Liquids and Solids AP CHEMISTRY

 Gas can be pictured as a substance whose component
particles are far apart, are in rapid random motion, and exert relatively small forces on each other.

The kinetic molecular model was constructed to account for the ideal behavior that most gases approach at high temperatures and low pressures.

Solids are obviously very different from gases. A gas has low density and high compressibility and completely fills its container. Solids have much greater densities, are compressible only to a very slight extent, and are rigid

a solid maintains its shape irrespective of its container.

These properties indicate that the components of a solid are close together and exert large attractive forces on each other.

The properties of liquids lie somewhere between those of solids and gases but not midway between, as can be seen from some of the properties of the three states of water.

For example, compare the enthalpy change for the melting of ice at 08C (the heat of fusion) with that for vaporizing liquid water at 1008C (the heat of vaporization):

 
These values show a much greater change in structure in going from the liquid to the gaseous state than in going from the solid to the liquid state. This suggests that there are extensive attractive forces among the molecules in liquid water, similar to but not
as strong as those in the solid state.

 
The conclusion is clear. The liquid and solid states show many similarities and are strikingly different from the gaseous state.

Intermolecular Forces
Let's consider the properties of the condensed states of matter (liquids and solids) and the forces that cause the aggregation of the components of a substance to form a liquid or a solid. These forces may involve covalent or ionic bonding, or they may involve weaker interactions usually called intermolecular forces (because they occur between, rather than within, molecules).

  
It is important to recognize that when a substance such as water changes from solid to liquid to gas, the molecules remain intact. The changes in states are due to changes in the forces among the molecules rather than in those within the molecules. In ice, as
we will see later in this chapter, the molecules are virtually locked in place, although they can vibrate about their positions. If energy is added, the motions of the molecules increase, and they eventually achieve the greater movement and disorder characteristic of liquid water.
 
The ice has melted. As more energy is added, the gaseous state is eventually reached, with the individual molecules far apart and interacting relatively little.However, the gas still consists of water molecules. It would take much energy to overcome the covalent bonds and decompose the water molecules into their component
atoms. This can be seen by comparing the energy needed to vaporize 1 mole of liquid water (40.7 kJ) with that needed to break the OOH bonds in 1 mole of water molecules (934 kJ).
 
Dipole–Dipole ForcesMolecules with polar bonds often behave in an electric field
as if they had a center of positive charge and a center of negative charge. That is, they exhibit a dipole moment. Molecules with dipole moments can attract each other electrostatically by lining up so that the positive and negative ends are close to each other 




This is called a dipole–dipole attraction. In a condensed state such as a liquid, where many molecules are in close proximity, the dipoles find the best compromise between attraction and repulsion. That is, the molecules orient themselves to maximize the interactions and to minimize and interactions.
 
Dipole–dipole forces are typically only about 1% as strong as covalent or ionic bonds, and they rapidly become weaker as the distance between the dipoles increases. At low pressures in the gas phase, where the molecules are far apart, these forces are relatively unimportant.
 
Hydrogen Bonding
Particularly strong dipole–dipole forces, however, are seen among molecules in which hydrogen is bound to a highly electronegative atom, such as nitrogen, oxygen, or fluorine. 

Reason for strength of hydrogen bonding
Two factors account for the strengths of these interactions: the great polarity of the bond and the close approach of the dipoles, allowed by the very small size of the hydrogen atom. Because dipole–dipole attractions of this type are so unusually strong, they are given a special name—hydrogen bonding. Figure   shows hydrogen bonding among water molecules, which occurs between the partially positive H atoms and the lone pairs on adjacent water molecules.  



Hydrogen bonding has a very important effect on physical properties. For example,the boiling points for the covalent hydrides of the elements in Groups 4A, 5A, 6A, and 7A are given in Fig.

Note that the nonpolar tetrahedral hydrides of Group 4A show a steady increase in boiling point with molar mass (that is, in going down the group), whereas, for the other groups, the lightest member has an unexpectedly high boiling point. Why? The answer lies in the especially large hydrogen bonding interactions that exist among the smallest molecules with the most polar X-H bonds. These
unusually strong hydrogen bonding forces are due primarily to two factors. 
One factor is the relatively large electronegativity values of the lightest elements in each group,which leads to especially polar X-H bonds. The second factor is the small size of the first element of each group, which allows for the close approach of the dipoles, further strengthening the intermolecular forces. Because the interactions among the molecules containing the lightest elements in Groups 5A and 6A are so strong, an unusually large quantity of energy must be supplied to overcome these interactions and separate the molecules to produce the gaseous state. These molecules will remain together in the liquid state even at high temperatures—hence the very high boiling points.
Hydrogen bonding is also important in organic molecules (molecules with a carbon chain backbone). For example, the alcohols methanol (CH3OH) and ethanol (CH3CH2OH) have much higher boiling points than would be expected from their
molar masses because of the polar O-H bonds in these molecules, which produce hydrogen bonding.



 London Dispersion Forces 

Even molecules without dipole moments must exert forces on each other. We know this because all substances—even the noble gases—exist in the liquid and solid states under certain conditions. The forces that exist among noble gas atoms and nonpolar
molecules are called London dispersion forces. To understand the origin of these forces, let’s consider a pair of noble gas atoms. Although we usually assume that the electrons of an atom are uniformly distributed about the nucleus, this is apparently not
true at every instant. As the electrons move about the nucleus, a momentary nonsymmetrical electron distribution can develop that produces a temporary dipolar arrangement of charge. The formation of this temporary dipole can, in turn, affect the electron
distribution of a neighboring atom. That is, this instantaneous dipole that occurs accidentally in a given atom can then induce a similar dipole in a neighboring atom Fig.


 

This phenomenon leads to an interatomic attraction that is relatively weak and short-lived but that can be very signifcant especially for large atoms.
For these interactions to become strong enough to produce a solid, the motions of the atoms must be greatly slowed down. This explains, for instance, why the noble gas elements have such low freezing points.
Note from Table 
 
 that the freezing point rises going down the group. The principal cause for this trend is that as the atomic number increases, the number of electrons increases, and there is an increased chance of the occurrence of momentary dipole interactions. We describe this phenomenon using the term polarizability, which indicates the ease with which the electron “cloud” of an atom can be distorted to give
a dipolar charge distribution. Thus we say that large atoms with many electrons exhibit a higher polarizability than small atoms. This means that the importance of London dispersion forces increases greatly as the size of the atom increases.
These same ideas also apply to nonpolar molecules such as H2, CH4, CCl4, and CO2.Since none of these molecules has a permanent dipole moment, their principal means of attracting each other is through London dispersion forces.