Hydrogen Bonding

Definition and General Features of Hydrogen Bonds

Hydrogen bonding is a specific, relatively strong type of intermolecular interaction that occurs when:

• A hydrogen atom is covalently bonded to a strongly electronegative atom: typically nitrogen (N), oxygen (O), or fluorine (F).
• The same hydrogen atom interacts (attracted electrostatically) with a lone pair on another electronegative atom (often N, O, or F) in the same or a different molecule.

In shorthand, hydrogen bonds can be represented as:
$$
\text{D–H} \cdots \text{A}
$$
where:

• D = hydrogen bond donor atom (usually N, O, or F),
• H = hydrogen atom covalently bonded to D,
• A = hydrogen bond acceptor atom (again usually N, O, or F) with a lone pair,
• the dots ⋯ indicate the hydrogen bond interaction.

Key points:

• Hydrogen bonds are stronger than most other intermolecular interactions (e.g. typical van der Waals forces), but weaker than covalent bonds.
• They are directional: the strongest interactions occur when D–H–A are arranged nearly linearly.
• They can be intermolecular (between different molecules) or intramolecular (within the same molecule).

Donors, Acceptors, and Typical Examples

Hydrogen Bond Donors

A hydrogen bond donor is a group where hydrogen is directly attached to a highly electronegative atom. Common donors:

• O–H groups: in water, alcohols, carboxylic acids
• N–H groups: in amines, amides, many biomolecules (e.g. proteins)
• F–H: in hydrogen fluoride (HF)

Examples:

• Water: each $ \text{H}_2\text{O} $ molecule has two O–H groups → can donate two hydrogen bonds.
• Ammonia: $ \text{NH}_3 $ has three N–H bonds → can donate up to three hydrogen bonds (in principle).

Hydrogen Bond Acceptors

A hydrogen bond acceptor is an atom with a lone pair and sufficient electronegativity, commonly:

• O in water, alcohols, ethers, carbonyl compounds, etc.
• N in amines, amides, nitriles, etc.
• F in HF and some inorganic fluorides.

Examples:

• In water, the oxygen atom has two lone pairs → each molecule can accept two hydrogen bonds.
• In an amide, the carbonyl oxygen is a good hydrogen bond acceptor.

Typical Molecular Motifs

Some of the most important hydrogen-bonding motifs:

• Between water molecules: $ \text{H}_2\text{O}\cdots\text{H–O}\text{H} $
• Between alcohols: $ \text{RO–H}\cdots\text{:O–R} $
• Between carboxylic acids: formation of cyclic dimers via two hydrogen bonds
• In amides and peptides: $ \text{N–H}\cdots\text{:O}=\text{C} $ interactions

Energetics and Strength of Hydrogen Bonds

Hydrogen bond strengths vary depending on the atoms involved and the molecular environment, but typically:

• Energies are on the order of 10–40 kJ·mol⁻¹ (often quoted ranges; “strong” hydrogen bonds can be somewhat larger).
• They are weaker than covalent bonds (hundreds of kJ·mol⁻¹), but often stronger than individual dispersion or dipole–dipole interactions.

Factors influencing hydrogen bond strength:

• Electronegativity of donor and acceptor atoms: O–H···O and N–H···O are common and fairly strong.
• Charge: hydrogen bonding involving charged species (e.g. between $ \text{H}_3\text{O}^+ $ and water, or between carboxylate anions and neutral donors) can be significantly stronger.
• Geometry: nearly linear D–H···A arrangements are stronger than bent ones.
• Environment: solvent polarity, neighboring groups, and the presence of competing hydrogen bond partners influence how strong an individual bond effectively is.

Structural Consequences of Hydrogen Bonding in Liquids and Solids

Hydrogen bonding can organize molecules into characteristic three-dimensional networks or ordered motifs.

Liquid Water and Ice

In liquid water:

• Each water molecule can, in principle, form up to four hydrogen bonds: two as donor (via two O–H groups) and two as acceptor (via two lone pairs on O).
• At room temperature, thermal motion constantly breaks and reforms hydrogen bonds, generating a dynamic network.

In ice:

• Hydrogen bonds are more ordered and nearly fully realized in a tetrahedral coordination.
• This open, tetrahedral arrangement makes ice less dense than liquid water, so ice floats.

These structural properties are direct consequences of hydrogen bonding patterns.

Association in Molecular Liquids and Solids

Many molecules form associates (dimers, chains, networks) in the condensed phase:

• Carboxylic acids often form cyclic dimers via two hydrogen bonds:
  $$
  \text{R–C(=O)–OH} \cdots \text{HO–C(=O)–R}
  $$
• Alcohols can form chains or rings of intermolecular hydrogen bonds.
• Hydrogen bonding frequently results in:
• Higher boiling and melting points than expected based on molar mass alone.
• Increased viscosity in liquids with extensive hydrogen bonding.

Crystal Structures

In crystals:

• Hydrogen bonds act as structure-directing interactions.
• They can link molecules into:
• Linear chains
• Two-dimensional sheets
• Three-dimensional networks

Because they are relatively strong and directional, hydrogen bonds are widely used in crystal engineering and supramolecular chemistry to design specific solid-state structures.

Intramolecular Hydrogen Bonding

Hydrogen bonds can also form within a single molecule, when a hydrogen bond donor and an acceptor are appropriately positioned.

Characteristics and consequences:

• Intramolecular hydrogen bonds can create “rings” within the molecule (e.g. six-membered pseudo-rings).
• They can stabilize certain conformations by locking parts of the molecule into place.
• They may compete with intermolecular hydrogen bonding:
• If intramolecular hydrogen bonding is strong, the molecule may engage less in intermolecular hydrogen bonding, influencing solubility and melting point.

Examples (generic motifs):

• An –OH group located near a carbonyl group within the same molecule may form an O–H···O=C hydrogen bond.
• In some aromatic compounds (e.g. salicylic acid-type structures), an internal O–H···O hydrogen bond significantly affects properties such as acidity and solubility.

Hydrogen Bonding and Physical Properties

Hydrogen bonding often explains why certain substances have unexpectedly high:

• Boiling points and melting points
• Heat capacities
• Viscosities
• Surface tensions

compared to structurally similar molecules that cannot form hydrogen bonds.

Boiling and Melting Points

Comparisons illustrate the impact:

• Water ($ \text{H}_2\text{O} $) vs. hydrogen sulfide ($ \text{H}_2\text{S} $):
• Both are group 16 hydrides, but $ \text{H}_2\text{O} $ has a much higher boiling point due to extensive O–H···O hydrogen bonding; $ \text{H}_2\text{S} $ cannot form comparable hydrogen bonds.
• Alcohols vs. alkanes:
• Ethanol ($ \text{CH}_3\text{CH}_2\text{OH} $) has a much higher boiling point than propane ($ \text{C}_3\text{H}_8 $) despite similar molar mass, because ethanol forms O–H···O hydrogen bonds while propane cannot.

Solubility

Hydrogen bonding plays a major role in solubility:

• Molecules that can form hydrogen bonds with water (e.g. small alcohols, amines) are usually well soluble in water.
• Molecules lacking hydrogen bond donors and acceptors tend to be poorly soluble in water, though they may be more soluble in nonpolar solvents.

Interactions involved:

• Solute–solvent hydrogen bonds (e.g. O–H or N–H from a solute with water molecules).
• Replacement of solvent–solvent hydrogen bonds by solute–solvent ones.

Hydrogen Bonding in Biological and Organic Structures

Hydrogen bonding is crucial in stabilizing the structures of biomolecules and many organic frameworks.

Proteins

In proteins, hydrogen bonding:

• Forms between backbone amide N–H and carbonyl C=O groups:
  $$
  \text{N–H} \cdots \text{O=C}
  $$
• Stabilizes secondary structures such as:
• α-helices: regular, intra-chain hydrogen bonds between every $i$-th and $(i+4)$-th amino acid.
• β-sheets: hydrogen bonds between neighboring strands of the polypeptide chain.

These hydrogen bonds are directional and form regular patterns, giving proteins their characteristic shapes.

Nucleic Acids (DNA and RNA)

In DNA:

• Complementary base pairs are held together by specific hydrogen bonding patterns:
• Between adenine (A) and thymine (T): two hydrogen bonds.
• Between guanine (G) and cytosine (C): three hydrogen bonds.
• These interactions:
• Help stabilize the double helix.
• Ensure specific base pairing, essential for accurate replication and transcription.

In RNA, similar hydrogen bonds stabilize internal base pairing and folded structures.

Polysaccharides and Other Biopolymers

Hydrogen bonds between:

• Hydroxyl groups in polysaccharides (e.g. cellulose) lead to:
• Strong intermolecular interactions.
• High tensile strength and low solubility in water for cellulose.
• Side chains and backbone groups in other biopolymers similarly create ordered, stable structures.

Special Types and Borderline Cases

Not all hydrogen bonds are equally straightforward; some special or borderline cases are often discussed.

Strong (Low-Barrier) Hydrogen Bonds

In some systems:

• The hydrogen is shared nearly symmetrically between two very similar atoms (e.g. O···H···O where both O atoms are similar).
• The hydrogen bond can be unusually strong, with partial covalent character and shorter distances.

These are sometimes referred to as low-barrier hydrogen bonds or short, strong hydrogen bonds and can affect properties like acidity and reaction pathways.

Non-Classical Hydrogen Bonds

Hydrogen bonding can sometimes occur with less typical donors or acceptors:

• C–H groups adjacent to highly electronegative substituents (e.g. in some aromatic or fluorinated compounds) can act as weak donors.
• Acceptors can include atoms like Cl, S, or π-systems (aromatic rings) under certain conditions.

These interactions are usually weaker than classical N–H, O–H, or F–H hydrogen bonds but can still influence structure and reactivity when many are present.

Summary of Key Characteristics

• Requires a hydrogen covalently bound to a strongly electronegative donor atom (N, O, F) and a lone pair on an acceptor atom (often also N, O, or F).
• Stronger and more directional than most other intermolecular forces, but weaker than covalent bonds.
• Leads to characteristic structural motifs in liquids and solids (networks, chains, sheets).
• Strongly influences physical properties such as boiling point, melting point, viscosity, and solubility.
• Essential for the structures and functions of many biological macromolecules, including proteins, nucleic acids, and polysaccharides.
• Can be intermolecular or intramolecular, ideal or distorted, classical or non-classical, but in all cases arises from a specific interaction involving hydrogen and electronegative atoms.