Attractive forces London dispersion

Polarizability is a measure of the ease with which the electrons of a molecule are distorted. It is the basis for evaluating the nonspecific attraction forces (London dispersion forces) that arise when two molecules approach each other. Each molecule distorts the electron cloud of the other and thereby induces an instantaneous dipole. The induced dipoles then attract each other. Dispersion forces are weak and are most important for the nonpolar solvents where other solvation forces are absent. They do, nevertheless, become stronger the larger the electron cloud, and they may also become important for some of the higher-molecular-weight polar solvents. Large solute particles such as iodide ion interact by this mechanism more strongly than do small ones such as fluoride ion. Furthermore, solvent polarizability may influence rates of certain types of reactions because transition states may be of different polarizability from reactants and so be differently solvated. 

Distinguish between dipole-dipole attraction forces and dispersion (London) forces. Is one type of these intermolecular forces stronger than the other Explain. 

All adsorption processes result from the attraction between like and unlike molecules. For the ethanol-water example given above, the attraction between water molecules is greater than between molecules of water and ethanol As a consequence, there is a tendency for the ethanol molecules to be expelled from the bulk of the solution and to concentrate at die surface. This tendency increases with the hydrocarhon chain-length of the alcohol. Gas molecules adsorb on a solid surface because of die attraction between unlike molecules. The attraction between like and unlike molecules arises from a variety of intermolecular forces. London dispersion forces exist in all types of matter and always act as an attractive force between adjacent atoms and molecules, no matter how dissimilar they are. Many oilier attractive forces depend upon die specific chemical nature of the neighboring molecules. These include dipole interactions, the hydrogen bond and the metallic bond. 

Weak attractive forces between nonbonded atoms are called van der Waals attractive forces, London forces, or dispersion forces, and are of great importance in determining the properties of liquids. They also can be expected to play a role in determining conformational equilibria whenever the distances between the atoms in the conformations correspond to the so-called van der Waals minima. 

London dispersion forces. London dispersion forces result from the attraction of correlated temporary dipoles. 

In contrast to dipole-dipole forces, London Dispersion interactions are much weaker in nature since they involve nonpolar molecules that do not possess permanent dipole moments. The only modes for molecular attraction are through polarization of electrons, which leads to the creation of small dipole-dipole interactions and mutual attractive forces. Since electron polarization occurs much more readily for electrons farther from the nucleus, this effect is more pronounced for molecules that are larger with a greater number of electrons, especially positioned on atoms with a high atomic number, consisting of more diffuse orbitals. These induced dipole forces are responsible for the liquefaction of gases such as He and Ar at low temperatures and pressures. The relative strength of London Dispersion forces is described by Eq. 3 ... 

All ions, atoms, and molecules are attracted by London dispersion forces. 

This chapter aims to give guidelines on how to use adsorption methods for the characterization of the surface area and pore size of heterogeneous catalysts. The information derived from these measurements can range from the total and available specific surface area to the pore sizes and the strength of sorption in micropores. Note that this spans information from a macroscopic description of the pore volume/specific surface area to a detailed microscopic assessment of the environment capable of sorbing molecules. In this chapter we will, however, be confined to the interaction between sorbed molecules and solid sorbents that are based on unspecific attractive and repulsive forces (van der Waals forces, London dispersion forces). 

Surface tension arises due to short range intermolecular forces. The most important ones are van der Waals forces, London dispersion forces, hydrogen and metallic bondings [1]. The contributions from the individual forces are assumed independent, and the effective surface tension are calculated as the linear sum of the individual force contributions. The different molecular attraction forces at the two sides of the interfaces induce a resulting attraction force at the interface. Imagine that the molecules at an interface between two fluids exist in a state different from that of the molecules in the interior of the fluid. The phase k molecules are (on the average) surrounded by phase k molecules on only one side within the interface, whereas the interior... 

The strength of London dispersion forces increases with the molecular masses of the molecules. The larger the molecular mass, the stronger the London attractive force. London forces exist in all substances. 

Three types of intermolecular attractive forces are known to exist between neutral molecules dipole-dipole forces, London dispersion forces, and hydrogenbonding forces. These forces are also called van der Waals forces after Johannes van der Waals, who developed the equation for predicting the deviation of gases from ideal behavior. (Section 10.9) Another kind of attractive force, the ion-dipole force, is important in solutions. All four forces are electrostatic in nature, involving attractions between positive and negative species. All tend to be less than 15% as strong as covalent or ionic bonds. 

As in a gas, particles in a liquid are in constant motion. However, the particles in a liquid are closer together than the particles in a gas are. Therefore, the attractive forces between particles in a liquid are more effective than those between particles in a gas. This attraction between liquid particles is caused by intermolecular forces, such as dipole-dipole forces, London dispersion forces, and hydrogen bonding. 

Van der Waals force Also called intermolecular forces, secondary valence forces, dispersion force, London dispersion force, or van der Waals attraction. It is an attractive force between two atoms or non-polar molecules, which arise because a fluctuating dipole moment in one molecule induces a dipole moment in the other, and the two dipole moments then interact. They are somewhat weaker than hydrogen bonds and far weaker than inter-atomic valences. Information regarding their numerical values is mostly semi-empirical, derived with the aid of theory from an analysis of physical and chemical data. 

The stability of nanoparticle suspensions is an important factor, as it determines the efficacy of the nanocapsules used in drug delivery applications. The stability is determined by the balance between the attractive van der Waals force and repulsive electrostatic force caused by the double layer of the oppositely charged ions [64]. Attraction and aggregation of nanoparticles are due to induced dipole—dipole forces (London dispersion forces). Induced dipole force is a part of the van der Waals forces. These forces result from multipoles formed in molecules, caused by quanmm induced instantaneous polarization. The formation of instantaneous dipoles occurs, because the electrons in adjacent molecules redistribute due to their correlated movements. 

Induced dipole/mduced dipole attraction (Section 2 17) Force of attraction resulting from a mutual and complemen tary polanzation of one molecule by another Also referred to as London forces or dispersion forces Inductive effect (Section 1 15) An electronic effect transmit ted by successive polanzation of the cr bonds within a mol ecule or an ion... 

Attractive and Repulsive Forces. The force that causes small particles to stick together after colliding is van der Waals attraction. There are three van der Waals forces (/) Keesom-van der Waals, due to dipole—dipole interactions that have higher probabiUty of attractive orientations than nonattractive (2) Debye-van der Waals, due to dipole-induced dipole interactions (ie, uneven charge distribution is induced in a nonpolar material) and (J) London dispersion forces, which occur between two nonpolar substances. 

Boiling Point. When describing the effect of alkane structure on boiling point in Section 2.17, we pointed out that van der Waals attractive forces between neutral molecules are of three types. The first two involve induced dipoles and are often refened to as dispersion forces, or London forces. 

The dispersion forces (London forces) which are always attractive. 

The ever present dispersion force (London attraction). 

In nature, the halogens exist as nonpolar diatomic molecules. London dispersion forces are the only forces of attraction acting between the molecules. These forces increase with increasing molecular size. 

The discussion thus far has focused on the forces between an array of atoms connected together through covalent bonds and their angles. Important interactions occur between atoms not directly bonded together. The theoretical explanation for attractive and repulsive forces for nonbonded atoms i and j is based on electron distributions. The motion of electrons about a nucleus creates instantaneous dipoles. The instantaneous dipoles on atom i induce dipoles of opposite polarity on atom j. The interactions between the instantaneous dipole on atom i with the induced instantaneous dipole on atom j of the two electron clouds of nonbonded atoms are responsible for attractive interactions. The attractive interactions are know as London Dispersion forces,70 which are related to r 6, where r is the distance between nonbonded atoms i and j. As the two electron clouds of nonbonded atoms i and j approach one another, they start to overlap. There is a point where electron-electron and nuclear-nuclear repulsion of like charges overwhelms the London Dispersion forces.33 The repulsive... 

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