Notes
Intermolecular forces (IMFs)
| Force | Found in | Strength |
|---|---|---|
| London dispersion | All molecules; bigger, more polarizable electron clouds have stronger LDFs | Weakest per interaction, but can add up |
| Dipole-dipole | Polar molecules | Medium |
| Hydrogen bonding | H bonded to N, O or F, attracted to another N, O or F | Strongest IMF |
| Ion-dipole | Ions dissolved in polar solvents | Strong |
- Stronger IMFs mean higher boiling points, lower vapor pressure, and higher surface tension.
- Types of solids: ionic, molecular, metallic, and covalent network (diamond, $\mathrm{SiO_2}$), which has very high melting points.
Gases
- Ideal gas law: $PV = nRT$. Use kelvin for temperature.
- Dalton's law: total pressure is the sum of partial pressures. $P_A = X_A P_{total}$, where $X_A$ is the mole fraction.
- Kinetic molecular theory: average kinetic energy depends only on temperature. At the same temperature, lighter molecules move faster.
- Real gases deviate from ideal at high pressure (molecules take up space) and low temperature (attractions matter).
Solutions and separation
- Molarity $\displaystyle M = \frac{\text{mol solute}}{\text{L solution}}$. Dilution: $M_1V_1 = M_2V_2$.
- "Like dissolves like": polar dissolves polar; nonpolar dissolves nonpolar.
- Chromatography separates by different attractions to the stationary and mobile phases; distillation separates by boiling point.
Light and matter (spectroscopy)
- $E = h\nu$ and $c = \lambda\nu$. Microwave radiation changes rotation, infrared changes vibration, and UV/visible changes electron energy levels.
- Beer–Lambert law: $A = \varepsilon bc$. Absorbance is proportional to concentration, which lets you find an unknown concentration from a calibration line.
Topics in this unit
Each topic has its own short notes and practice.
- 3.1Intermolecular and Interparticle Forces
- 3.2Properties of Solids
- 3.3Solids, Liquids, and Gases
- 3.4Ideal Gas Law
- 3.5Kinetic Molecular Theory
- 3.6Deviation from Ideal Gas Law
- 3.7Solutions and Mixtures
- 3.8Representations of Solutions
- 3.9Separation of Solutions and Mixtures
- 3.10Solubility
- 3.11Spectroscopy and the Electromagnetic Spectrum
- 3.12Properties of Photons
- 3.13Beer-Lambert Law
Practice
Water forms hydrogen bonds.
Both are nonpolar; London dispersion forces increase with polarizability.
$\displaystyle V = \frac{nRT}{P} = 2.00(0.0821)(273) \approx 44.8$ L.
$P_{O_2} = 0.25 \times 4.0$.
Molecules are close together and attractions matter more.
$M_1V_1 = M_2V_2$ gives $V_2 = 200$ mL total, so add 150 mL.
Absorbance is proportional to concentration: half the absorbance means half the concentration.
Every carbon is covalently bonded in a continuous network.
Free response
Write your answer first, then open the scoring guide.
Free-response question
Ethanol ($\mathrm{C_2H_5OH}$) and dimethyl ether ($\mathrm{CH_3OCH_3}$) have the same molar mass.
(a) Which has the higher boiling point? Explain using intermolecular forces.
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Ethanol. It has an O–H bond, so it forms hydrogen bonds between molecules. Dimethyl ether has no H on its oxygen, so it has only dipole-dipole and dispersion forces, which are weaker.
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(b) A 0.500 mol sample of ethanol vapor is in a 10.0 L flask at 400. K. Calculate its pressure.
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$\displaystyle P = \frac{nRT}{V} = \frac{0.500(0.0821)(400.)}{10.0} = 1.64$ atm.
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(c) Would the measured pressure likely be higher or lower than the ideal value? Explain.
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Lower. Attractions between ethanol molecules (hydrogen bonds) pull molecules toward each other, so they hit the walls with less force than an ideal gas would.
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