As was the case for gaseous substances, the kinetic molecular theory may be used to explain the behavior of solids and liquids. In the following description, the term particle will be used to refer to an atom, molecule, or ion. Consider these two aspects of the molecular-level environments in solid, liquid, and gaseous matter:.
The differences in the properties of a solid, liquid, or gas reflect the strengths of the attractive forces between the atoms, molecules, or ions that make up each phase. The phase in which a substance exists depends on the relative extents of its intermolecular forces IMFs and the kinetic energies KE of its molecules. IMFs are the various forces of attraction that may exist between the atoms and molecules of a substance due to electrostatic phenomena, as will be detailed in this module.
Figure 1 illustrates how changes in physical state may be induced by changing the temperature, hence, the average KE, of a given substance.
Figure 1. Transitions between solid, liquid, and gaseous states of a substance occur when conditions of temperature or pressure favor the associated changes in intermolecular forces. Note: The space between particles in the gas phase is much greater than shown. Figure 2. Condensation forms when water vapor in the air is cooled enough to form liquid water, such as a on the outside of a cold beverage glass or b in the form of fog.
Figure 3. Gaseous butane is compressed within the storage compartment of a disposable lighter, resulting in its condensation to the liquid state.
As an example of the processes depicted in this figure, consider a sample of water.
3 Types of Intermolecular Forces
When gaseous water is cooled sufficiently, the attractions between H 2 O molecules will be capable of holding them together when they come into contact with each other; the gas condenses, forming liquid H 2 O.
For example, liquid water forms on the outside of a cold glass as the water vapor in the air is cooled by the cold glass, as seen in Figure 2. We can also liquefy many gases by compressing them, if the temperature is not too high.
The increased pressure brings the molecules of a gas closer together, such that the attractions between the molecules become strong relative to their KE. Consequently, they form liquids. Butane, C 4 H 10is the fuel used in disposable lighters and is a gas at standard temperature and pressure.
Finally, if the temperature of a liquid becomes sufficiently low, or the pressure on the liquid becomes sufficiently high, the molecules of the liquid no longer have enough KE to overcome the IMF between them, and a solid forms. A more thorough discussion of these and other changes of state, or phase transitions, is provided in a later module of this chapter.
This simulation is useful for visualizing concepts introduced throughout this chapter.Intermolecular forces or IMFs are physical forces between molecules. In contrast, intramolecular forces are forces between atoms within a single molecule. Intermolecular forces are weaker than intramolecular forces. The strength or weakness of intermolecular forces determines the state of matter of a substance e. There are three major types of intermolecular forces: London dispersion forcedipole-dipole interaction, and ion-dipole interaction.
Here's a closer look at these three intermolecular forces, with examples of each type. The London dispersion force, the force between two nonpolar molecules, is the weakest of the intermolecular forces.
The electrons of one molecule are attracted to the nucleus of the other molecule, while repelled by the other molecule's electrons. A dipole is induced when the electron clouds of the molecules are distorted by the attractive and repulsive electrostatic forces.
Example: A second example of London dispersion force is the interaction between nitrogen gas N 2 and oxygen gas O 2 molecules. The electrons of the atoms are not only attracted to their own atomic nucleus, but also to the protons in the nucleus of the other atoms. Dipole-dipole interaction occurs whenever two polar molecules get near each other.
The positively charged portion of one molecule is attracted to the negatively charged portion of another molecule. Since many molecules are polar, this is a common intermolecular force. A hydrogen atom of one molecule is attracted to an electronegative atom of another molecule, such as an oxygen atom in water.
Ion-dipole interaction occurs when an ion encounters a polar molecule. In this case, the charge of the ion determines which part of the molecule attracts and which repels. A cation or positive ion would be attracted to the negative part of a molecule and repelled by the positive part. An anion or negative ion would be attracted to the positive part of a molecule and repelled by the negative part.
Van der Waals forces are the interaction between uncharged atoms or molecules. The forces are used to explain the universal attraction between bodies, the physical adsorption of gases, and the cohesion of condensed phases. The van der Waals forces encompass intermolecular forces as well as some intramolecular forces including Keesom interaction, the Debye force, and the London dispersion force.
Share Flipboard Email. Anne Marie Helmenstine, Ph. Chemistry Expert. Helmenstine holds a Ph. She has taught science courses at the high school, college, and graduate levels. Facebook Facebook Twitter Twitter. Updated December 07, Key Takeaways: Intermolecular Forces Intermolecular forces act between molecules. In contrast, intramolecular forces act within molecules. Examples of intermolecular forces include the London dispersion force, dipole-dipole interaction, ion-dipole interaction, and van der Waals forces.The properties of liquids are intermediate between those of gases and solids, but are more similar to solids.
In contrast to intra molecular forces, such as the covalent bonds that hold atoms together in molecules and polyatomic ions, inter molecular forces hold molecules together in a liquid or solid. Intermolecular forces are generally much weaker than covalent bonds. Despite this seemingly low value, the intermolecular forces in liquid water are among the strongest such forces known! Given the large difference in the strengths of intra- and intermolecular forces, changes between the solid, liquid, and gaseous states almost invariably occur for molecular substances without breaking covalent bonds.
The properties of liquids are intermediate between those of gases and solids but are more similar to solids. Intermolecular forces determine bulk properties such as the melting points of solids and the boiling points of liquids. Liquids boil when the molecules have enough thermal energy to overcome the intermolecular attractive forces that hold them together, thereby forming bubbles of vapor within the liquid.
Similarly, solids melt when the molecules acquire enough thermal energy to overcome the intermolecular forces that lock them into place in the solid. Intermolecular forces are electrostatic in nature; that is, they arise from the interaction between positively and negatively charged species. Like covalent and ionic bonds, intermolecular interactions are the sum of both attractive and repulsive components.
Because electrostatic interactions fall off rapidly with increasing distance between molecules, intermolecular interactions are most important for solids and liquids, where the molecules are close together. These interactions become important for gases only at very high pressures, where they are responsible for the observed deviations from the ideal gas law at high pressures.
For more information on the behavior of real gases and deviations from the ideal gas law. In this section, we explicitly consider three kinds of intermolecular interactions: There are two additional types of electrostatic interaction that you are already familiar with: the ion—ion interactions that are responsible for ionic bonding and the ion—dipole interactions that occur when ionic substances dissolve in a polar substance such as water.
The first two are often described collectively as van der Waals forces. Polar covalent bonds behave as if the bonded atoms have localized fractional charges that are equal but opposite i. If the structure of a molecule is such that the individual bond dipoles do not cancel one another, then the molecule has a net dipole moment.
On average, however, the attractive interactions dominate. In addition, the attractive interaction between dipoles falls off much more rapidly with increasing distance than do the ion—ion interactions.
Their structures are as follows:. Compare the molar masses and the polarities of the compounds. Compounds with higher molar masses and that are polar will have the highest boiling points.
The first compound, 2-methylpropane, contains only C—H bonds, which are not very polar because C and H have similar electronegativities. It should therefore have a very small but nonzero dipole moment and a very low boiling point. As a result, the C—O bond dipoles partially reinforce one another and generate a significant dipole moment that should give a moderately high boiling point. The C—O bond dipole therefore corresponds to the molecular dipole, which should result in both a rather large dipole moment and a high boiling point.
Thus far we have considered only interactions between polar molecules, but other factors must be considered to explain why many nonpolar molecules, such as bromine, benzene, and hexane, are liquids at room temperature, and others, such as iodine and naphthalene, are solids. What kind of attractive forces can exist between nonpolar molecules or atoms?
This question was answered by Fritz London —a German physicist who later worked in the United States. InLondon proposed that temporary fluctuations in the electron distributions within atoms and nonpolar molecules could result in the formation of short-lived instantaneous dipole momentswhich produce attractive forces called London dispersion forces between otherwise nonpolar substances.Intermolecular Forces.
Water is the only substance we routinely encounter as a solid, a liquid, and a gas. At low temperatures, it is a solid in which the individual molecules are locked into a rigid structure. As we raise the temperature, the average kinetic energy of the molecules increases, which increases the rate at which these molecules move.
There are three ways in which a water molecule move: 1 vibration, 2 rotation, and 3 translation. Water molecules vibrate when H--O bonds are stretched or bent. Rotation involves the motion of a molecule around its center of gravity. Translation literally means to change from one place to another. It therefore describes the motion of molecules through space. To understand the effect of this motion, we need to differentiate between intramolecular and intermolecular bonds.
The covalent bonds between the hydrogen and oxygen atoms in a water molecule are called intramolecular bonds. The prefix intra - comes from the Latin stem meaning "within or inside. The bonds between the neighboring water molecules in ice are called intermolecular bondsfrom the Latin stem meaning "between. The intramolecular bonds that hold the atoms in H 2 O molecules together are almost 25 times as strong as the intermolecular bonds between water molecules.
All three modes of motion disrupt the bonds between water molecules. As the system becomes warmer, the thermal energy of the water molecules eventually becomes too large to allow these molecules to be locked into the rigid structure of ice. At this point, the solid melts to form a liquid in which intermolecular bonds are constantly broken and reformed as the molecules move through the liquid.
Eventually, the thermal energy of the water molecules becomes so large that they move too rapidly to form intermolecular bonds and the liquid boils to form a gas in which each particle moves more or less randomly through space. The difference between solids and liquids, or liquids and gases, is therefore based on a competition between the strength of intermolecular bonds and the thermal energy of the system.
At a given temperature, substances that contain strong intermolecular bonds are more likely to be solids. For a given intermolecular bond strength, the higher the temperature, the more likely the substance will be a gas. The kinetic theory assumes that there is no force of attraction between the particles in a gas. If this assumption were correct, gases would never condense to form liquids and solids at low temperatures. In the Dutch physicist Johannes van der Waals derived an equation that not only included the force of attraction between gas particles but also corrected for the fact that the volume of these particles becomes a significant fraction of the total volume of the gas at high pressures.
The van der Waals equation is used today to give a better fit to the experimental data of real gases than can be obtained with the ideal gas equation. But that wasn't van der Waals's goal. He was trying to develop a model that would explain the behavior of liquids by including terms that reflected the size of the atoms or molecules in the liquid and the strength of the bonds between these atoms or molecules.
The weak intermolecular bonds in liquids and solids are therefore often called van der Waals forces. These forces can be divided into three categories: 1 dipole-dipole, 2 dipole-induced dipole, and 3 induced dipole-induced dipole.
Many molecules contain bonds that fall between the extremes of ionic and covalent bonds. The difference between the electronegativities of the atoms in these molecules is large enough that the electrons aren't shared equally, and yet small enough that the electrons aren't drawn exclusively to one of the atoms to form positive and negative ions. The bonds in these molecules are said to be polarbecause they have positive and negative ends, or poles, and the molecules are often said to have a dipole moment.
HCl molecules, for example, have a dipole moment because the hydrogen atom has a slight positive charge and the chlorine atom has a slight negative charge. Because of the force of attraction between oppositely charged particles, there is a small dipole-dipole force of attraction between adjacent HCl molecules.Danny R. There are two intermolecular forces that are available right now. The London Forces, also are known as the London Dispersion Force, is known to be a type of force that you can get between the various atoms and molecules that are available.
Dipole forces, on the other hand, is a type of force that is considered to be strong. This is between one end of the polar molecule and the negative end of another polar molecule.
The forces are usually strong and are usually stronger as compared to the London Forces. It is ideal to know the various forces that are available just to be sure.
Forgot your password? Speak now. Login Sign Up Free. Quiz Maker All Products. Discuss Science Chemistry. What are the intermolecular forces present in the molecule CH2F2? Post Your Answer. The first thing that you need to do first is to figure out what type of compound you are looking at.
You should realize that CH2F2 is known to be a polar molecule. The type of force that you can expect to have is dipole interaction. Aside from this type of interaction, you can expect that there are also going to be other forces that are available. If you would check something that is related to this molecule, which is CF4, you will realize that the structure is tetrahedral. This means that the force that will be the strongest for this one will be the London Dispersion Force.
Continue Reading. James Answered Mar 16, The correct answers to this question are London forces and dipole-dipole forces. London forces may also be known as LDF or London dispersion forces. The name of the forces comes after Fritz London, who was a physicist.
London forces are forces that act between molecules and atoms. LDF's are apart of the van der Waals forces. Dipole-dipole forces are forces that are attractive. They are a positive end of a polar molecule and a negative end of a polar molecule.Work in groups on these problems. You should try to answer the questions without referring to your textbook.
If you get stuck, try asking another group for help. Most substances can exist in either gas, liquid, or solid phase under appropriate conditions of temperature and pressure. The phase that we see under ordinary conditions room temperature and normal atmospheric pressure is a result of the forces of attraction between molecules or ions comprising the substance. The strength of these attractions also determines what changes in temperature and pressure are needed to effect a phase transition.
The behavior of a pure substance in any of its phases is altered when it is mixed with other substances to make solutions. Solutions are homogeneous mixtures, which can occur in any phase. But most often in chemistry we are dealing with solutions that are in the liquid phase. An understanding of the processes by which solutions form and of how their properties differ from their pure-substance components is useful in many real-life applications of materials.
The tendency of a substance to be found in one state or the other under certain conditions is largely a result of the forces of attraction that exist between the particles comprising it.
We will concentrate on the forces between molecules in molecular substances, which are called intermolecular forces. Forces that exist within molecules, such as chemical bonds, are called intramolecular forces.
The greater the strength of the intermolecular forces, the more likely the substance is to be found in a condensed state ; i. As we have seen, the model of an ideal gas assumes that the gas particles molecules or atoms have virtually no forces of attraction between them, are widely separated, and are constantly moving with high velocity and kinetic energy.
In truth, there are forces of attraction between the particles, but in a gas the kinetic energy is so high that these cannot effectively bring the particles together. With stronger intermolecular forces or lower kinetic energy, those forces may draw molecules closer together, resulting in a condensed phase.
Going from gas to liquid to solid, molecular velocities and particle separations diminish progressively as structural order increases. In the case of liquids, molecular attractions give rise to viscositya resistance to flow.
Solids have stronger intermolecular forces, making them rigid, with essentially no tendency to flow. Although the mix of types and strengths of intermolecular forces determines the state of a substance under certain conditions, in general most substances can be found in any of the three states under appropriate conditions of temperature and pressure. Changing those conditions can induce a change in the state of the substance, called a phase transition.Intermolecular forces IMF are the forces which mediate interaction between atomsincluding forces of attraction or repulsion which act between atoms and other types of neighboring particles, e.
Intermolecular forces are weak relative to intramolecular forces — the forces which hold a molecule together. For example, the covalent bondinvolving sharing electron pairs between atoms, is much stronger than the forces present between neighboring molecules. Both sets of forces are essential parts of force fields frequently used in molecular mechanics.
The investigation of intermolecular forces starts from macroscopic observations which indicate the existence and action of forces at a molecular level.
These observations include non-ideal-gas thermodynamic behavior reflected by virial coefficientsvapor pressureviscositysuperficial tension, and absorption data. The first reference to the nature of microscopic forces is found in Alexis Clairaut 's work Theorie de la Figure de la Terre. Information on intermolecular forces is obtained by macroscopic measurements of properties like viscosity, pressure, volume, temperature PVT data. The link to microscopic aspects is given by virial coefficients and Lennard-Jones potentials.
A hydrogen bond is the attraction between the lone pair of an electronegative atom and a hydrogen atom that is bonded to an electronegative atom, usually nitrogenoxygenor fluorine. However, it also has some features of covalent bonding: it is directional, stronger than a van der Waals force interaction, produces interatomic distances shorter than the sum of their van der Waals radiiand usually involves a limited number of interaction partners, which can be interpreted as a kind of valence.
The number of Hydrogen bonds formed between molecules is equal to the number of active pairs. The molecule which donates its hydrogen is termed the donor molecule, while the molecule containing lone pair participating in H bonding is termed the acceptor molecule. The number of active pairs is equal to the common number between number of hydrogens the donor has and the number of lone pairs the acceptor has.
Though both not depicted in the diagram, water molecules have two active pairs, as the oxygen atom can interact with two hydrogens to form two hydrogen bonds. Intramolecular hydrogen bonding is partly responsible for the secondarytertiaryand quaternary structures of proteins and nucleic acids. It also plays an important role in the structure of polymersboth synthetic and natural.
The attraction between cationic and anionic sites is a noncovalent, or intermolecular interaction which is usually referred to as ion pairing or salt bridge. Most salts form crystals with characteristic distances between the ions; in contrast to many other noncovalent interactions salt bridges are not directional and show in the solid state usually contact determined only by the van der Waals radii of the ions. Dipole—dipole interactions are electrostatic interactions between molecules which have permanent dipoles.
This interaction is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved. These interactions tend to align the molecules to increase attraction reducing potential energy.
An example of a dipole—dipole interaction can be seen in hydrogen chloride HCl : the positive end of a polar molecule will attract the negative end of the other molecule and influence its position.Practice Exercise p 436 Intermolecular Forces
Polar molecules have a net attraction between them. Often molecules contain dipolar groups of atoms, but have no overall dipole moment on the molecule as a whole. This occurs if there is symmetry within the molecule that causes the dipoles to cancel each other out. This occurs in molecules such as tetrachloromethane and carbon dioxide.