Hydrogen Bonding

• Electronegativity and Polarity: Atoms like $N$, $O$, and $F$ are so electronegative that they pull electron density away from the hydrogen atom, leaving it with a significant partial positive charge ($\delta^+$).

• Electrostatic Attraction: The nearly 'naked' proton of the hydrogen atom is strongly attracted to the concentrated negative charge of a lone pair on a nearby electronegative atom.

• Directionality: Hydrogen bonds are most stable and strongest when the donor atom, the hydrogen atom, and the acceptor atom are arranged in a linear fashion ($180^\circ$ angle).

• Bond Density: The strength of the interaction is also determined by the number of available lone pairs and hydrogen atoms; for instance, water forms more extensive networks than ammonia due to its balanced ratio of donors and acceptors.

Identifying Potential Hydrogen Bonding

• Step 1: Check for Donors: Look for hydrogen atoms directly attached to $N$, $O$, or $F$. If hydrogen is only attached to $C$, it cannot act as a donor.
• Step 2: Check for Acceptors: Ensure the molecule (or a neighboring one) has an $N$, $O$, or $F$ atom with at least one lone pair of electrons.
• Step 3: Predict Physical Properties: If hydrogen bonding is present, expect higher boiling points, higher viscosity, and greater solubility in polar solvents like water compared to similar-sized non-polar molecules.

• Intermolecular vs. Intramolecular: Hydrogen bonds are typically forces between molecules (intermolecular), whereas covalent bonds are the forces within a molecule (intramolecular) that hold atoms together.

• The 'NOF' Rule: Always remember that hydrogen bonding only occurs with Nitrogen, Oxygen, and Fluorine. If you see a molecule like $CH_3F$, be careful: the $H$ is bonded to $C$, not $F$, so it cannot form hydrogen bonds with itself.

• Boiling Point Trends: If asked to explain why a smaller molecule has a higher boiling point than a larger one, check for hydrogen bonding. It is the most common reason for 'anomalous' jumps in boiling point graphs for Group 15, 16, and 17 hydrides.

• Solubility Logic: Substances that can form hydrogen bonds with water (like ethanol) are generally highly soluble, even if they have a non-polar 'tail', because the $H$-bonding overcomes other factors.

• The 'Any Hydrogen' Error: A common mistake is assuming any hydrogen atom can participate. Hydrogen atoms bonded to Carbon ($C-H$) are not polar enough to form hydrogen bonds.

• Bond Strength Confusion: Students often think hydrogen bonds are as strong as covalent bonds. In reality, they are about $10$ times weaker, which is why they can be broken by simple heating (boiling) while covalent bonds remain intact.

• Lone Pair Neglect: Forgetting that an acceptor needs a lone pair can lead to errors in predicting the maximum number of bonds a molecule can form.