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London Forces or van der Waals Forces    Dipole-Dipole Attraction    Hydrogen Bonding

 

van der Waals Forces and London Dispersion Forces

 
 
A momentary dipole is all it takes to cause weak attraction in nonpolar molecules. Cl2, Br2 and I2 are all nonpolar so any attractions that occur (liquid and solid phases) are caused by van der Waals forces.

 

Simplified Explanation

van der Waals Forces are very weak forces of attraction between molecules resulting from: 

  • London Forces (Fritz London 1930),  are weak IMF that arise from momentary dipoles occurring due to uneven electron distributions in neighboring molecules as they approach one another

  • the weak residual attraction of the nuclei in one molecule for the electrons in a neighboring molecule. 

****This weak intermolecular attraction is why nonpolar gases in our atmosphere can form liquids under extreme pressure and low temperatures.

The more electrons that are present in the molecule, the stronger the dispersion forces will be. 

Dispersion forces are the only type of intermolecular force operating between non-polar molecules, for example, dispersion forces operate between 

  • hydrogen (H2) molecules

  • chlorine (Cl2) molecules

  • carbon dioxide (CO2) molecules

  • dinitrogen tetroxide (N2O4) molecules

  • methane (CH4) molecules

 

A More Complex Explanation

The attractive forces that exist between molecules are responsible for many of the bulk physical properties exhibited by substances. Some compounds are gases, some are liquids, and others are solids. The melting and boiling points of pure substances reflect these intermolecular forces, and are commonly used for identification. Of these two, the boiling point is considered the most representative measure of general intermolecular attractions. Thus, a melting point reflects the thermal energy needed to convert the highly ordered array of molecules in a crystal lattice to the randomness of a liquid. The distance between molecules in a crystal lattice is small and regular, with intermolecular forces serving to constrain the motion of the molecules more severely than in the liquid state. Molecular size is important, but shape is also critical, since individual molecules need to fit together cooperatively for the attractive lattice forces to be large. Spherically shaped molecules generally have relatively high melting points, which in some cases approach the boiling point, reflecting the fact that spheres can pack together more closely than other shapes. This structure or shape sensitivity is one of the reasons that melting points are widely used to identify specific compounds.
Boiling points, on the other hand, essentially reflect the kinetic energy needed to release a molecule from the cooperative attractions of the liquid state so that it becomes an unencumbered and relative independent gaseous state species. All atoms and molecules have a weak attraction for one another, known as van der Waals attraction. This attractive force has its origin in the electrostatic attraction of the electrons of one molecule or atom for the nuclei of another, and has been called London dispersion force.


The following animation illustrates how close approach of two neon atoms may perturb their electron distributions in a manner that induces dipole attraction. The induced dipoles are transient, but are sufficient to permit liquefaction of neon at low temperature and high pressure.

In general, larger molecules have higher boiling points than smaller molecules of the same kind, indicating that dispersion forces increase with mass, number of electrons, number of atoms or some combination thereof. The following table lists the boiling points of an assortment of elements and covalent compounds composed of molecules lacking a permanent dipole. The number of electrons in each species is noted in the first column, and the mass of each is given as a superscript number preceding the formula.

 

# Electrons

Molecules & Boiling Points C

10

20Ne  –246   ;   16CH4  –162

18

40Ar  –186   ;   32SiH4  –112   ;   30C2H6  –89   ;   38F2  –187

34-44

84Kr  –152   ;   58C4H10  –0.5   ;   72(CH3)4C  10   ;   71Cl2  –35   ;   88CF4  –130

66-76

114[(CH3)3C]2  106   ;   126(CH2)9  174   ;   160Br2  59   ;   154CCl4  77   ;   138C2F6  –78

Two ten electron molecules are shown in the first row. Neon is heavier than methane, but it boils 84 lower. Methane is composed of five atoms, and the additional nuclei may provide greater opportunity for induced dipole formation as other molecules approach. The ease with which the electrons of a molecule, atom or ion are displaced by a neighboring charge is called polarizability, so we may conclude that methane is more polarizable than neon. In the second row, four eighteen electron molecules are listed. Most of their boiling points are higher than the ten electron compounds neon and methane, but fluorine is an exception, boiling 25 below methane. The remaining examples in the table conform to the correlation of boiling point with total electrons and number of nuclei, but fluorine containing molecules remain an exception.
The anomalous behavior of fluorine may be attributed to its very high electronegativity. The fluorine nucleus exerts such a strong attraction for its electrons that they are much less polarizable than the electrons of most other atoms.

Of course, boiling point relationships may be dominated by even stronger attractive forces, such as those involving electrostatic attraction between oppositely charged ionic species, and between the partial charge separations of molecular dipoles. Molecules having a permanent dipole moment should therefore have higher boiling points than equivalent nonpolar compounds, as illustrated by the data in the following table.

 

# Electrons

Molecules & Boiling Points C

14-18

30C2H6  –89   ;   28H2C=CH2  –104   ;   26HC≡CH  –84  ;   30H2C=O  –21  ;   27HC≡N  26  ;   34CH3-F  –78

22-26

42CH3CH=CH2  –48  ;   40CH3C≡CH  –23  ;   44CH3CH=O  21  ;   41CH3C≡N  81  ;   46(CH3)2O  –24  ;   50.5CH3-Cl  –24  ;   52CH2F2  –52

32-44

58(CH3)3CH  –12   ;   56(CH3)2C=CH2  –7   ;   58(CH3)2C=O  56   ;   59(CH3)3N  3   ;   95CH3-Br  45  ;   85CH2Cl2  40  ;   70CHF3  –84

In the first row of compounds, ethane, ethene and ethyne have no molecular dipole, and serve as useful references for single, double and triple bonded derivatives that do. Formaldehyde and hydrogen cyanide clearly show the enhanced intermolecular attraction resulting from a permanent dipole. Methyl fluoride is anomalous, as are most organofluorine compounds. In the second and third rows, all the compounds have permanent dipoles, but those associated with the hydrocarbons (first two compounds in each case) are very small. Large molecular dipoles come chiefly from bonds to high-electronegative atoms (relative to carbon and hydrogen), especially if they are double or triple bonds. Thus, aldehydes, ketones and nitriles tend to be higher boiling than equivalently sized hydrocarbons and alkyl halides. The atypical behavior of fluorine compounds is unexpected in view of the large electronegativity difference between carbon and fluorine.

from-www.cem.msu.edu/~reusch/ VirtualText/physprop.htm

 

Van der Waals equation

Johannes Van der Waals was interested in the kinetic theory of gases and fluids, and his primary work was to develop an equation which applied to real gases, unlike that of Robert Boyle which assumes that there are no attractive forces between molecule and that molecules have zero volume.
In reality, molecules have a small volume and attractive forces exist between them. Van der Waals introduced these properties into the theory by means of two constants, which were specific to each gas:

(P + a/v2)(v-b) = RT  , where a and b are constant for a particular gas.

Van der Waals was awarded a Nobel prize in 1910 for his work on the equation of state of gases and liquids.

The weak, electrostatic attractions between atoms were named Van der Waals forces in his honor.

from-http://www.chemsoc.org/timeline/pages/1920.html

London Forces or van der Waals Forces    Dipole-Dipole Attraction    Hydrogen Bonding

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