Surveying the Chapter: Page 442 Chapter 11 Intermolecular Forces, Liquids, and Solids We begin with a brief comparison of solids, liquids, and gases from a molecular perspective, which reveals the important roles that temperature and play in determining the physical state of substance. We then examine the main kinds of intermolecular forces that occur within and between substances: o o o o We will learn that the nature and the strength of the intermolecular forces between molecules are largely responsible for many properties of liquids, including their, which is a measure of a liquid s resistance to flow, and, which is a measure of a liquid s resistance to increase its surface area. We ll explore, the transitions of matter between the gaseous, liquid, and solid states, and the energy changes that accompany phase changes. We ll examine the that exists between a liquid and its gaseous state and introduce the idea of. A liquid boils when its vapor pressure equals the pressure acting on the surface of the liquid. In a the equilibria among the gaseous, liquid, and solid phases are displayed graphically. Finally, we ll examine solids. Orderly arrangements of molecules or ions in three dimensions characterize. We ll examine how the structure of a crystalline solid can be conveyed in terms of its and how simple molecules and ions are most efficiently arranged in three dimensions. Solids can be characterized and classified according to the attractive forces between their component atoms, molecules, or ions. We examine four such classes: Section 11.1 A molecular Comparison of Gases, Liquids, and Solids The fundamental difference between states of matter is the between particles. 1
Gas: Total, much empty space; particles have complete freedom of motion particles far apart. Liquids:, particles or clusters of particles are free to move relative to each other; particles close together. Solids: ; particles are essentially in fixed positions particles close together. Because in the solid and liquid states particles are closer together, we refer to them as. The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: Section 11.2 Intermolecular Forces o The attractions molecules are not nearly as strong as the attractions that hold compounds together. o They are, however, strong enough to control physical properties such as o o o o o These intermolecular forces as a group are referred to as o van der Waals Forces A fourth type of force, are an important force in solutions of ions. The strength of these forces is what makes it possible for substances to dissolve in solvents. Dipole-Dipole Interactions Molecules that have are attracted to each other. The end of one is attracted to the end of the other and vice-versa. These forces are only important when the molecules are. The the molecule, the is its boiling point. 2
London Dispersion Forces o While the electrons in the 1s orbital of helium would each other (and, therefore, tend to stay far away from each other), it does happen that they wind up on the same side of the atom. o At that, then, the helium atom is, with an excess of electrons on the left side and a shortage on the right side. o Another helium nearby, then, would have a in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1. o London dispersion forces, or dispersion forces, are between an. o These forces are present in molecules, whether they are polar or nonpolar. o The tendency of an electron cloud to distort in this way is called. Factors Affecting London Forces The of the molecule affects the strength of dispersion forces: (like n-pentane tend to have dispersion forces than (like neopentane). This is due to the increased in n-pentane. The strength of dispersion forces tends to with. Larger atoms have larger electron clouds, which are easier to polarize. Which Have a Greater Effect: Dipole-Dipole Interactions or Dispersion Forces? If two molecules are of comparable size and shape, interactions will likely be the dominating force. If one molecule is much larger than another, will likely determine its physical properties. Boiling Points as a Function of Molecular Weight The BP of the group 4A (bottom) and 6A (top) hydrides are shown as a function of molecular weight. In general, the BP. How Do We Explain This? The nonpolar series (SnH4 to CH4) follow the expected trend. The polar series follows the trend from H2Te through H2S, but water is quite an anomaly. 3
Hydrogen Bonding The dipole-dipole interactions experienced when H is bonded to are unusually strong. We call these interactions. Hydrogen bonding arises in part from the. Also, when hydrogen is bonded to one of those very electronegative elements, the is exposed. Predicting the Types and Relative Strengths of Intermolecular Forces List the substances Barium Chloride Hydrogen molecule Carbon Monoxide Hydrogen fluoride Neon In order of increasing boiling points. 4
Practice Exercise A. Identify the intermolecular forces present in the following substances, and B. Select the substance with the highest boiling point: CH 3 CH 3 CH 3 OH CH 3 CH 2 OH Section 11.3 Some Properties of Liquids Intermolecular Forces Affect Many Physical Properties o The strength of the between particles can greatly affect the properties of a substance or solution. Viscosity o Resistance of a liquid to flow is called. o It is related to the ease with which molecules can move past each other. o Viscosity with stronger intermolecular forces and with higher temperature. Surface Tension o Surface tension results from the experienced by the molecules on the surface of a liquid. o These are forces. Section 11.4 Phase Changes (Up arrows are endo, down arrows are exo) 5
Energy Changes Associated with Changes of State o : Energy required to change a solid at its melting point to a liquid. o : Energy required to change a liquid at its boiling point to a gas. o Notice that the heat of vaporization is always larger than its heat of fusion. o The heat of fusion, or enthalpy of fusion, for ice is. o The heat of vaporization, or enthalpy of vaporization, for water is o The heat of sublimation is the sum of heats of vaporization and fusion. o For water = approx Energy Changes Associated with Changes of State The heat added to the system at the goes into pulling the molecules farther apart from each other. The temperature of the substance does not rise during the phase change. Calculating H for Temperature and Phase Changes Calculate the enthalpy change upon converting 1 mol of ice at -25 o C to water vapor (steam) at 125 o C under a constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.09 J/g-K, 4.18 J/g-K, and 1.84 J/g-K respectively. For H 2 O, Hfus = 6.01 kj/mol and Hvap = 40.67 kj/mol. 6
Practice Exercise What is the enthalpy change during the process in which 100 g of water at 50.0 o C is cooled to ice at -30 o C? A gas normally liquefies at some point when pressure is applied to it. The highest temperature at which a distinct liquid phase can form is referred to as the. The is the pressure required to bring about liquefaction at this critical temperature. Section 11.5 Vapor Pressure At any temperature, some molecules in a liquid have enough to escape. As the temperature rises, the fraction of molecules that have enough energy to escape. As more molecules escape the liquid, the pressure they exert increases. The liquid and vapor reach a state of : liquid molecules evaporate and vapor molecules condense. The of a liquid is the temperature at which its equals. The boiling point is the temperature at which its vapor pressure is 760 torr. Example Problem Use Figure 11.24 to estimate the boiling point of diethyl ether under an external pressure of 0.80 atm. At what external pressure will ethanol have a boiling point of 60oC? 7
Section 11.6 Phase Diagrams Phase diagrams display the state of a substance at various and and the places where exist between phases. The is the liquid-vapor interface. It starts at the the point at which all three states are in. It ends at the above this critical temperature and critical pressure the are indistinguishable from each other. Each point along this line is the of the substance at that pressure. The is the interface between liquid and solid. The at each pressure can be found along this line. the substance cannot exist in the liquid state. Along the the solid and gas phases are in ; the point at each pressure is along this line. Phase Diagram of Water Note the critical temperature and critical pressure: These are due to the strong forces between water molecules. The slope of the solid liquid line is. This means that as the pressure is increased at a temperature just below the, water goes from a solid to a liquid. 8
Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 at normal pressures. The low critical temperature and critical pressure for CO 2 make supercritical CO 2 a good solvent for extracting substances (such as caffeine). Section 11.7 Structures of Solids We can think of solids as falling into two groups: particles are in highly ordered arrangement. no particular order in the arrangement of particles. Attractions in Ionic Crystals In, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions. Because of the order in a crystal, we can focus on the repeating pattern of arrangement called the. 9
There are several types of basic arrangements in crystals, such as the ones shown above We can determine the of an ionic solid by determining how many ions of each element fall within the unit cell. Section 11.8 Bonding in Solids 10
Covalent-Network and Molecular Solids are an example of a covalent-network solid in which atoms are covalently bonded to each other. They tend to be hard and have high melting points. Covalent-Network and Molecular Solids is an example of a molecular solid in which atoms are held together with van der Waals forces. They tend to be softer and have lower melting points. Metallic Solids Metals are, but the attractions between atoms are too strong to be van der Waals forces. In metals, valence electrons are throughout the solid. 11