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Intermolecular Forces and Liquid Properties

An overview of the different types of intermolecular forces and their effect on the properties of liquids.

Image by Aaron Burden

The phases or states of matter are gas, liquid, and solid. The state of matter of a substance depends on how fast the particles in the substance are moving (kinetic energy) and the interactions between particles (intermolecular forces). The fundamental difference between the states of matter is the strength of the intermolecular forces (IMFs) and whether the IMFs can overcome the particle's kinetic energy. A condensed phase of matter is one where the particles are close together, such as with solids and liquids. Therefore, it stands to reason that the IMFs in solids are more substantial than in liquids, while gases have the weakest IMFs. Below is a summary of the characteristics of each state of matter.

Table 1: The properties of the states of matter

Forces within and in-between Molecules

Intramolecular: Intra means the same. Intramolecular interactions are the bonds that hold atoms together within a molecule or compound and are very strong.

 

Intermolecular: Inter means different. Intermolecular forces are attractive forces between two molecules (or two portions of the same molecule if it is sufficiently large). This force is generally significantly weaker than intramolecular forces, with the strength of the force depending on the magnitude of charges (or partial charges), with the idea that opposites attract.

 

Intermolecular forces (IMFs) hold molecules together; the more substantial the IMFs, the better the molecules stick to each other. Since the basis of IMFs is charge attractions, the greater the charge and subsequent attraction, the stronger the intermolecular force. The intermolecular forces include dispersion forces (ion-induced dipole, dipole-induced dipole, induced dipole-induced dipole), dipole-dipole, hydrogen-bonding, and ion-dipole forces.

Figure 1: Strengths of intra and intermolecular forces in kJ per mole.

Different species exhibit different amounts of charge. For example, ions gain or lose electrons and possess a full, whole number charge (such as +2 or -1). Conversely, polar molecules have not gained or lost electrons; they merely have an unequal sharing of electrons. The electron density imbalance across polar molecules results in partial positive and partial negative areas within the molecule, not complete charges as is seen with ions.

Electrostatic Forces

Electrostatic forces, also known as Coulomb forces, are the attraction or repulsion of particles or objects because of their electric charge. Two like electric charges, positive-positive or negative-negative, repel each other along a straight line between their centers, while opposites, positive-negative, attract. Because the electrostatic force is the force between charged particles, IMFs are a consequence of the electrostatic force. Therefore, understanding the electrostatic force will give an individual an understanding of how IMFs work.

 

The charge of the particles involved and the distance between the particles determines the magnitude of the electrostatic force:

Figure 2: Electrostatic forces (F21 and F12) between two charged particles (q1 and q2).

London Dispersion Forces

The most prevalent IMF is the London dispersion force (LDF), which occurs in all polar and nonpolar molecules. LDFs (also called just dispersion forces) are electrostatic interactions caused by the attraction of temporary dipoles to other temporary dipoles (or other charged species). A temporary dipole is a short-lived dipole resulting from the movement of electrons in an atom or molecule (remember that electrons are always moving). This movement of electrons can cause a momentary shift in electron density, creating an instantaneous, but not permanent, dipole within a molecule. Sometimes these dipoles are instantaneous, occurring due to a natural shift of electrons. Sometimes, the temporary dipoles are induced (caused by interaction with a charged species attracting or repelling the electrons within the atom or molecule). A temporary dipole can also cause a neighboring molecule to shift its electron density (electrons repel electrons) and induce a dipole in that other molecule.

Figure 3: The formation of a temporary dipole in nonpolar molecules.

The above image shows nonpolar molecules forming temporary or instantaneous dipoles. Note that any species with electrons can create a temporary dipole (even polar molecules) as it is a momentary shift in the electron cloud, and all atoms and molecules have an electron cloud. However, species with more significant electron clouds make temporary dipoles that are longer lived with partial charges with large magnitudes.

Figure 4: Instantaneous dipoles averaging out to a zero dipole moment.

The ease at which a molecule can have an instantaneous dipole depends upon the number of electrons floating around the structure. More electrons mean it's easier to shift them (the electrons) and cause a temporary dipole.  There are three types of LDFs: induced dipole-induced dipole, dipole-induced dipole, and ion-induced dipole.

Induced dipole-induced dipole interactions result from one species with electrons interacting with another species with electrons, causing the shift of electron density in the molecules and inducing dipoles (temporary dipoles). The more electrons a species has, the stronger and longer lived the induced dipoles. The forces due to temporary dipoles are reasonably weak, though the strength varies significantly due to significant variances in molar masses and the number of electrons (1-10 kJ/mol). Additionally, this force is seen in all molecules, though it is the only IMF that exists between nonpolar molecules. Below is an image of the induced-dipole induced dipole force between two molecules or atoms.

Figure 5: Induced dipole-induced dipole forces.

The next LDF is the dipole-induced dipole interaction, an electrostatic force between a permanent dipole and the induced dipole of a nonpolar molecule. With the dipole-induced dipole force, a molecule with a permanent dipole shifts the electron cloud on the nonpolar molecule to create a temporary dipole. It is another weak interaction (2 – 10 kJ / mol), and is the IMF between polar molecules and other molecules (usually, we consider the other molecules to be nonpolar ones). Below is an image depicting the interaction.

Figure 6: Dipole-induced dipole forces.

Ion-induced dipole is an electrostatic force between an ion and the induced dipole of a nonpolar molecule. When an ion comes in contact with a nonpolar molecule (with no net dipole), a cation or an anion can cause the electrons within the nonpolar molecule to shift, inducing a dipole. Cations attract the molecule's electrons, while anions repel them. Ion-induced dipole is a weak to moderate force (3 – 15 kJ/mol) and is the force that exists between ions and molecules (generally nonpolar molecules). The interaction between a nonpolar molecule and a cation (Na+) or anion (Cl-) is shown in the image below.

Figure 7: Ion-induced dipole forces.

Polarizability is the ease with which the electron cloud can be distorted by an outside force. When an electron cloud can easily be distorted (shifted to cause an induced dipole), it is said to be polarizable. Large, spongy molecules (those with great molar masses and surface areas) are more polarizable and have more significant dispersion forces. Below shows an image of the polarizability in both F2 and I2.

Figure 8: Polarizability in F2 and I2 molecules.

Figure 9: Polarizability and molar mass trend in group 7A diatomics.

The molecule F2, , a gaseous molecule at room temperature, is not very polarizable due to its small molar mass, ~38 g/mol. In contrast, I2, which has a much greater molar mass, ~254 g/mol, is significantly more polarizable and, as such, exists as a solid at room temperature. From this, we can see that the more polarizability a molecule is, the higher its melting and boiling points. This trend of increased polarizability leading to increased melting and boiling points. Furthermore, the trend of increased polarizability leading to increased melting and boiling points is evident in all molecules, not just the halogen. For instance, if you were to look at the boiling points of alkanes, increased polarizability (due to increases in molecular weight and surface area) raises the alkane's boiling point (see the table below).

Table 2: Boiling points of hydrocarbons with different molecular weights and surface areas.

Dipole-Dipole Forces (Including Hydrogen Bonding)

The dipole-dipole IMF is an electrostatic force between the partial charge of one end of a polar molecule and the partial charge of one end of another polar molecule. The dipole-dipole force exists between all (and only) polar molecules and is considered a weak IMF (2-15 kJ/mol). The attraction occurs between the positive pole of one of the molecules with the negative pole of another. Like induced dipole-induced dipole forces, the molecules involved in dipole-dipole forces can be the same (pure substances) or different (mixtures). 

The image to the left shows the dipole-dipole interaction between two formaldehyde molecules.  The bond between oxygen and carbon is polar because of the difference in electronegativity. As a result of the bond dipole (and overall molecular dipole), the more electronegative oxygen atom of one formaldehyde molecule has a negative partial charge (the negative pole of the dipole) and is attracted to the partially positive carbon (forms the positive pole due to carbon's lower electronegativity).

Figure 10: Dipole-dipole forces

Hydrogen bonding is a super strong dipole-dipole interaction (10-40 kJ/mol). The interaction results from an attraction between a hydrogen atom bound to a highly electronegative atom and another electronegative atom bearing a nonbonding electron pair. The electronegative atom pulls so much electron density away from the hydrogen that it shrinks its radius, allowing the lone pairs of N, O, or F to get close to the electron-deficient hydrogen atom, forming the "bond." Hydrogen bonding is a more potent attraction than dipole-dipole or London dispersion forces. Note that for a hydrogen bond to be possible, a molecule (or two molecules together) must possess a hydrogen bond donor and a hydrogen bond acceptor. A hydrogen bond donor is any hydrogen atom directly attached to an exceedingly electronegative atom (nitrogen, oxygen, or fluorine). The hydrogen bond acceptor must be a highly electronegative atom (N, O, F) and have an attached pair of nonbonding electrons.

hydrogen bond examples.bmp

Figure 11: Hydrogen bonding in pure substances (a. water) and mixtures (b. water and formaldehyde; c. water and ammonia).

Hydrogen bonding in a. pure water, b. in a mixture of formaldehyde and water, and c. in a mixture of ammonia and water. Hydrogen bond donors are colored red, while hydrogen bond acceptors are blue (any atoms that are black are unable to participate in hydrogen bonding).

You can use the flowchart below to assist you in determining the dominant intermolecular force in a pure substance or mixture. Please note that the flowchart is for the DOMINANT IMF, but more than one IMF is often possible. For instance, everything has dispersion forces, and all polar molecules have dipole-dipole interactions.

Flowchart 1: Determining the dominant IMF in a pure substance or a mixture.

Comparing and Identifying Intermolecular Forces

Shape and molecular mass have the most significant influence on dispersion forces. Therefore, if molecules are similar in mass and shape (consider factors such as the molecules being linear or branched), the magnitude of their dispersion forces will be nearly equal.

 

If mass and shape are similar, the molecule's polarity will be the determining factor in the strength of the IMFs of that molecule. IMF strength will increase with increasing polarity (think electrostatic force).

 

Suppose substances have dramatically different molar masses or shapes (and there is no hydrogen bonding). In that case, the strength of the dispersion forces will determine the strength of the intermolecular forces within the substance. The strength of intermolecular forces increases with increasing molar mass and surface area (linear molecules have larger surface areas than branched molecules).

 

The strength of IMFs in a molecule increases with increasing numbers of hydrogen bond sites. The more possible hydrogen bonds a molecule can make, the stronger the IMFs within that molecule.

Below is a table comparing the strengths of the different bonds and intermolecular forces.

Table 3: All the types of interactions between particles and what kinds of substances contain them.

Identify and rank the strength of the IMF in the following pure substances or mixtures? (Highlight text to see the answer)

Nitrogen (N2 (g)) = London dispersion forces (induced dipole – induced dipole) = weakest, 5

H2O(l) = Hydrogen bonding, dipole-dipole, LDF = 2
Heptane (C7H16 (l)) = LDFs (induced dipole – induced dipole) = 4
H2O(l) & NaCl = Ion – dipole, LDF = strongest, 1
N2 (g) & H2O(l) = London dispersion forces (dipole – induced dipole) = 3

Liquid Properties

Intermolecular forces play a direct role in the properties of gases, liquids, and solids. For example, the IMFs determine what phase of matter a substance will be in at a given temperature, in addition to determining the particles' behavior. In the section below, we will look at key liquid properties (viscosity, surface tension, and capillary action, and the role IMFs play in those properties.

The image to the right shows someone pouring out a viscous liquid. The viscosity, or resistance of a liquid to flow, is responsible for some fluids, such as honey, having a characteristic thickness. A thicker liquid will flow more slowly than a thinner one. For instance, motor oils come in many different thicknesses, which a consumer can determine by looking at the XW-XX notation on the bottle (W is the viscosity rating at cold temperatures; a larger value indicates a greater viscosity).

Figure 12: The viscosity of oil.

motor oil.jpg

Figure 13: Viscosities of different motor oils.

As mentioned above, viscosity is the resistance of a liquid to flow, and has an SI unit of kilograms per meters per second. The basis for the viscosity or thickness of a liquid is the IMFs between particles and the temperature. Specifically, the greater the intermolecular forces the greater the viscosity, which in turn means the lower the temperature, the greater the viscosity (remember that a larger temperature indicates a greater amount of kinetic energy, resulting in weaker intermolecular forces). The way that viscosity works is that when molecules are greatly attracted to each other, they stick to each other more, decreasing the ability for the molecules to slide past one another (resistance to flow). Shape and flexibility also play a role. Molecules that are more likely to become entangled (like with unbranched, flexible molecules), the harder it is for those moleculs to slide past one another (think spaghetti noodles). Two examples of highly viscous fluids are shown below. In the case of glycerol, there are lots of hydrogen bonding positions (donors in red, acceptors in blue), which holds molecules together. With synthetic motor oils, the long hydrocarbon chains intertwine owing to their structure and strong dispersion forces (due to the large size of the molecules).

Figure 14: Molecular structures of synthetic motor oil and glycerol.

The next liquid property is surface tension, which refers to a fluid's elastic tendency to acquire the minimum surface area possible. Surface tension arises from an imbalance in intermolecular forces at the boundary between liquid and another substance (usually air). The liquid's surface area decreases to minimize the interaction between the liquid and the other substance due to dissimilar IMFs, which reduces the energy of the fluid. Since surface tension is a measure of the energy of the liquid's surface area, the unit of surface tension is joules per square meter (J⋅m−2). A high value for surface tension indicates that the liquid has a high resistance to surface penetration, particularly when compared to the ability to penetrate (move through) the bulk of the liquid. 

Surface tension is based on the strength of IMFs alone (unlike viscosity, which has a shape component), leading to water having a higher surface tension compared to most other liquids due to the whole molecule being able to hydrogen bond. Another liquid with very high surface tension is mercury, which has metallic bonds holding the particles together, making it hard for objects to penetrate the surface of liquid mercury. The images on the right, which show a paperclip (top photo) and water strider (bottom photo) lying on the water's surface instead of penetrating, are good examples of water's high surface tension. Note that many bugs also have hairs covering their feet that have a waxy (nonpolar) coating, which also contributes to them being able to walk on water.

Figure 15: Molecular view of a water droplet showing how surface tension works. (Image: anonymous by request, license: CC BY-NC-SA 2.0.)

The image above depicts what happens on the molecular level to water molecules on the surface of a water droplet (left circle) and those in the bulk of the liquid (right circle). The water molecules in the bulk of the liquid are surrounded by other water molecules, so they have hydrogen bonding interactions all around them. Surface water molecules have air above (air molecules are nonpolar, and the strongest IMF between air and water is dipole-induced dipole) and water molecules below (hydrogen bonding). The strong intermolecular forces between the surface molecules and the bulk liquid and a strong desire to minimize contact with the nonpolar air molecules create a net pull downward. That downward pull is responsible for water's considerable surface tension. A way to think about surface tension is to imagine the surface water molecules acting like a tight film over the top of the water, preventing small objects from penetrating.

Figure 16: Celery sticks in dyed water showing how capillary action works in plants. (Image from PBS Kids for Parents

In the image above, water (and the dye within the water) is drawn up through the celery's vascular system, coloring the leafy vegetables in various colors. The thin, vein-like tubes in celery and other plants are called xylems, and they move water up from the plant's roots to the stem and leaves. The image provides an excellent visual of capillary action.

Capillary action is responsible for liquids flowing up a narrow tube (capillary), often overcoming other external forces, such as gravity. The definition of capillary action is the rise of liquids up a very narrow tube (the thinner the tube, the greater the rise of the liquid; see image to the right) and is an interplay between cohesive and adhesive forces.

wss-property-capillary-action.jpg

Figure 18: The glass test tube on the left is filled with water, while the glass test tube on the right is filled with liquid mercury. This is because water and glass have similar IMFs, which results in strong adhesive forces, while mercury and glass have dissimilar IMFs (weak adhesive forces but strong cohesive forces), resulting in the mercury pulling away from the glass surface. (Image: anonymous by request, license: CC BY-NC-SA 2.0.)

Cohesive forces (the tendency in liquids to resist separation) and adhesive forces (attractive forces between dissimilar molecules) are a consequence of IMFs, with cohesion resulting from IMFs binding molecules of a pure substance together, and adhesive forces arising from  IMFs binding different molecules together. For capillary action to occur, there need to be strong IMFs in the liquid and strong IMFS between the liquid and the capillary walls. For instance, water will go up the walls in a glass tube without difficulty due to the hydrogen bonding in water and the dipole-dipole interactions between the glass and the water (glass is made of SiO2, which is very polar), producing a concave meniscus (image on the left). In contrast, mercury in a glass tube will have a convex meniscus due to significant cohesive forces and comparatively weak adhesive forces (image on the left).

Identify the IMFs, relative viscosities, and relative surface tensions based on your knowledge of IMFs (Highlight text to see the answer).

IMFs                                      hydrogen bonding            dispersion                                dipole-dipole

                                                           dipole-dipole                                                                              dispersion
                                                              dispersion

Viscosity                                            2                                  3 (weakest)                               1 (strongest)

 

 

Surface tension                   1 (strongest)                   3 (weakest)                                             2

Phase Changes

The conversion from one state of matter to another is called a phase change or state change. The six types of phase changes, shown in the image below, are melting ↔ freezing (solid ↔ liquid), vaporization ↔ condensation (liquid ↔ gas), and sublimation ↔ deposition (solid ↔ gas). 

Because different states of matter have different associated energies, energy must be inputted or removed from the system to cause a phase change. For example, consider a solid versus a liquid: The particles in a solid are locked in place, held tightly together by IMFs, while in a liquid, the particles are free-flowing (a fluid state of matter), with weaker IMFs and more kinetic energy. Therefore, the kinetic energy needs to increase (energy inputted into the system – the solid) to convert from a solid to a liquid to allow for more movement and to overcome the IMFs that want to lock the molecules in place. When converting from liquid to gas, even more energy needs to be added to the system as there are no intermolecular forces in a volume of gas and significant amounts of kinetic energy.

Heating and Cooling curves.bmp

Heating Curve

Cooling Curve

The image above shows an example of a heating curve (top chart) and a cooling curve (bottom chart). On the curves, lines with a slope other than zero (the lines that are not flat) represent the energy increase or decrease as a substance changes temperature. Calculating the energy in these regions requires us to take advantage of the calorimetry equation we learned in chemistry 1; specifically, q = m·Cs·ΔT. In the equation, colloquially referred to as q equals mCAT, q is the heat in joules (J), m is mass in grams (g), Cs is the specific heat in J·g⁻¹·°C⁻¹, and ΔT is the change in temperature in K or °C. The lines on the curves with a zero slope signify the energy associated with a state change. The slopes are zero because the temperature does not change during a phase change, but the amount of energy does. These lines also represent the equilibrium between the two states. Essentially, any time a substance is at the temperature of its phase change (i.e., boiling point or melting point), the two phases are in a state of dynamic equilibrium. With dynamic equilibria, the rate of the forward process (or forward reaction) occurs at the same rate as the reverse process. Therefore, a liquid and a gas are in dynamic equilibrium at the boiling point of the liquid because evaporation and condensation co-occur at equivalent rates. When calculating the energy of a state change, you will multiply the enthalpy by the amount of substance that is changing state (remember that enthalpy is an extrinsic property and is therefore based on amount).

ΔHfusion = + 6.01 kJ/mol

ΔHvaporization = + 40.7 kJ/mol

Cs,ice = 2.09 J/(g·°C)

Cs,water = 4.18 J/(g·°C)

Cs,steam = 1.84 J/(g·°C).

Vaporization and Vapor Pressure

The image above shows an example of a heating curve (top chart) and a cooling curve (bottom chart). On the curves, lines with a slope other than zero (the lines that are not flat) represent the energy increase or decrease as a substance changes temperature. Calculating the energy in these regions requires us to take advantage of the calorimetry equation we learned in chemistry 1; specifically, q = m·Cs·ΔT. In the equation, colloquially referred to as q equals mCAT, q is the heat in joules (J), m is mass in grams (g), Cs is the specific heat in J/(g⁻¹·°C), and ΔT is the change in temperature in K or °C. The lines on the curves with a zero slope signify the energy associated with a state change. The slopes are zero because the temperature does not change during a phase change, but the amount of energy does. These lines also represent the equilibrium between the two states. Essentially, any time a substance is at the temperature of its phase change (i.e., boiling point or melting point), the two phases are in a state of dynamic equilibrium. With dynamic equilibria, the rate of the forward process (or forward reaction) occurs at the same rate as the reverse process. Therefore, a liquid and a gas are in dynamic equilibrium at the boiling point of the liquid because evaporation and condensation co-occur at equivalent rates. When calculating the energy of a state change, you will multiply the enthalpy by the amount of substance that is changing state (remember that enthalpy is an extrinsic property and is therefore based on amount).

References

  1. Petrucci, Ralph H., et al. General Chemistry: Principles and Modern Applications. Upper Saddle River, NJ: Prentice Hall, 2007.

  2. Brown, Theodore L., et al. Chemistry : the Central Science. New York: Pearson, 2017

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