2024Chem. IMFs ↓↑

intermolecular forces (imf’s): attractive forces between molecules that hold substances in condensed states (liquid or solid).

• depends on imf strength (aka polarity + molar mass) !!

• high imf = liquid or solid @ room temp.

• low imf = gas @ room temp.

intramolecular force: attractive forces within molecules like metallic, covalent, or ionic bonding.

• holds a compound together !!

melting point / boiling point: ↑ bp and ↑ mp = ↑ imf (more energy is required to overcome imf’s going from solid to liquid).

• substances w/ bp’s LESS THAN 23°C are GASES at room temp.

• substances w/ bp’s BETWEEN 23°C — 100°C are LIQUIDS at room temp.

the IMF forces !!

• ion-ion (intramolecular force)

  • VERY strong.

  • ionic… salt… ALWAYS solid @ room temp.

• ion-dipole

  • occurs when an ionic substance is mixed with a polar (covalent) substance

  • aka. salt dissolution

• hydrogen bonding

  • polar molecules containing hydrogen DIRECTLY bond to ↑ electronegative atoms N, O, or F.

  • creates a “super” dipole-dipole force where the hydrogen atom acts as a “bridge” between two NOF atoms.

  • substances Hbonding capable tend to have ↑’er boiling points / melting points than comparable substances that aren’t Hbond capable.

  • Hbonds aren’t true chem. bonds but rather intermolecular associations

• dipole-dipole

  • ALL polar molecules have permanent dipoles and ↑’er IMF’s than comparably sized non polar molecules

  • determines solubility… “like dissolves like”…

• london dispersion forces (ldf’s)

  • present between ALL molecules, most significant in non-polar substances.

  • results from fluctuations in electron distributions of each electron cloud.

  • strength is determined by how polarizable it’s electron cloud is…

    • larger cloud = more polarizable = stronger LDF

  • neopentane vs. pentane

    • neopentane has a lower electron cloud surface area and is “bulky”

    • pentane has a higher electron cloud surface area and is “long”

↑ imf’s = ↑ bp (liq.) + ↑ mp (sol.) because more thermal energy is required to break the attractive forces BETWEEN the molecules in their condensed states.

surface tension: the energy required to raise surface area - ↑’s with increasing IMF strength.

• ex. h2o droplets form spheres because h2o has a tendency to minimize surface area to lower inherent molecular potential energy.

  • molecules @ liquid surface have ↑’er potential energy than interior molecules as the exterior ones have fewer “neighbors” to interact with… therefore making them less stable… therefore giving them higher potential energies.

• spherical shapes have the LOWEST surface energy to volume ratio of any shape.

viscosity: the resistance of a liquid to flow.

vapor pressure + dynamic equilibrium

• dynamic equilibrium is when the ratio of condensation and the ratio of vaporization becomes equal. the pressure of a gas in dynamic equilibrium with it’s liquid is it’s vapor pressure. the vapor pressure of a liquid depends on the imf’s present in the liquid and the temperature of the system.

• ↑ imf’s = ↓ vapor pressure

• ↑ vapor pressure = ↑ vaporization rate

vaporization (liquid to gas)

• thermal energy overcomes imf’s to produce a state/phase change. molecules are in constant motion due to thermal energy and ↑’er temperature = ↑’er average kinetic energy of the collection of molecules. molecules @ the surface of a liquid are held less tightly than they are in the interior due to fewer “neighbor” interactions. therefore, some break free of liquid phase and become a gas.

volatility of liquids:

• volatile liquids have WEAK imf’s, therefore evaporate at higher rates @ room temp.

• nonvolatile liquids have STRONG imf’s, therefore vaporize slower @ room temp.

condensation: the reverse process of vaporization… exothermic change

enthalpy of vaporization: the amount of heat required to vaporize or condense ONE mole of a liquid to a gas.

use q = mC△T for temperature change !!

• q = heat

• m = mass

• C = specific heat capacity

  • (how much energy is needed to raise temp. of single mole of a specific substance by a single celsius

• △T = final(T) - initial(T)

use q = n△H for a phase change

• q = heat

• n = moles

• (△H = q / n)

segment 1 (+ slope)

• temp△ NOT a phase△

• Csice = 2.09J / g • °C

  • q = 1.5E3J = 1.5KJ

segment 2 (0 slope)

• heat added doesn’t temp△ because the heat is absorbed by the phase△

• Hfus = 6.02kJ / mole

  • q = 6.02kJ

segment 3 (+ slope)

• calculate heat

• CsH2o = 4.18J / g • °C

  • q = 7.52kJ

segment 4 (0 slope)

• phase△ liquid to gas needs △Hvap enthalpy of vaporization.

  • △HvapH2o = 4.07kJ / mole

    • q = 4.07kJ

parts of a phase diagram:

• sublimation curve > fusion curve (vertical between) > vaporization curve

  • solid to the left of fusion curve

  • liquid to the right of fusion curve

• fusion curve doesn’t continue past triple point