unit 3 #

  • intermolecular forces are the forces between molecules (H₂-H₂)

    • london dispersion forces (LDF)

      • alllllll molecules exhibit LDFs (even polar molecules!!!)

      • caused by a temporary dipole (also is temporary)

        • also called an induced dipole-induced dipole attraction (when a polar molecule induces a dipole in a nonpolar molecule, between polar and nonpolar!!!)

      • weakest intermolecular forces

      • the strength of the LDF depends on how easily the electrons can disperse (polarizability, discussed in 2.1)

        • the larger the electron cloud, the stronger the LDF

        • so larger molecules exhibit stronger LDFs

        • same molar mass → same LDF strength

      • also depends on surface area. if two molecules have the same molar mass, but one has a larger surface area, that one exhibits greater LDFs

      • main (& only) force between nonpolar molecules (& noble gases!!!)

    • dipole-dipole forces

      • attractions to the opposite partial charges of other molecules

      • the molecules orient themselves in a way so that the δ+ side of one molecule is facing the δ- side of another (& vice versa)

        • this happens so attraction is maximized & repulsion is minimized

      • the strength of the force depends on the magnitude of the dipole

        • the greater the dipole moment, the greater the force

      • between polar molecules

    • hydrogen bonds

      • bonds between H & N/O/F

      • subtype of dipole-dipole force

        • they’re so much stronger than normal dipole-dipole forces so they were given their own name

      • strongest intermolecular force

      • between polar molecules

  • when a substance changes from solid to liquid to gas & vice versa, the molecules in that substance remain intact

    • these changes in state are due to the breaking of intermolecular forces, not intramolecular forces

      • H₂O(l) → H₂O(g)

      • same substance, no chemical alteration, just state change yo

    • so molecules with hydrogen bonds rather than dipole-dipole forces have significantly higher boiling points, because their intermolecular forces are stronger

    • the breaking of intramolecular forces results in chemical changes

  • dissolution is both physical and chemical by the way

    • physical because NaCl(s) → NaCl(aq)

    • chemical because the ionic bonds between Na⁺ and Cl⁻ have to be broken

  • types of solids:

    • ionic solids

      • discussed in 2.3

    • molecular solids

      • formed of nonmetals only

      • relatively low melting & boiling points

    • metallic solids

      • metallic bonding (valence electrons are in a sea of mobile / delocalized electrons)

      • good conductors

      • malleable & ductile

    • network covalent solids

      • formed by distinct atoms all bonded together covalently in a 3d network

        • FIXED BOND ANGLES

      • formed by carbon & metalloids (boron, silicon, germanium, etc)

        • SiO₂, SiC, BC

      • very high melting points

      • very hard

      • poor conductors

  • solids

    • molecules in fixed locations

    • molecules are vibrating in their places

  • liquids

    • molecules slide past each other

    • molecules at the surface can evaporate, they always are

  • gases

    • molecules move randomly in straight lines between collisions

    • barely any intermolecular forces

  • ideal gas: gas that strictly obeys the gas laws (mainly we mean the ideal gas law)

    • real gases can behave like ideal gases at high temperature & low pressure

  • the gas laws!!!

    • charles’ law

      • volume & temperature

        • pressure & number of moles constant

      • V ∝ T

      • V₁/T₁ = V₂/T₂

    • gay-lussac’s law

      • pressure & temperature

        • volume & number of moles constant

      • P ∝ T

      • P₁/T₁ = P₂/T₂

    • boyle’s law

      • pressure & volume

        • temperature & number of moles constant

      • P ∝ 1/V

      • P₁V₁=P₂V₂

    • avogadro’s law

      • volume & number of moles

        • pressure & temperature constant

      • V ∝ n

      • V₁/n₁=V₂/n₂

  • PV=nRT (ideal gas law)

    • P: pressure (atm)

    • V: volume (L)

    • n: number of moles (mol)

    • R: ideal gas constant / 0.082 (L.atm/mol.k)

    • T: temperature (kelvins)

  • Pᴛ=Pᴀ+Pʙ+Pᴄ+…

    • Pᴛ: total pressure of gases in the container

    • Pᴀ and Pʙ and Pᴄ are the partial pressures of gases a, b, and c

      • partial pressure: the pressure exerted by each gas if it were alone in the container

  • Xɪ=nɪ/nᴛ

    • Xɪ: mole fraction of gas i

    • nɪ: number of moles of gas i

    • nᴛ: total number of moles in the mixture

  • Xɪ=Pɪ/Pᴛ

    • Xɪ: mole fraction of gas i

    • Pɪ: partial pressure of gas i

    • Pᴛ: total pressure of gases in the container

  • absolute zero: 0 kelvins, -273.15°C (impossible to obtain)

  • kinetic molecular theory: explains the properties of an ideal gas

    • particles are so small compared to the distances separating them, so the volume of individual particles is negligible

    • particles are in constant motion → the collisions of the particles with the walls of the container are the cause of the pressure exerted by the gas

    • particles exert no force on each other (no attraction, no repulsion)

    • avg K.E. of a collection of gas particles ∝ temp in Kelvins

  • all particles are in continuous random motion

    • (vibration in solids, sliding in liquids, collisions and movement in gases)

  • gaseous particles move in straight lines with constant velocity between collisions

    • when the particles collide, they don’t stick together. this is because their collisions are elastic

  • K.E=m(v)²/2

    • K.E: kinetic energy

    • m: mass

    • v: velocity

  • so greater mass means more frequent collisions and shorter distance between collisions

  • difference in sizes of atoms is kinda negligible - like if you were to say that a molecule of hydrogen can fit through a hole better than a molecule of oxygen because it’s smaller, you’d be wrong. instead, the hydrogen passes through the hole faster because its molar mass is less than the molar mass of oxygen, so it has a higher speed

  • maxwell-boltzmann distributions: (x - particle speed (basically velocity) / y - frequency of that speed to be found within the atoms of a molecule (or molecules of a compound idk man)

    • hot/cold

      • cold: higher peak nearer to the origin (slower avg movement)

      • hot: more flat hill spread out over all speeds (more particles with higher speeds)

    • heavy/light

      • heavy: higher peak nearer to the origin (slower avg movement)

      • light: more flat hill spread out over all speeds (more particles with higher speeds bc lighter particles move faster)

  • the ideal gas law (PV=nRT) is only valid for ideal gases. the rule does not apply for real gases

  • gases are not ideal at:

    • low temperatures

    • high pressures

    • (or like particles with significant intermolecular forces or molecular sizes)

  • WE WANT:

    • HIGH TEMP

    • LOW PRESSURE

    • INSIGNIFICANT IMF

  • LARGER INTERMOLECULAR FORCES → LESS IDEAL GAS

  • solutions are homogeneous mixtures

  • C = n/V

    • C: concentration (M)

    • n: moles of SOLUTE (solid obv) (mol)

    • V: volume of SOLUTION (L or dm³)

  • when diluting: use C₁V₁=C₂V₂

  • particulate models!!!

  • in solutions specifically, the water molecules play a part in how the model looks or whatever

  • orientation of water molecules

    • the δ⁻ oxygen → oriented towards cation

    • the δ⁺ hydrogen → oriented towards anion

  • how do we separate the components of a liquid solution (eg ink)?

    • chromatography!!!

  • chromatography:

    • place a drop of the mixture to be separated on a piece of paper. draw a line (origin line) on the x axis that the dot is situated on. place this paper in a solvent that this mixture is soluble in. if the mixture isn’t soluble in the solvent, the dot won’t budge and nothing will happen. make sure that the dot (of the mixture) isn’t immersed in the solvent, otherwise it’ll dissolve as normal. the solvent will begin making its way up the paper, and so will the dot’s components. the components that appear on the paper at first are the ones that are least attracted to the solvent, meaning they have opposite polarity to the substance. the components that appear last are most attracted to the solvent, meaning they have the same polarity as the solvent. also means that these substances are most soluble in the solvent.

    • mobile phase: solution

    • stationary phase: paper

    • ps alcohols are polar (hydrogen bonding, also the OH makes them geometrically asymmetric)

  • distillation:

    • simple:

      • used to separate a solute and a solvent (solute dissolved in the solvent, obviously)

      1. boil mixture in flask

      2. liquid begins boiling & becomes gas

      3. liquid passes through the condenser to be condensed back into a liquid

      4. pure liquid collected in a beaker

      5. process repeated until no liquid is left in the beginning flask

    • fractional:

      • used to separate 2 miscible liquids with a close boiling point

      1. boil mixture in flask at a boiling point higher than both liquids’ boiling points

      2. they will evaporate and reach the fractionating column filled with glass beads. these glass beads are at a temperature between the two liquids’ boiling points. meaning they condense the liquid with the higher boiling point and allow the one with the lower boiling point to move on to the condenser.

      3. the liquid (or gas now, technically) with the lower boiling point enters the condenser, where it condenses back into a liquid

      4. that liquid is collected as the distillate in a beaker

  • remember that when assessing the boiling points of two compounds, do not discuss intramolecular forces (covalent/ionic bonds). those are not broken by boiling (they’re broken chemically). discuss intermolecular forces (LDF, hydrogen bonds, dipole-dipole, etc)

  • ionic solutes dissolve in polar solvents

  • polar solutes dissolve in polar solvents

  • nonpolar solutes dissolve in nonpolar solvents

  • ^ all these are due to intermolecular attractions between solutes and solvents

  • the electromagnetic spectrum:

    • gamma rays / x-rays / ultraviolet / light (visible) / infrared / microwaves / radio waves

    • ^ in order of decreasing energy

      • decreasing frequency

      • increasing wavelength

    • all travel in waves at the speed of light

  • ultraviolet/visible radiation: transitions in electronic energy levels (gives atoms the energy to jump from one energy level to another)

  • infrared radiation: transitions in molecular vibrational levels (causes atoms to vibrate, gives info about bonds/bond order)

  • microwave radiation: transitions in molecular rotational levels (causes atoms to rotate, gives info about polarity)

  • photoelectric effect: phenomenon in which electrons are emitted from the surface of a metal when light strikes it

  • wavelength ∝ 1/frequency

  • c = λv

    • c: speed of light (2.998×10⁸ m/s)

    • λ: wavelength (m)

    • v: frequency (Hz, which is s⁻¹)

  • E = hv

    • E: energy of the photon absorbed or emitted by an atom or molecule (J)

    • h: Planck’s constant (6.626×10⁻³⁴ J.s)

    • v: frequency (Hz, which is s⁻¹)

  • beer-lambert law: there is a linear relationship between the concentration & absorbance of a solution

  • A = εbc

    • A: absorbance

    • ε: molar absorptivity

    • b: light path

    • c: concentration (M)

  • molar absorptivity: how intensely a sample absorbs light of a specific wavelength

  • beer’s law plot:

    • x axis → concentration (M)

    • y axis → absorbance