Overview
This lecture discusses intermolecular forces, focusing on dipole-dipole interactions and hydrogen bonding. It explains how these forces operate between molecules, and examines their impact on boiling points and solubility in water. Several examples are provided to illustrate these concepts, including acetone, carbon monoxide, water, ammonia, methanol, ethanol, dimethyl ether, butanol, octanol, pentane, and neopentane.
Dipole-Dipole Interactions
- Dipole-dipole interactions occur between polar molecules, which have regions of partial positive and partial negative charge.
- Example: In acetone, the oxygen atom (with two lone pairs) carries a partial negative charge, while the central carbon atom has a partial positive charge. When two acetone molecules are near each other, the negative oxygen of one is attracted to the positive carbon of the other, creating a dipole-dipole interaction.
- Carbon monoxide is another example: the carbon atom has a partial positive charge and the oxygen atom a partial negative charge. When two CO molecules are close, the positive end of one is attracted to the negative end of the other, resulting in a dipole-dipole interaction.
- These interactions occur between separate molecules, not within a single molecule.
- The attraction is due to the general principle that opposite charges attract, pulling the molecules together.
Hydrogen Bonding
- Hydrogen bonding is a special, stronger type of dipole-dipole interaction.
- It occurs when hydrogen is directly bonded to nitrogen, oxygen, or fluorine.
- Example: In water, the oxygen atom has a partial negative charge and the hydrogen atom a partial positive charge. The oxygen of one water molecule is attracted to the hydrogen of another, forming a hydrogen bond.
- Hydrogen bonds are responsible for holding water molecules together and are always between separate molecules, not within a single molecule.
- Hydrogen bonding significantly increases both the boiling point and the solubility of a substance in water.
Effects on Boiling Point and Solubility
- Molecules with hydrogen bonds, such as ammonia (NH₃) and methanol (CH₃OH), have higher boiling points and are highly soluble in water due to strong intermolecular attractions.
- Ethanol (C₂H₅OH) contains hydrogen bonds, making it highly polar, with a higher boiling point and greater water solubility than dimethyl ether (CH₃OCH₃), which, despite being polar, lacks hydrogen bonds because its hydrogens are not directly attached to oxygen.
- Ethanol’s boiling point: ~78°C; Dimethyl ether’s boiling point: -23°C.
- As the hydrocarbon chain length increases (e.g., from ethanol to butanol to octanol), boiling point rises due to increased London dispersion forces, but water solubility decreases because the nonpolar region becomes larger.
- Ethanol is highly soluble in water; butanol is less soluble; octanol, with a long nonpolar chain, is not soluble in water and mixes better with nonpolar substances like oil.
- The presence of an -OH group (as in alcohols) increases water solubility, but as the number of nonpolar C-H bonds increases, solubility drops.
- Methanol, with the smallest nonpolar region, has the highest solubility in water among these alcohols, while octanol, with the largest nonpolar region, has the highest boiling point.
Structure and Boiling Point of Alkanes
- Pentane and neopentane are constitutional isomers: they have the same molecular formula (C₅H₁₂) but different structures.
- Straight-chain alkanes (like pentane) have higher boiling points than branched alkanes (like neopentane) because straight chains have more surface area, leading to stronger London dispersion forces.
- Branched alkanes (neopentane) have less surface area, resulting in weaker intermolecular forces and lower boiling points.
- Neither pentane nor neopentane is soluble in water, as they lack -OH groups and are nonpolar.
- In general, molecules made only of carbon and hydrogen (e.g., methane, ethane, propane) are nonpolar and do not mix with water.
Key Terms & Definitions
- Intermolecular Forces: Forces or interactions that occur between separate molecules, not within a single molecule.
- Dipole-Dipole Interaction: Attraction between the positive end of one polar molecule and the negative end of another.
- Hydrogen Bonding: A strong type of dipole-dipole interaction that occurs when hydrogen is bonded to nitrogen, oxygen, or fluorine, and is attracted to these atoms in neighboring molecules.
- London Dispersion Forces: Weak, temporary attractions caused by momentary shifts in electron distribution in nonpolar molecules; increase with molecular size and surface area.
- Isomers: Molecules with the same molecular formula but different structural arrangements (connectivity).
Action Items / Next Steps
- Review the molecular structures of acetone, water, ethanol, dimethyl ether, butanol, octanol, pentane, and neopentane to identify polar regions and possible hydrogen bonding.
- Practice identifying the types of intermolecular forces present in various molecules, focusing on recognizing hydrogen bonds, dipole-dipole interactions, and London dispersion forces.
- Compare boiling points and water solubility of molecules with similar formulas but different structures, paying attention to the effects of hydrogen bonding, chain length, and branching.