Chapter 13: Intermolecular Forces, Liquids, and Solids

Transcription

1 Chapter 13: Intermolecular Forces, Liquids, and Solids The Kinetic-Molecular Theory of Matter Recall the kinetic-molecular theory of gases (Chapter 12). Kinetic-molecular theory applies to solids and liquids too: 1. Molecules are in constant motion in random directions. 2. Molecules move at different speeds, but the average speed is proportional to temperature. All molecules have the same average kinetic energy at a given temperature. In gases, the distances between molecules are much larger than the size of the molecules themselves. As such, most of the motion of gas molecules is (moving from one location to another). In solids, molecules are packed in a three-dimensional solid lattice. They cannot easily change location, so their kinetic energy is primarily due to motion. This motion is averaged about the molecule s location in the lattice. In liquids, molecules are much closer together than in gases, but they are still able to change locations (unlike in solids). As such, liquid molecules undergo significant amounts of translational, rotational and vibrational motion.

2 Intermolecular Forces The kinetic energy of molecules tends to keep them as separated and disorganized as possible, so there must be some opposing force(s) pulling them together particularly in solids and liquids. These attractions are intermolecular forces. Intermolecular forces are similar to, but weaker than, those involved in bonding. There are four kinds of intermolecular forces, all of which are based on. The strength of an intermolecular force is determined by the strength and permanence of the dipoles involved (and by the distance between the dipoles). Type of Intermolecular Force Approximate Strength (kj/mol) Ion-Dipole Dipole-Dipole (includes hydrogen bonding) Dipole-Induced Dipole Induced Dipole-Induced Dipole (aka London dispersion forces)

3 Hydrogen bonding is a special case of dipole-dipole attraction. This is due to the very large dipole moment that occurs when a hydrogen atom is bonded to a very small, very electronegative atom (i.e. F, or N). The large dipole moment results in a particularly strong dipole-dipole attraction. In hydrogen bonding, both the atom bonded to H and the atom attracted to H must be either F, or N. They do not, however, have to be the same element. e.g. Induced dipoles occur when two molecules approach closely enough that the electron cloud of one is repelled by the electron cloud of the other. The electrons move to the far end of the molecule, setting up a temporary dipole. These induced dipoles attract each other (and one induced dipole can induce another to form). Large atoms are most susceptible to forming induced dipoles. (most polarizable) Why? Intermolecular forces between neutral molecules are referred to as a group as van der Waals forces.

4 Properties of Liquids Liquids are much denser than gases. As in solids, there is virtually no free volume between molecules in a liquid. As such, they are virtually incompressible a property used to advantage in hydraulics. The molecules in a liquid are in constant motion translational, vibrational and rotational. As in gases, the motion of any one molecule is random, but the average speed is proportional to temperature. The statistical picture for the energy distribution in liquids is identical to that for ideal gases: Enthalpy of Vaporization When you leave wet dishes on a draining board overnight, they will usually be dry the next morning. The water has evaporated even though your kitchen was C below the boiling point of water! How is this possible? At any given temperature, some of the liquid molecules will have enough energy to escape the intermolecular forces holding the molecules together.

5 Heating the liquid increases the rate of vaporization by increasing the proportion of molecules having at least this much energy. Since heating increases the rate of vaporization, this is an process with a enthalpy. This is consistent with the observation that rapid vaporization cools any remaining liquid. The molecules with enough energy to evaporate do so, reducing the average kinetic energy of the remaining liquid molecules (thereby reducing the liquid s temperature). During slow evaporation, heat is steadily resupplied from the environment so we don t notice this loss. The enthalpy of vaporization is defined as the energy required to vaporize a liquid. Its opposite, the enthalpy of condensation, is of equal magnitude but opposite sign. e.g. H 2 (l) H 2 (g) H vap = 40.7 kj/mol H 2 (g) H 2 (l) H cond = 40.7 kj/mol Vapour Pressure Because there will always be at least a few molecules in a liquid with enough energy to evaporate, there will always be a layer of gas molecules immediately above a liquid. The pressure exerted by these gas molecules is referred to as vapour pressure. When a liquid is placed in a sealed container, it will eventually reach an equilibrium at which the rate of evaporation is the same as the rate of condensation. In this situation, an equilibrium vapour pressure (often shortened to vapour pressure ) can be measured. This is a physical property specific to a given liquid at a given temperature, and it increases exponentially with temperature.

6 When the vapour pressure of a liquid is equal to the atmospheric pressure, the liquid begins to boil. The pressure exerted by gas molecules is equal to the pressure of the atmosphere pushing down on the liquid s surface. As such, bubbles of gas can form within the liquid (not just at its surface), they rise, and we observe boiling. The temperature at which a liquid boils is referred to as its boiling point; however, this property is pressure-dependent. Reducing the pressure will decrease a liquid s observed boiling point while increasing pressure will increase the observed boiling point. To allow us to compare boiling points of different compounds, we therefore define the temperature at which the vapour pressure of a liquid is equal to as a liquid s normal boiling point. Liquids with a high vapour pressure/low boiling point are referred to as. e.g. Looking at the vapour pressure curves above, (a) What is the approximate vapour pressure of ethanol at room temperature (20 C)? (b) Are liquid ethanol and its vapour in equilibrium when the temperature is 60 C and the vapour pressure is 600 mmhg? If not, does liquid evaporate to form more vapour, or does vapour condense to form more liquid.

7 Plotting vapour pressure vs. temperature gives an exponential curve (see previous page). Plotting the natural logarithm of the vapour pressure versus the reciprocal of temperature gives a linear graph. The equation of the line is the Clausius-Clapeyron equation: H ln P = - o vap + C RT This equation is more useful in its comparative form (i.e. the equation for calculating the slope of the linear graph): P ln 1 H = - o vap P 2 R T 1 T 2 W here P 1 and P 2 are vapour pressures (in the same units as each other), H vap is the enthalpy of vaporization (in J/mol), R is the ideal gas constant (in J mol -1 K -1 ) and T 1 and T 2 are temperatures (in K). e.g. If diethyl ether has a vapour pressure of 534 mmhg at C and a vapour pressure of 57.0 mmhg at C, what is its enthalpy of vaporization?

8 Critical Temperature and Pressure As the temperature of a liquid is increased, its vapour pressure continues to increase, and it is necessary to have an increasingly high atmospheric pressure in order to maintain the liquid state. There comes a temperature at which no amount of pressure is able to compress gas to liquid; the energy of the molecules is just too high. This is the critical temperature. The point on a phase diagram at which the critical temperature and critical pressure meet is referred to as the critical point. At temperatures higher than the critical temperature, high pressures give dense gases called supercritical fluids. The molecules in supercritical fluids are packed together almost as tightly as in liquids but they have enough energy to overcome intermolecular forces and move freely. As such, they can make excellent solvents (especially supercritical C 2 which has the added advantage of being much more environmentally friendly than most organic solvents). When is it important to know a substance s critical point? e.g. Choosing a refrigerant. Refrigerators work by compressing a gas to liquid. As the liquid is released into areas of lower pressure, it evaporates, consuming energy and cooling the fridge. If a potential refrigerant has a very low critical temperature, this process will not work because it will be impossible to condense the gas to a liquid. Freon-12, a common refrigerant, has a critical temperature of ~112 C.

9 Surface Tension Molecules within a liquid are attracted to all of the molecules around them via intermolecular forces. Molecules on the surface of a liquid can only interact with the molecules immediately below them and beside them. As such, the surface molecules feel a net attraction toward the interior of the liquid and act as a skin for the liquid. The energy required to break through this surface skin is referred to as surface tension. H H H H H H H H H H H H H H smooth surface; lots of intermolecular attractions H H H H H H H H H H H H H H surface tension broken; intermolecular attractions interfered with Some insects take advantage of surface tension to literally walk on water! Surface tension is also the reason why liquids tend to bead spheres have the smallest surface area for a given volume. Capillary Action Capillary action is due to the same intermolecular forces as surface tension; however, the attractions are between molecules of a polar liquid (often water) and a polar solid (often glass). Glass consists of polar Si- bonds so, if a narrow glass tube is placed in water, the water molecules are attracted to the glass. ther water molecules are attracted to the initially stuck molecules, and are pulled H H Si H H H H Si H H H H Si H H H H simplified structure for glass (it's much more 3-dimensional!)

10 up by surface tension. This continues until the force of gravity pulling water molecules down is equal to the force of surface tension pulling water molecules up. Capillary action is responsible for the meniscus typically observed when using pipettes, burettes, etc.. Viscosity Viscosity is defined as the resistance of liquids to flow. Liquids with high viscosities tend to be thick and difficult to pour. This is because the strong intermolecular forces in viscous liquids cause the molecules to stick together. Viscous liquids tend to fall into one of two categories (or both): 1. large molecules 2. polar molecules Summary Exercises 1. butane methanol helium (CH 3 CH 2 CH 2 CH 3 ) (CH 3 H) (He) (a) Rank the pure substances above in order of increasing strength of intermolecular forces: (b) Which of these pure substances would you expect to be gases at 25 C and 1 atm?

11 2. The graph below shows vapour pressure curves for carbon disulfide (CS 2 ) and nitromethane (CH 3 N 2 ). (a) What type of intermolecular forces exist in the liquid phase of each compound? (b) What are the vapour pressures of CS 2 and CH 3 N 2 at 40 C? (c) What are the normal boiling points of CS 2 and CH 3 N 2? (d) At what temperature does CS 2 have a vapour pressure of 600 mmhg? (e) At what temperature does CH 3 N 2 have a vapour pressure of 60 mmhg?

12 Properties of Solids As in liquids, the molecules in solids are packed tightly together. Unlike in liquids, the molecules in solids do not experience translational motion only vibrational. Solid molecules are held tightly in place by strong intermolecular forces. There are two main categories of solids: Crystalline solids, in which the molecules are ordered, and Amorphous solids, in which the molecules are disordered. Crystalline Solids Crystalline solids are made up of repeating units built upon each other like bricks in a wall. An individual brick is called the. Mathematicians have determined that there are seven possible shapes for a unit cell (aka seven crystal systems). These include cubic unit cells and hexagonal unit cells: Crystalline solids can be analyzed by x-ray crystallography, in which an x-ray is passed through a crystal. The crystal acts as a diffraction grating (the x-rays can pass through gaps in the crystal structure but not through the atoms themselves), and analysis of the resulting diffraction pattern allows a chemist to determine the structure of the crystal.

13 There are four different kinds of crystalline solids, each distinguished by the type of chemical bonding involved: Ionic solids: positive and negative ions held together by electrostatic attractions Metallic solids: metal atoms held together by electrostatic attractions between nuclei and a sea of electrons Molecular solids: molecules held together by van der Waals attractions Network solids: covalently bonded networks of atoms For a more detailed look at each type of solid, see (and learn!) Table 13.6 in Kotz. Amorphous Solids The molecules in amorphous solids have no long-range order and can be thought of as frozen liquids. The most widely known amorphous solid is glass. Interestingly enough, crystallization results in the failure of glass. Melting Points and Enthalpies of Fusion Just like liquids, the physical properties of solids can be explained using intermolecular forces. Converting a solid to a liquid requires input of enough energy to overcome the intermolecular forces holding each atom/molecule in the solid in place. This process can be called either melting or fusion. Thus, melting points (or fusion points) are higher for solids with stronger intermolecular forces. In general, ionic solids and transition metals have the highest melting points; polar molecular solids have intermediate melting points; and nonpolar molecular solids have low melting points.

14 The enthalpy of fusion is defined as the energy required to liquefy a solid. Its opposite, the enthalpy of crystallization, is of equal magnitude but opposite sign. e.g. H 2 (s) H 2 (l) H fus = 6.02 kj/mol H 2 (l) H 2 (s) H cryst = 6.02 kj/mol Solids with high enthalpies of fusion have high melting points (and vice versa). Sublimation Some solids convert directly from the solid to the gas phase without passing through the liquid stage. This is called sublimation. The most familiar examples are probably carbon dioxide ( dry ice ) and iodine. The enthalpy of sublimation is defined as the energy required to sublime a solid. Phase Diagrams The equilibrium vapour pressure curves we looked at earlier this chapter can be extended to describe the complete range of states of matter. This is a phase diagram, and each state of matter is referred to as a. The phase diagram to the right is that for carbon dioxide (C 2 ). Note that the slope of the solid-liquid boundary line is positive, indicating that the liquid is less dense than the solid. Also, note the point where the three phases (gas, liquid and solid) meet. This is the.

15 Below the triple point, it is not possible for a substance to exist in the liquid phase. For C 2, this means that liquid C 2 cannot exist at pressures below 5 atm or temperatures below -57 C. According to its phase diagram, the critical temperature for C 2 is and the critical pressure is. The phase diagram to the right is that for water. Note the unusual negative slope of the solid-liquid boundary line, indicating that the solid is less dense than the liquid. This unusual property is due to the strongly hydrogen bonded structure of the solid phase of water. This relatively low density of ice (compared to liquid water) is what allows fish to survive in partially frozen lakes and rivers over the winter. If ice was more dense than water, lakes would freeze from the bottom up, and there would be no liquid water left for the fish. Because ice is less dense than water, lakes essentially freeze from the top down, so there is often a layer of liquid water left at the bottom.

16 Important Concepts from Chapter 13 kinetic-molecular theory of matter intermolecular forces (four categories, hydrogen bonding) properties of liquids o vapour pressure o boiling vs. evaporation o enthalpy of vaporization (Clausius-Clapeyron equation) o surface tension o capillary action o viscosity properties of solids o five different types of solids (see Table 13.6 in Kotz) o enthalpy of fusion o enthalpy of sublimation phase diagrams o predicting changes of state under different conditions o critical point (and supercritical fluids) o triple point