unit 3: intermolecular forces

intermolecular forces

forces that act between molecules

london dispersion forces/dispersion forces

temporary asymmetric distribution of electrons that cause short-lived dipoles in all atoms and molecules

dipole-induced dipole interactions

a polar molecule induces a diple in a nearby nonpolar molecule

dipole-dipole interactions

attraction between poles of two polar molecules

hydrogen bonding

the interactions between a hydrogen covalently bonded to an N, O, or F molecule with another N, O, or F atom on another molecule

states of matter

atoms and molecules with stronger intermolecular forces require more energy to pull apart

vapor pressure

the pressure exerted by vapor at thermodynamic equilibrium with its liquid/solid phase in a closed system

ionic solids

• form a crystal/lattice structure and have high melting and boiling points

  • coulomb’s law applies → larger size of the atoms in the ions decreases lattice energy

    • lattice energy increases as ion charge increases and as the distance between atoms decreases (smaller atoms have higher lattice energy)

  • poor electrical conductors as a solid → but, good conductors when aqueous or molten since the ions can move freely

  • typically brittle and crack/shatter when enough pressure/stress is applied because the ions shift out of their lattice

covalent network

• each atom is connected to others with strong covalent bonds

  • carbon & silicon usually do this since they can form 4 bonds

  • high melting/boiling points and poor electrical conductivity (in all forms)

metallic solids

• form a crystal array of nuclei with core electrons while the valence electrons are delocalized and move throughout the solid

  • very good conductors

  • malleable, ductile; don’t crack/break because the metallic bond is nondirectional

    • the electrons aren’t affected much by force so they contribute to the malleable metallic properties

  • high melting/boiling points

molecular solids

• not held together by chemical bonds, but instead are held together by weak IMFs

  • low melting/boiling points, poor conductors

solids

• come in 2 structure types:

  • crystalline: highly ordered, repeating structure

  • amorphous: lacking a single repeating pattern

• both types are hard and rigid

ideal gas law

• an equation used to approximate the macroscopic properties of gases

  • PV = nRT

    1. P = pressure (in atm/mmHg/torr, 1 atm = 760 mmHg = 760 torr)

    2. V = volume (usually liters)

    3. n = number of gas particles (moles)

    4. R = gas constant → MUST match pressure units

    5. T = temperature (in Kelvin → add 273 to Celsius units)

total pressure

in a gas mixture, (P total) is equal to the sum of the partial pressures of each gas in the mixture

partial pressure of a gas

the pressure exerted by one gas in a mixture

kinetic molecular theory

the KMT for gases is a model that helps explain the macroscopic properties of gases

ideal gas assumptions

1. temperature is roughly equal to kinetic energy (K = 1/2mV^2)

2. gas particles have no IMFs

3. gas particles take up negligible volumes

4. gas particles are in constant random motion and do not lose energy when they collide

diffusion

the movement of gas particles from areas of high concentration to low concentration

effusion

the movement of gas particles through a small hole/mean free path (only one particle can move through at a time)

deviations from ideal behavior!

1. intermolecular forces become significant, causing the gas particles to stick together, resulting in less-than-ideal pressures

   • extremely high pressures where gas particles are pressed close together

   • low temperatures or temperatures close to boiling point → cannot escape IMF pulls

   • gas particles with stronger IMFs deviate more often

2. the volume of gas particles becomes significant at extremely tiny volume/pressure

heterogeneous mixture

mixture of 2+ substances with an inconsistent composition throughout

homogenous mixture

solution, a mixture of 2+ substances where the composition is consistent throughout

molarity

number of moles of solute in 1 liter of solution

chromatography

• separates chemical species by taking advantage of IMF strength within and among the components of the solution

  • separates a mixture into component parts

paper chromatography

• the mixture sample is placed on a piece of filter paper, which is dipped into a liquid

  • solvent moves up the paper through capillary action

mobile phase

stationary phase

mixture sample

retention (Rf) value

solute distance traveled ÷ solvent distance traveled (will be <1)

thin layer chromatography

replace paper with special-coated, thin glass

column chromatography

• components travel down instead of up

  • fractions/elutions: collected molecules

distillation

a laboratory technique that separates a mixture of 2+ liquids into their component parts by tracking advantage of differing boiling points

distillate

collected liquid

electromagnetic radiation

• has properties of both waves and particles

  • travels at the speed of light

  • quantized energy, comes in photons

microwave absorption

can change the rotational level of the molecule

infrared wave absorption

can change the vibrational energy level of the molecule

ultraviolet wave absorption

can increase the electron energy level/move electrons to an excited state

photoelectric effect

the emission of electrons (photoelectrons) when electromagnetic radiation hits a material (usually solid elemental metal)

E = hv

(energy = frequency x planck’s constant)

c = λv

(speed of light constant = wavelength x frequency)

spectrophotometer

used to measure how solute concentration in a solution affects the amount of light absorbed by a solution