12 -- Part 1: Intermolecular Forces: Liquids and Solids
The properties of liquids are related to intermolecular forces.
Discuss how temperature changes affect the state of a substance.
The triple, melting, boiling, and critical points are on a phase diagram.
Give one example of each of the different types of solid.
The Born-Fajans-Haber cycle can be used.
Solids, liquids, and gases were compared at the macroscopic and microscopic levels.
Two of the many natural slow down when we make ice cubes by placing water in a tray in a freezer.
The water becomes ice when the intermolecular forces between the mol phenomena take over.
Chapter include the more energy from the surroundings is absorbed by the water molecule, which ordered structure of the solid overcome the intermolecular forces within the ice cube and enter the liquid compared with the liquid state.
We wanted conditions in which state and the variation of the intermolecular forces to be insignificant in our study of gases.
This approach allowed us to density with the state of describe gases with the ideal gas equation.
We need to identify the various intermolecular forces before we can describe the other states of matter.
Some interesting properties of liquids are related to the strengths of these forces.
Intermolecular forces cause gas behavior to depart from ideality at high pressures and low temperatures.
A gas condenses to a liquid when the forces are strong.
The intermolecular forces keep the molecule in close proximity to each other so that they are confined to a definite volume.
There are intermolecular forces that contribute to the term a(n>V)2 in the van der Waals equation for nonideal gases.
The dipole moment and polar izability are essential for describing the physical basis of attractive intermolecular forces.
The distribution of electron density within a molecule is described by these properties.
We'll review some of the points we made about these two properties before we discuss different types of intermolecular interactions.
A molecule has a dipole moment when the centers of positive and negative charge don't match.
It's important to distinguish the atom.
Its electron cloud is spread out.
The polarizability of a molecule is related to how easily it is displaced and the volume of its charge cloud.
A polar molecule has a permanent dipole moment, so it tends to line up with the positive end of one dipole directed toward the negative ends of neighboring dipoles.
The partial ordering of the molecule can cause a substance to persist at higher than expected temperatures.
Figure 12-2 depicts the molecule and the CHF as an example of the influence of dipole-dipole interactions on physical properties.
The approach of a polar molecule will cause the electronic charge cloud of a molecule to be distorted.
The interaction between a dipole and an inducer causes the attraction between a pair of polar molecules to be slightly greater than predicted by the interaction of the unmodified dipole moments.
The blue arrow shows the dipole moment.
The difference in boiling point is a result of one difference in the dipole moments.
The intermolecular attractions between CHF3 molecule include 3.34 * 10-30 C m.
The intermolecular interactions described in the previous section are important, but they are usually not the most important.
The result is a dipole attraction.
The dynamic nature of polarizability is an important characteristic that we have not yet mentioned.
It is1-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-65561-6556 A normally non polar species becomes temporarily polar due to the displacement of electrons.
The molecule has an instantaneous dipole moment.
After this, electrons in a neighboring molecule may be displaced.
The events lead to an intermolecular force of attraction.
As polarizability increases, dispersion forces become stronger.
The data in Table 12.1 suggests that substances made up of larger, more polarizable molecule tend to have higher boiling points.
The data in Table 12.1 shows that polarizability increases with mass.
This correlation is not new.
A molecule with a lot of atoms has a lot of electrons.
A large molecule has a large charge cloud.
The melting and boiling points of molecule tend to increase with increasing mass because of the correlation between polarizability and mass.
The mole has a boiling point of 4 K, while radon has a boiling point of 222 U.
polarizability volume is sometimes referred to as polarizability.
The units of polarizability have the units of volume.
The measure of the atomic or molecular volume is provided by polarizability.
Some features that we cannot explain by the types of intermolecular forces considered to be at this point are shown in Hydrogen Bonding Figure 12-4, in which the boiling points of a series of similar compounds are plotted as a function of molecular mass.
The polarizabilities of C H and CO are expected to be 888-609- 888-609- 888-609-.
We have to look for a different factor to base our prediction on.
The next thing to consider is the chemistry of the molecule.
The electronegativity difference between C and H is so small that we expect butane to be nonpolar.
A strong carbon-to-oxygen dipole is found in one of the molecules.
It is helpful to sketch the structure of a molecule to see if symmetri Butane and acetone cal features cause bond dipoles to cancel.
The diagrams show the acetone molecule as a polar molecule.
The O bond can't be offset by other bonds.
acetone is polar.
One polar and one non indicates regions of high polar, and we expect the polar substance to have the higher boiling negative electrostatic point.
When comparing the properties of different substances, we must consider the various types of intermolecular forces and the factors that affect the strength of each type of force.
Although it wasn't important, the three-dimensional shape of a molecule is a very important consideration and it is usually necessary to sketch the molecule's structure to see how it plays a role.
The expected order of increasing boiling point is C8H18, CH3 CH2 CH3 and C6H5CHO.
There are exceptions in three groups.
The boiling points of NH3 and H2O are the same as those of any other hydride in their group.
The main points in the figure are outlined.
The H atom is between the two F atoms.
The H atom is weakly bonded to the F atom of a nearby molecule.
The F atom has a pair of electrons.
The temperature is about 180deg.
F bonds to an atom that attracts electron density away from the H nucleus.
The H nucleus, a protons, can be attracted to a lone pair of electrons by a neighboring molecule.
In general, a hydrogen bond is depicted as X-H Y--, where the three many of the HF molecule are dots.
The fragment is known as the hydro.
The hydrogen as part of the hydrogen to one F atom is accepted by each H atom as a single bond.
Compared with other intermolecular forces, hydrogen another F atom through a bonds are relatively strong, having energies of the order of 15 to 40 kJ mol-1.
The strength of single covalent bonds is greater than 150 kJ mol-1.
Water is the most common substance in which hydrogen bonding occurs.
Water is held in a rigid but open structure by hydrogen bonds.
Only a small portion of the hydrogen bonds are broken when ice is melted.
The fusion of ice 16.01 kJ mol-12 has a relatively low heat.
If all the hydrogen bonds were to break, it would be much more.
Some of the hydrogen bonds are broken when ice is melted.
The increase in density when ice is melted is accounted for by the fact that the water molecule is more compactly arranged.
The number of H2O molecule per unit volume in the liquid is greater than in the solid.
The acceptor is the H water molecule.
Each water molecule is linked to four others through hydrogen bonds.
There is an arrangement.
Each H atom is located along a line joining two O atoms, but closer to one O atom than the other.
O atoms are arranged in hexagonal rings.
The hexagonal shapes of snowflakes are similar to this pattern.
The liquid has more water in it than the solid.
The hydrogen bonds break when liquid water is heated above the melting point.
The density of the liquid water continues to increase.
The maximum density of liquid water is 3.98 degC.
The density of the water decreases as the temperature increases.
The ice Solid and liquid densities over the top of the lake insulate the water below from further heat loss.
The ice cubes are floating.
Without hydrogen bonding, all lakes would freeze from the bottom up; on liquid water, fish, small bottom-feeding animals, and aquatic plants would not know one; ice is less dense survive the winter.
Solid paraffin sinks to the bottom of the beaker because of Hydrogen Bonding.
Water's properties are affected by hydrogen bonding.
There are many others.
The heat of vaporization is low because not all the hydrogen bonds are disrupted.
Hydrogen bonding can be used to explain certain trends in viscosity.
The OH group can form a hydrogen bond to the alcohol molecule.
We expect the diol to flow more slowly because of the stronger intermolecular forces.
The three common alcohols are compared.
1 N s m-2 is 10 P and 1000 cP.
There is a chance that a molecule with an H atom of salicylic acid showing covalently bonding to one highly negative atom and another highly positive atom nearby in the same molecule.
The OH group is too far away from the molecule.
There is no hydrogen bonding in this situation.
Some chemical reactions in living matter involve complex structures and need to be broken and re- formed.
We will find both inter and inter hydroxybenzoic acid.