partial pressures of each gas in a mixture of gases is equal to the total pressure

• when gases collected “over water”

mole fraction: XA = moles A/total moles

partial pressure A = XA * total pressure

3.5: kinetic molecular theory

kinetic molecular theory(KMT): theoretical model that describes the nature of ideal gases

1. the volume of gas particles can be ignored because they are so small

2. gas particles are in constant, random motions

3. particles are assumed to have no attractive/repulsive forces between them(IMFs can be ignored)

4. average kinetic energy of a sample of a gas is proportional to the kelvin temp. of the gas(KE = 1/2mv²)

maxwell boltzmann distribution: shows the distribution of the kinetic energies of particles at a given temp.

graham’s law: rate 1/ rate 2 = sqrt(molar mass 1/ molar mass 2)

3.6: deviations from ideal gas law

gases behave ideally under ordinary conditions-high temp. & low pressure

KMT assumes volumes of gas molecules are insignificant

• under high temp. they are moving quickly & under low pressure they are distant from each other

not all real gases behave ideally at high pressures & low temps

• when pressure is increased the particles are closer together

  • volume of gas molecules are significant

• as temp. decreases the particles move slower

  • IMFs become significant

low temp

• molecules move slower & have less energetic collisions

• “clump” together more

• collide with container less

  • decreases pressure

non zero molecular volume makes the actual volume greater than predicted

intermolecular attractions make the pressure less than predicted