Intermolecular Forces – Definition, Types, Equations, Examples
Characteristics of chemical systems are observable when they represent the bulk properties of matter. For example, an individual molecule does not boil, while a bulk boils. Collections of water molecules have wetting properties while individual molecules do not. Water, just like all matter, can exist in different states. It exists as ice in the solid state, it exists as a liquid, and in the gaseous state as vapor or steam. The physical properties of all three states of matter are very different. However, even if physical properties are changing, the chemical properties do not change. But we must note that the rate of chemical reactions somewhat depends on the physical state.
Therefore, to understand the different states of matter and their structures, we need to understand what are intermolecular forces, molecular interactions, their natures, the effect of thermal energy, and the motion of particles. The balance of all these parameters defines the physical state of a substance.
The forces of attraction and repulsion between interacting atoms and molecules are called intermolecular forces.
These intermolecular forces are responsible for most of the chemical and physical properties of matter. For example, the greater the intermolecular forces, the higher is the boiling point. We can safely conclude that the boiling point of an element is directly proportional to the strength of its intermolecular forces.
Van der Waals forces are attractive intermolecular forces. They vary to a large extent in magnitude and are categorized accordingly. They include dispersion forces or London forces, dipole-induced dipole forces, and dipole-dipole forces. Hydrogen Bonding is dipole-dipole interaction but considered separately because only a few elements can participate in hydrogen bond formation.
Note: Ion-dipole forces are also intermolecular forces, but are not considered under van der Waals forces. We can use the boiling point of an element as the parameter to categorize the intermolecular forces.
Dispersion Forces or London Forces
Dispersion Forces or London Forces are those which arise due to the movement of electrons, creating a temporary positive and negative charge. It takes place mostly in atoms and non-polar molecules that are electrically symmetrical and have no dipole moment.
As we can see in the image above, we first consider two atoms A and B that have no dipole moment. Next, we can see Atom A with an instantaneous dipole with more electron density on the right-hand side, while Atom B is with an induced dipole. In the final part, we can observe how Atom A has more electron density on the left-hand side. Atom B is with an induced dipole.
London forces operate for a short distance (~500 pm) and their magnitude depends largely on the polarisability of the particle. It is the weakest force out of all the intermolecular forces.
Polar Molecules are those molecules that have a net dipole because of opposing charges on either end. They have dipole-dipole interactions as attractive forces. Their permanent dipoles due to varying electronegativities of the atoms are associated with a covalent bond. The partially positive of the polar molecule attracts the partially negative part of another molecule.
For example, HCl molecules have dipole-dipole interactions.
Hydrogen is less electronegative as compared to Chlorine. Thus, Chlorine acquires a partially negative charge while Hydrogen gets a partially positive charge. Thus, dipole-dipole interaction takes place between the HCl molecules.
As mentioned before, this is a special case of dipole-dipole interaction. This is found in mostly molecules in which highly polar O-H, H-F, or N-H bonds are present.
Hydrogen Bonding is mostly regarded to be limited to Nitrogen, Oxygen, and Fluorine, but in some cases, species such as Chlorine also participate in Hydrogen bonding.
These are just like dipole-dipole interactions, with the only exception being that they rise between polar molecules and ions. The strength of the ion-dipole interactions depend on the following:
- The size of the polar molecule.
- The charge and size of the ion.
- The magnitude of the dipole moment.
For example, when NaCl is mixed with water, the H2O molecules (polar) are attracted to the Na (sodium) and Cl (Chlorine) ions in the beaker.
Ion Induced Dipole Interactions
In ion-induced dipole interactions, an ion is polarized by a non-polar molecule. The non-polar molecules behave as induced dipoles as they obtain a charge.
This interaction between the ion and the induced dipole is ion-induced dipole interaction.
Dipole Induced Dipole Interactions
The dipole-induced dipole interactions are similar to ion-induced interactions, with the exception being that the non-polar are converted to induced dipoles due to the presence of a polar molecule.
These forces are between the polar molecules that have a permanent dipole and the molecules lacking permanent dipole.
Question 1: What is the difference between Intermolecular Forces and Thermal Interactions?
The difference between Intermolecular Forces and Thermal Interactions are as follows:
The forces of attraction and repulsion between interacting atoms and molecules are called intermolecular forces. The total measure of the sum of the kinetic energy of all the atoms and molecules is called thermal energy. These forces are due to the dipole of one or both the given molecules. These forces are due to the motion of particles. Greater intermolecular forces lead to the substance being in the solid state. Greater thermal interactions will not allow the substance to remain in the solid state. These forces hold particles together. These forces keep particles apart. They do not have any effect due to the temperature, but the boiling point of the substance is directly proportional to the strength of the forces. Thermal energy is directly dependent/proportional to temperature. Intermolecular forces are weak in the gaseous phase and strongest in solid state. Thermal forces are weak in solid state and high in the gaseous phase. The volume of the given matter is less when the intermolecular forces are high. The volume of the given matter is more when the thermal interactions are high. Gases can’t be liquefied on compression only due to the strong intermolecular forces. Gases can be easily liquefied by reducing the thermal energy by lowering the temperature.
Question 2: What is Hydrogen Bonding?
Hydrogen Bonding is a special unique case of dipole-dipole interaction. Dipole-dipole interactions are the attractive forces on polar molecules. Hydrogen Bonding is found in mostly molecules in which highly polar O-H, H-F, or N-H bonds are present. It is mostly regarded to be limited to Nitrogen, Oxygen and Fluorine, but in some cases, species such as Chlorine also participate in Hydrogen bonding.
Question 3: Explain the interaction energy of Dispersion forces and Dipole-Dipole Interactions.
Dispersion Forces or London Forces are those which arise due to the movement of electrons, creating a temporary positive and negative charge. London forces operate for a short distance (~500 pm) and their magnitude depends largely on the polarisability of the particle. The interaction energy of the dispersion forces is inversely proportional to the sixth power of the distance between two particles.
Interaction Energy α 1/x6, where x is the distance between two given particles.
Dipole-dipole interactions are the attractive forces on polar molecules. They are between polar molecules that are mostly of two types. The first type is stationary, while the other is rotating.
The interaction energy of dipole-dipole interactions is inversely proportional to the third power of the distance between two particles in the case of stationary polar molecules.
Interaction Energy α 1/x3, where x is the distance between two given particles.
The interaction energy of dipole-dipole forces varies inversely to the sixth power of the distance between two particles in the case of rotating polar molecules.
Interaction Energy ∝ 1/x6, where x is the distance between two given particles.
Question 4: Why does ice have a lower density than water?
Hydrogen Bonding affects the physical properties of compounds. Ice has hydrogen bonding as intermolecular forces. Thus, it has a lower density than water because of hydrogen bonding and cage-like structure of ice.
Ice has hexagonal three-dimensional crystal structure (as per X-ray crystallographic data). This hexagonal crystal structure is formed due to intermolecular hydrogen bonding.
When ice melts, most of the hydrogen bonds break and some of the empty spaces are occupied by water molecules. Liquid water molecules are thus more closely packed together than molecules in ice. Thus, ice has a lower density than water.
Question 5: Water has maximum density at 4o Celsius. Why?
Water molecules in ice exist in a crystal lattice with a lot of empty space.
When ice melts into liquid water, the density of the water increases as the structure starts to break and collapse. As we increase the temperature, the molecules start moving faster and get further apart. As the temperature increases, the density decreases. At temperatures nearing 0oC, water still has several ice-like clusters. As the temperature of warm water decreases, the water molecules slow down and the density increases. At 4oC, the clusters start forming. The molecules are still slowing down and coming closer together, but the formation of clusters makes the molecules be further apart. Cluster formation is the bigger effect, so the density starts to decrease. Thus, the density of water is a maximum at 4oC.
Question 6: What is the strength of hydrogen bonds dependent on?
Strength of Hydrogen Bonds depends on the coulombic interaction between the lone-pair electrons of one electronegative atom of a molecule and the hydrogen atom of another molecule.
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