Consider two polar molecules water H₂O (molecular mass 18 and dipole moment 1.8 D) and nitrosyl fluoride ONF (molecular mass 49 and dipole moment 1.8 D). As their dipole moments are the same, their boiling points should be comparable. But boiling point of water is 100° C while that of nitrosyl fluoride is – 56° C. Thus there is a large difference between their boiling points. This difference is due to a special type of bonding between water molecules called hydrogen bonding.

The large difference in electronegativities of the pairs of bonded atoms N-H, O-H, F-H and Cl-H establish highly polar covalent bonds in which hydrogen acquires exceptionally large positive charge and another electronegative atom acquires exceptionally large negative charge on it. Due to this, there is an attraction between the positively charged hydrogen atom of one molecule and negatively charged atom of another molecule. This interaction is called hydrogen bonding. This concept was introduced by Latimer and Rodebush in 1920. A hydrogen bond is an electromagnetic attraction created between a partially positively charged hydrogen atom attached to a highly electronegative atom and another nearby electronegative atom.

This interaction is represented by a dotted line. The energy of the hydrogen bond varies between 10 to 100 kJ mol^-1. This energy is significant. Hence it is a very important factor which influences the bulk properties of a substance.

Variation of Boiling Points of Hydrides:

Following graph shows a variation of boiling points of hydrides of groups 4A, 5A, 6A, 7A.

From the graph, we can see that the boiling point of ammonia (hydride of nitrogen), water (hydride of oxygen) and hydrogen fluoride (hydride of fluorine) are exceptionally higher in their group. It is due to strong molecular hydrogen bonding.

We can see that methane (hydride of carbon) has a low boiling point due to the absence of hydrogen bonding. In case of inert gases, the trend is increasing in boiling points with the increase in atomic mass. It is due to dispersion forces increase with the mass of the molecule.

Nature of Hydrogen Bonding:

• A hydrogen bond is a type of dipole-dipole interaction; it is not a true chemical bond it is a mere electrostatic attraction.
• These attractions can occur between molecules (intermolecularly) or within different parts of a single molecule (intramolecularly).
• Intramolecular hydrogen bonds are those which occur within one single molecule. This occurs when two functional groups of a molecule can form hydrogen bonds with each other. In order for this to happen, both a hydrogen donor an acceptor must be present within one molecule, and they must be within close proximity of each other in the molecule. For example, intramolecular hydrogen bonding occurs in ethylene glycol (C₂H₄(OH)₂) between its two hydroxyl groups due to the molecular geometry.
• Intermolecular hydrogen bonds occur between separate molecules in a substance. They can occur between any number of like or unlike molecules as long as hydrogen donors and acceptors are present an in positions in which they can interact. For example, intermolecular hydrogen bonds can occur between NH₃ molecules alone, between H₂O molecules alone, or between NH₃ and H₂O molecules.
• Hydrogen bond never involves more than two atoms.
• With the increase in electronegativity and decrease in the size of the atom to which hydrogen is covalently attached, the strength of the bonding increases.
• Hydrogen bonds are generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds.

Effects of Hydrogen Bonding:

• Dissociation: Due to hydrogen bonding in HF, in aqueous solution, hydrogen fluoride (HF) dissociates and gives the difluoride ion instead of fluoride ion. The molecules of HCl, HBr, HI  do not form hydrogen bonding. This explains the non -existence of compounds like KHCl₂ , KHBr₂, KHI₂.
• Association: The molecules of carboxylic acids exist as dimer because of the hydrogen bonding. The molecular masses of such compounds are found to be double than those calculated from their simple molecular formula.
• High melting and Boiling point, Less Volatility: As extra energy is required to break hydrogen bonds, the compounds having hydrogen bonding show abnormally high melting and boiling points. The unusually high boiling points of hydrogen fluoride (w.r.t other hydrogen halides), water (w.r.t. H₂S, H₂Se and H₂Te), ammonia (w.r.t PH₃), ethanol (w.r.t. ethers) are due to the existence of hydrogen bonding.
• Solubility: Solubility increases with the increase in easiness to form a hydrogen bond. Lower alcohols are soluble in water because of the hydrogen bonding which can take place between water and alcohol molecule.
• Viscosity and surface tension: The substances having hydrogen bonding exist as an associated molecule. Hence their flow becomes comparatively difficult. Thus they have higher viscosity and high surface tension.

Importance of Hydrogen Bonding:

• These bonds occur in inorganic molecules, such as water, and organic molecules, such as DNA and proteins.
• The two complementary strands of DNA are held together by hydrogen bonds between complementary nucleotides (A&T, C&G).
• In water it contributes to its unique properties, including its high boiling point (100 °C) and surface tension.
• Intramolecular hydrogen bonding is partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids. The hydrogen bonds help the proteins and nucleic acids form and maintain specific shapes.

Science > Chemistry > States of Matter > Hydrogen Bonding

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Intermolecular forces are the forces of attraction between neighbouring molecules, situated at a distance much closer in comparison with their molecular diameter. They vary from state to state of matter. These forces exist in all states of matter. Liquid and solid states of matter are regarded as a condensed state. These forces arise due to the interaction between molecules. These forces are collectively called as van der Waal’s forces. The intermolecular forces are much weaker than the intramolecular forces. The melting and boiling points of the substances depend on the magnitude of intermolecular forces. Larger the magnitude of these forces higher is the melting point and boiling point of the substance.

These forces due to the interaction between molecules are further classified as a) Dipole-dipole interaction b) Dipole-induced dipole interaction c) Dispersion forces or London Forces. Besides van der Waal’s forces, there are two more intermolecular attraction forces called as ion-dipole interaction and hydrogen bonding.

Dipole-Dipole Interaction Between Molecules:

This effect was studied by Keesom in 1912, hence these forces are also called Keesom forces and also referred as orientation effect. Compounds having atoms of different electro-negativities act as a dipole (e.g. SO₂, HCl, NO₂, etc.) Dipole-dipole interaction results in a force of attraction between neighbouring molecules having a permanent dipole moment. This interaction is due to the electrostatic force.

Consider polar molecule HCl, a compound of electropositive element hydrogen and electronegative element chlorine. Due to which chlorine pulls shared electron in the bond towards itself acquiring small negative charge (δ^–) and hydrogen acquires an equal positive charge (δ⁺). Thus the HCl molecule becomes an electrical dipole.

Due to these, all molecules align in a way that oppositely charged ends come close to one another as shown.

Other examples of this type of interaction are HBr, H₂S, NH₃.

Characteristics of Dipole-Dipole Interaction:

• Dipole-dipole interaction results in a force of attraction between neighbouring molecules having a permanent dipole moment due to electrostatic force.
• The strength of dipole-dipole interaction depends upon the dipole moment of the interacting molecule.
• Besides dipole-dipole interaction, the polar molecules interact with the London forces. The combined effect of the two increases the total intermolecular forces.
• The dipole-dipole interaction becomes significant when the distance between the dipoles is less than 500 pm.
• In dipole-dipole interactions for mobile dipoles, the interaction energy is inversely proportional to the sixth power of the distance between the two interacting molecules.

Interaction energy ∝ – 1/r⁶

• The negative sign indicates that energy is released when the two dipoles approach each other.
• In dipole-dipole interactions for stationary dipoles (as in solids), the interaction energy is inversely proportional to the cube of the distance between the two interacting molecules.

Interaction energy ∝ – 1/r³

• The negative sign indicates that energy is released when the two dipoles approach each other.
• The energy for dipole-dipole interaction lies between 4.18 J – 12.55 J

Dipole-Induced Dipole:

This interaction was studied by Debye (1920). This interaction is also called an inductive effect or induction effect. These type of forces operate between polar (μ > 1) and nonpolar (μ = 1) molecules.

In such interaction, the permanent dipole of the polar molecules induces dipole on the nonpolar molecule by deforming or polarizing its electronic cloud.

The interaction energy of these forces is inversely proportional to the sixth power of the distance between the two interacting molecules.

The strength of the forces depends on the distance between the molecules and the polarizability of the nonpolar molecule.

Example: Interaction between NH₃ (polar) and C₆H₆ (nonpolar).

The ease with which an atom or non-polar molecule’s electron cloud can be distorted is called its polarizability. In smaller molecules the electron cloud is near to the nucleus hence the electron cloud cannot be distorted easily. Hence smaller molecules are less polarizable than the larger molecules. In a group of a periodic table, the polarizability increases down the group while in the period the polarizability decreases as we move from left to right in a period of a periodic table.

Characteristics of Dipole-Induced Dipole Interaction:

• In such interaction, the permanent dipole of the polar molecules induces dipole on the nonpolar molecule by deforming or polarizing its electronic cloud.
• The interaction energy of these forces is inversely proportional to the sixth power of the distance between the two interacting molecules.
• The strength of the forces depends on the distance between the molecules and the polarizability of the nonpolar molecule.
• Polarizability decides the ease of condensation of gases containing atoms or non-polar molecules.

Dispersion Forces or London Forces:

This force of attraction was first proposed by the German physicist Fritz London, and for this reason force of attraction between two temporary dipoles is known as London forces. Another name for this force is dispersion force.

Atoms and nonpolar molecules are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed. But a dipole may develop momentarily even in such atoms and molecules due to polarizability. Polarizability is an ease with which the arrangement of electrons in the atom or molecule can be disturbed.

Suppose we have two nonpolar atoms ‘A’ and ‘B’ in the close vicinity of each other. It may so happen that momentarily electronic charge distribution in one of the atoms say ‘A’, becomes unsymmetrical i.e., the charge cloud is more on one side than the other.

This results in the development of instantaneous dipole on the atom ‘A’ for a very short time. This instantaneous or transient dipole distorts the electron density of the other atom ‘B’, which is close to it and as a consequence, a dipole is induced in the atom ‘B’. The temporary dipoles of atom ‘A’ and ‘B’ attract each other. Similarly, temporary dipoles are induced in molecules also.

Characteristics of Dispersion Forces:

• In this interaction, the atoms or nonpolar molecules may develop a dipole momentarily due to polarizability.
• These forces are always attractive and interaction energy is inversely proportional to the sixth power of the distance between two interacting particles.  where r is the distance between two particles.
• These forces are important only at short distances (~500 pm) and their magnitude depends on the polarizability of the particle.
• The strength of these forces increases with the increase in molecular mass, molecular size, number of electrons and surface area of the molecule.
• The liquid state of helium and methane is due to the presence of dispersion forces. Other examples are N₂, H₂, CO₂.
• The magnitude of dispersion forces depends on the molecular mass, molecular size, molecular geometry.
• The dispersion forces increase with the increase in the molecular mass and molecular size of an atom. A more spacious arrangement of atoms in molecules results in greater dispersion forces than the compact arrangement.

Ion-Dipole Interactions:

Ion-Dipole Forces are involved in solutions where an ionic compound is dissolved into a polar solvent. Hence the system exhibiting ion-dipole interactions are solutions and not pure substances.

Cations are smaller than anions. The charge on cations is more concentrated. Due to this, the interaction between a cation and the negative end of the polar molecule are stronger than the corresponding interaction between the anion and the positive end of the polar molecule.

Example: Interaction between the ions formed by dissolving NaCl in water and water molecules. ions of NaCl dissociates because of the attraction between the separated ions and oppositely charged poles of water.

Hydration is a process by which each ion is surrounded by a number of water molecules with an oppositely charged pole of solvent water oriented in the direction of the ion concerned.

Ion-dipole and ion-induced dipole forces operate much like dipole-dipole and induced dipole-dipole interactions. However, ion-dipole forces involve ions instead of solely polar molecules. Ion-dipole forces are stronger than dipole interactions because the charge of an ion is much greater than the charge of a dipole.

The strength of ion-dipole interaction depends on the charge and the size of the ion and the magnitude of the dipole moment and the size of the dipole. Ion-dipole bonding is also stronger than hydrogen bonding. An ion-dipole force consists of an ion and a polar molecule aligning so that the positive and negative charges are next to one another, allowing for maximum attraction.  Intermolecular ion-dipole forces are much weaker than covalent or ionic bonds.

Science > Chemistry > States of Matter > Interaction Between Molecules

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