Unit 3: Properties of Substances and Mixtures

21 terms. Say what each one means, then open the card to check.

• Intermolecular forces

  Attractions between separate molecules. They are much weaker than the covalent bonds inside a molecule, and they set boiling points and solubility.

  Example. Boiling water breaks the attractions between H₂O molecules; the O–H bonds inside each molecule stay intact.

• London dispersion forces

  Attractions between temporary dipoles that form when electron clouds shift. Every molecule has them, and they grow with the number of electrons.

  Example. Iodine, I₂, is a solid at room temperature while fluorine, F₂, is a gas, because I₂ has far more electrons and a more polarizable cloud.

• Polarizability

  How easily a particle's electron cloud is distorted. Larger clouds with more electrons are more polarizable and have stronger London forces.

  Example. Xenon atoms are far more polarizable than helium atoms, so xenon boils at a much higher temperature.

• Dipole-dipole forces

  Attractions between the positive end of one polar molecule and the negative end of another.

  Example. Acetone molecules line up so that the partly negative oxygen of one is near the partly positive carbon of the next.

• Hydrogen bond

  A strong dipole attraction between an H bonded to N, O or F and a lone pair on an N, O or F atom of a nearby molecule.

  Example. In liquid ammonia, the H on one NH₃ molecule is attracted to the lone pair on the N of a neighbor.

• Vapor pressure

  The pressure of a vapor above its liquid when evaporation and condensation balance in a closed container. Weaker forces mean higher vapor pressure.

  Example. Acetone evaporates faster than water at room temperature and has the higher vapor pressure, because acetone has no O–H group and cannot hydrogen bond to itself.

• Covalent network solid

  A solid in which every atom is joined to its neighbors by covalent bonds in a continuous network, so it is very hard and melts very high.

  Example. Diamond is a network of C–C bonds that runs through the whole crystal, and it does not melt until well above 3000 °C.

• Molecular solid

  A solid made of separate molecules held by intermolecular forces, so it is usually soft, melts low and does not conduct.

  Example. Solid carbon dioxide, dry ice, turns to gas at −78 °C because only weak London forces hold its molecules.

• Ideal gas law

  PV = nRT links pressure, volume, moles and kelvin temperature for a gas whose particles have no volume and no attractions.

  Example. One mole of an ideal gas at 1.00 atm and 273.15 K takes up 1.00 × 0.08206 × 273.15 ÷ 1.00 = 22.4 L.

• Partial pressure

  The pressure one gas in a mixture would exert alone. By Dalton's law, the partial pressures add to the total.

  Example. If air at 1.00 atm is 21% oxygen by moles, the partial pressure of oxygen is about 0.21 atm.

• Kinetic molecular theory

  A model of gases: tiny particles in constant random motion, with elastic collisions and an average kinetic energy that depends only on kelvin temperature.

  Example. At the same temperature, helium atoms and xenon atoms have the same average kinetic energy, so the light helium atoms move much faster.

• Maxwell-Boltzmann distribution

  A graph of how many gas particles have each speed. At a higher temperature, or for a lighter gas, the curve is broader and its peak moves to a higher speed.

  Example. The speed curve for a sample of neon at 500 K peaks at a higher speed and is flatter than the curve for the same neon at 300 K.

• Real gas behavior

  Real gases depart from PV = nRT at high pressure, where particle volume matters, and at low temperature, where attractions matter.

  Example. Ammonia, a polar gas, behaves less ideally than helium near its condensing point, because its molecules attract one another.

• Molarity

  Moles of solute per liter of solution, written M. It is not moles per liter of solvent.

  Example. Dissolving 0.0300 mol of glucose and adding water to make exactly 0.250 L of solution gives 0.0300 ÷ 0.250 = 0.120 M.

• Dilution

  Adding solvent lowers the concentration but keeps the moles of solute the same, so M₁V₁ = M₂V₂.

  Example. Diluting 10.0 mL of 6.00 M HCl to 60.0 mL gives 6.00 × 10.0 ÷ 60.0 = 1.00 M.

• Like dissolves like

  A solute dissolves well when its attractions to the solvent are similar to the attractions within each pure substance.

  Example. Oil, which is nonpolar, does not mix with water, but it dissolves in hexane, which is also nonpolar.

• Chromatography and distillation

  Ways to separate a mixture by differences in attractions: chromatography by attraction to a stationary phase, distillation by boiling point.

  Example. Heating a mixture of acetone (boiling point 56 °C) and water (100 °C) boils the acetone off first, so it can be collected as the distillate.

• Photon energy

  The energy of one photon of light, E = hν = hc ÷ λ. Shorter wavelength means higher frequency and more energy.

  Example. A photon of 250 nm ultraviolet light carries 6.626 × 10⁻³⁴ × 2.998 × 10⁸ ÷ (250 × 10⁻⁹) = 7.95 × 10⁻¹⁹ J.

• Regions of the spectrum

  Microwaves change how molecules rotate, infrared changes how bonds vibrate, and visible and ultraviolet light move electrons to higher energy levels.

  Example. An infrared spectrum of a sample can show a C=O bond, because that bond absorbs at a typical vibration frequency.

• Beer-Lambert law

  A = εbc: absorbance equals molar absorptivity times path length times concentration, so at one wavelength and one cell absorbance is proportional to concentration.

  Example. If a 0.10 M solution of a dye has an absorbance of 0.30, a 0.20 M solution of the same dye in the same cell has an absorbance of 0.60.

• Wavelength of maximum absorbance

  The wavelength a substance absorbs most strongly. It is the top of the peak in the substance's absorbance spectrum.

  Example. A blue dye solution absorbs mostly orange light, so its absorbance spectrum peaks near 600 nm.