Why can water molecules form 4 hydrogen bonds?

Water (H2O) is a unique molecule with special properties that allow it to form up to four hydrogen bonds with neighboring water molecules. This ability to form multiple hydrogen bonds gives rise to many of water’s distinctive physical and chemical characteristics.

What is a Hydrogen Bond?

A hydrogen bond is an electrostatic attraction between a hydrogen atom on one molecule and an electronegative atom (usually oxygen or nitrogen) on another molecule. It is not as strong as a covalent bond, but stronger than van der Waals forces.

For a hydrogen bond to occur, three criteria must be met:

  • A hydrogen atom must be covalently bonded to an electronegative atom (usually O, N, or F). This creates a polar covalent bond with a slightly positive charge on the H atom.
  • The hydrogen atom must be close to another electronegative atom with a lone pair of electrons.
  • The electronegative atoms must be small enough to facilitate close contact between the molecules.

When these criteria are met, the positive charge on the hydrogen interacts with the lone pair electrons on the electronegative atom, creating a strong dipole-dipole interaction or hydrogen bond.

Hydrogen Bonding in Water

In the water molecule, the oxygen atom has six valence electrons. Two of these electrons participate in covalent bonds with the hydrogen atoms, leaving four unbonded lone pairs on the oxygen. This allows water to act as both a hydrogen bond donor and acceptor.

The O-H bonds in water are polar, meaning the oxygen exerts a partial negative charge and the hydrogens carry a partial positive charge. This charge difference allows the hydrogen atoms on one water molecule to electrostatically interact with the lone pairs on the oxygen atom of another water molecule.

Hydrogen bonding in water results in several important properties:

  • High boiling point – Hydrogen bonds need a lot of energy to break, so water has a boiling point of 100°C, which is higher than expected for such a small molecule.
  • High heat capacity – More energy is needed to disrupt the hydrogen bonds during heating and cooling.
  • Cohesive behavior – Hydrogen bonds hold water molecules together, giving water surface tension.
  • Solvent capabilities – The polar nature of water allows it to dissolve polar substances by forming hydration shells via hydrogen bonding.

How Many Hydrogen Bonds Can a Water Molecule Form?

An individual water molecule can participate in a maximum of four hydrogen bonds – two as a hydrogen bond donor and two as a hydrogen bond acceptor. This is due to water’s tetrahedral molecular geometry.

The two O-H covalent bonds form an angle of about 104.5°. In a tetrahedral arrangement, a central atom is bonded to four outer atoms located at the corners of a tetrahedron. This allows the two hydrogen atoms and two lone pairs on the oxygen atom to orient themselves in a tetrahedral geometry.

This tetrahedral arrangement allows each hydrogen atom and each lone pair to participate in hydrogen bonding with neighboring water molecules. The two hydrogen atoms can act as hydrogen bond donors, while the two lone pairs can act as hydrogen bond acceptors.

2 Donor Bonds

The partially positive hydrogen atoms on one water molecule are electrostatically attracted to the lone pairs on oxygen atoms of adjacent water molecules. This allows each H atom to form a hydrogen bond to an adjacent water molecule, for a total of two donor hydrogen bonds.

2 Acceptor Bonds

The two lone pairs on the oxygen atom of a central water molecule can form hydrogen bonds with the partially positive H atoms on adjacent water molecules. This allows the central water to form two hydrogen bonds as an acceptor, one from each lone pair.

Network of Hydrogen Bonds

Within liquid water and solid ice, the tetrahedral arrangement and hydrogen bonding capacity of water molecules give rise to an extensive, intermolecular network of hydrogen bonds that extend in three dimensions. Each water molecule accepts two and donates two hydrogen bonds.

This organized network of hydrogen bonds is responsible for many of water’s emergent properties like high boiling point, strong surface tension, and the unique density profile of ice.

In the liquid state, water molecules are constantly moving and exchanging hydrogen bond partners. This highly dynamic and rapidly changing hydrogen bond network allows water to flow as a liquid while retaining a high degree of short-range order and cohesion between molecules.

Covalent Bonding in Water

In addition to intermolecular hydrogen bonding, the chemical properties of water originate from the covalent bonds within an individual H2O molecule. There are two strong O-H covalent bonds between the oxygen atom and each hydrogen atom.

Some key properties of the O-H covalent bonds are:

  • Very polar – oxygen exerts a strong electronegative pull on the shared electrons
  • Short bond length – 95 pm
  • High bond energy – about 460 kJ/mol

The polarity of the O-H bonds allows water to act as both a hydrogen bond donor and acceptor. The partial charges created by these polar covalent bonds facilitate the intermolecular hydrogen bonding interactions with neighboring water molecules.

Resonance Structures

The covalent bonding in water can be represented by two resonance structures:

Resonance structures depict the dynamic movement of electrons within a molecule. For water, the actual structure is a hybrid of these two resonance forms. The position of the bonding electrons fluctuates between the three atoms.

These resonance structures illustrate how the electrons are simultaneously shared between the O-H covalent bonds and the oxygen lone pairs. This electron movement facilitates hydrogen bond formation and breakdown, assisting in water’s network flexibility and solvent capabilities.

Orbital Hybridization

The electronic configuration of the oxygen atom is 1s2 2s2 2p4. To form two covalent bonds and two lone pairs, the oxygen atom undergoes sp3 hybridization.

In sp3 hybridization, one 2s orbital combines with three 2p orbitals to produce four new sp3 hybrid orbitals oriented in a tetrahedral geometry. These four orbitals contain one lone pair each and form the sigma bonds to the hydrogen atoms at an angle of 104.5°.

This hybridization and geometry allows water to achieve tetrahedral symmetry and maximize the potential for hydrogen bond formation and networking.

Quantum Mechanical Model

Modern quantum mechanical models can accurately describe the electron distribution and bonding patterns in the water molecule.

Quantum mechanics calculates electron probability distributions rather than fixed electron positions. It predicts the following electron arrangement for water:

  • Two covalent bonds form between the 1s orbital on each hydrogen and a sp3 orbital on oxygen.
  • The four sp3 hybrid orbitals are filled with two lone pairs.
  • Due to oxygen’s higher electronegativity, the bonding electron density is concentrated around the oxygen nucleus.
  • The electron pairs occupy tetrahedrally-arranged sp3 orbitals to minimize repulsion.

This model corroborates the tetrahedral structure and demonstrates how water’s orbital configuration enables hydrogen bonding capabilities.

Molecular Orbital Theory

Molecular orbital (MO) theory provides another model of bonding in water. Instead of localized electrons, molecular orbital theory describes bonding in terms of the combination and interactions between the atomic orbitals of the bonded atoms.

In water, MO theory predicts bond formation from the overlap of the H 1s orbitals and the O 2p orbitals. The key molecular orbitals are:

  • Two covalent bonding orbitals
  • Two covalent antibonding orbitals
  • Four lone pair orbitals

This delocalized model also supports the tetrahedral arrangement and hydrogen bond capabilities derived from a four-lobed electron density distribution.

VSEPR Theory

Valence shell electron pair repulsion (VSEPR) theory is a simple model that predicts molecular geometry based on minimizing repulsions between electron pairs. It accurately predicts water’s tetrahedral shape.

In water, VSEPR theory considers:

  • Two bonding electron pairs in the O-H covalent bonds
  • Two lone electron pairs on the oxygen atom

To minimize repulsions, the four electron pairs orient themselves tetrahedrally. This angle maximizes distance between the electron pairs while ensuring space for four neighboring water molecules.

By incorporating lone pair repulsions, VSEPR theory arrived at the tetrahedral model that facilitates hydrogen bonding over 100 years before modern computational chemistry.

Water Phase Diagrams

The hydrogen bonding abilities of water molecules give rise to unique phase change behavior that can be depicted on phase diagrams.

Unlike most substances, water expands upon freezing. This results in the density anomaly where ice is less dense than liquid water.

The hydrogen bond network forms an open, crystalline lattice in ice that is less dense than liquid water. This density behavior is the opposite of most materials, which contract and become more dense upon freezing.

At high pressures, the density anomaly disappears as the hydrogen bond network becomes more compact and coordinated in various high-pressure ice phases.

Understanding these phase changes and density anomalies relies heavily on water’s capacity to form an interconnected matrix of hydrogen bonds between molecules.

Triple Point

The distinct hydrogen bonding abilities of water also give rise to the triple point – a single temperature and pressure at which water co-exists in equilibrium as gas, liquid, and solid phases simultaneously.

For water, the triple point occurs at 0.01°C and a partial vapor pressure of 0.006129 bar. At this point, adding or removing heat will transition between solid, liquid, and gaseous states while maintaining equilibrium.

The well-defined triple point is a result of water’s hydrogen bonding networking capabilities in each distinct phase.

Water as a Solvent

Water’s ability to form hydrogen bonds makes it an excellent solvent, able to dissolve a wide range of polar substances.

When a solute dissolves in water, its polar molecules interact with water via hydrogen bonds. This allows the solute to mix homogeneously at a molecular level with the solvent.

Some examples of water’s solvent capabilities:

Solute Solute Polarity Solvation Mechanism
Sodium chloride Ionic Ion-dipole interactions
Ethanol Polar covalent Hydrogen bonding
Sucrose Polar covalent Hydrogen bonding

Water itself also interacts readily with ions and other polar molecules via hydrogen bonding. This enables water to dissolve a wide range of substances while resisting separation into distinct phases.


Pure water undergoes a small degree of self-ionization through two simultaneous reactions:

H2O ⇌ H+ + OH
H2O + H2O ⇌ H3O+ + OH

About two out of every billion water molecules are ionized at any given time. The hydrogen and hydroxide ions produced can facilitate further chemical reactions.

Self-ionization occurs when protons are transferred between water molecules, enabled by hydrogen bonding interactions. This generates an equilibrium mixture of hydrated protons and hydroxide ions.

The ion concentrations are described by the water dissociation constant Kw = [H+][OH-] = 10-14 at 25°C. This equilibrium facilitates many acid-base reactions in aqueous solutions.

Micellular Solubilization

Water’s hydrogen bonding capabilities allow it to interact with amphiphilic molecules and promote the formation of supramolecular micelle structures in aqueous solutions.

Micelles consist of aggregation of amphiphilic molecules such as surfactants and lipids. The hydrophobic tails orient towards the micelle center, while the hydrophilic heads remain in contact with surrounding water molecules.

Water forms hydrogen bonds with the hydrophilic micelle heads, enabling dissolution. The non-polar hydrophobic cores provide regions for solubilizing non-polar solutes in an aqueous environment.

Detergents utilize this ability of water to form microscopic emulsions and solubilize both polar and non-polar substances.

Hydrophobic Effects

The hydrogen bonding between water molecules produces powerful hydrophobic effects – the tendency of non-polar substances to aggregate rather than dissolve in water.

Non-polar molecules cannot form hydrogen bonds with water. This distorts the highly organized hydrogen bond network, forcing the non-polar substances to cluster together and minimize contact with water.

Key hydrophobic effects in water include:

  • Insolubility of non-polar gases like methane
  • Separation of oil and water into distinct phases
  • Protein folding – non-polar residues cluster inside away from water
  • Self-assembly of lipid bilayer membranes

These behaviors minimize disruption to water’s hydrogen bond network by excluding non-polar molecules.


Water owes many of its characteristic properties and solvent capabilities to its ability to form a maximum of four hydrogen bonds per molecule. Two covalent O-H bonds allow it to donate two hydrogen bonds, while two lone pairs permit it to accept two H-bonds.

This tetrahedral arrangement maximizes water’s capacity for intermolecular hydrogen bonding and facilitates the formation of an extensive 3D network of hydrogen bonds in liquid water and solid ice.

The presence of this hydrogen bond network explains water’s unusually high boiling point, solvent properties, phase behaviors, and role in supporting life. Understanding the hydrogen bonding capabilities of water is key to unraveling its many anomalies and distinctive characteristics.

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