: The Liquid and Solid States 10-1 10.1 Changes of State How do solids, liquids and gases differ? Figure 10.4 10-2 10.1 Changes of State : transitions between physical states Vaporization/Condensation (l) (g) Freezing/Melting (also called fusion) (l) (s) Sublimation/Deposition (s) (g) 10-3 1
10.1 Changes of State: l g : occurs because some surface liquid molecules have high enough kinetic energy to escape to the gas phase. Heat is required to maintain the temperature needed for vaporation. Endothermic is vaporization at temperatures other than boiling points. The heat needed comes from the environment. : Transition from a gas to a liquid Occurs when gas particles cannot escape the container and thus, come into contact with a liquid An exothermic process (energy is released) The reverse 10-4 Copy right of Mc Graw-Hill vaporization Educ ation. Permis sion required for reproduction or dis play. 10.1 Changes of State: Equilibrium is a state in which opposing processes occur at equa rates. An equilibrium is designated by a double arrow, such as: vaporization liquid gas condensation The gas produced by evaporation exerts a pressure on the liquid below it. A liquid placed in a closed container will evaporate only until an equilibrium is reached. : the pressure exerted by the vapor at equilibrium with it s liquid. 10-5 10.1 Changes of State: Equilibrium Vapor pressure depends on molecular structure and temperature. 10-6 2
10.1 Changes of State: Equilibrium Boiling occurs when the vapor pressure equals the external pressure of the atmosphere. : The temperature at which vapor pressure equals the external pressure of the atmosphere. is exactly 1 atm. : What is the normal boiling point for propane? : the boiling point when the external pressure. What is its boiling point at 0.40 atm? 10-7 10.1 Changes of State: l s : the conversion of a liquid to a solid. The average kinetic energy of a liquid decreases when it is cooled. If the average kinetic energy becomes low enough, then the molecules become fixed in position in the solid state. Exothermic process substance. : the temperature at which this occurs for a : the freezing point when the external pressure is exactly 1 atm. : Phase change from solid to liquid Reverse of freezing Also called fusion Endothermic process : Same temperature as freezing point 10-8 10.1 Changes of State: l s When ice melts, the water molecules stay intact but flow past one another. When the ionic compound NaCl melts, ionic bonds are broken and ions flow freely. 10-9 3
10.1 Changes of State: s g : Under certain conditions, substances can go directly from the solid state to the gas state. CO 2 and I 2 do this under normal conditions. Endothermic : A transition directly from gas to solid without passing through the liquid state Reverse of sublimation Exothermic After iodine sublimes upon heating, it then undergoes deposition back to the solid state when it hits the cold tube above it. 10-10 : Changes in the State of Matter Identify the process shown in the following diagram: 10-11 10.1 Changes of State : shows the phase changes of a substance as heat is removed or added at constant pressure Figure 10.14 10-12 4
10.1 Changes of State: Cooling Curve Notice that only one change occurs in a heating/cooling curve at a time: either phase is changing or temperature is changing Figure 10.14 A 10-13 : Cooling & Heating Curves The molecular-level diagram represents two physical states of a substance. Identify where each state would predominate on the heating curve given to the right. 10-14 10.1 Changes of State: Energy Changes : With constant phase, energy change can be calculated with: : At constant temperature, energy change can be calculated with: a substance : Energy required to melt 1 mole of : Energy required to evaporate 1 mole of a substance Units are J/mol or kj/mol 10-15 5
: Energy for Phase Changes Calculate the heat required to boil 125 g of water at its boiling point, 100.0 C. The molar heat of vaporization of water is 4.07 10 4 J/mol. 10-16 10.1 Changes of State: Energy Changes For the reverse processes, the amount of heat is the same, but heat is released instead of absorbed, so the values are the same but with opposite sign. For water: ΔH f = 6.01 kj/mol What is the heat of solidification (freezing)? ΔH s = -6.01 kj/mol ΔH V = 40.7 kj/mol What is the heat of condensation (gas to liquid)? ΔH c = -40.7 kj/mol 10-17 10.2 Intermolecular Forces Why is CO 2 a gas at room temperature and H 2 O a liquid? Why do liquids have different boiling points? What causes molecules to stick together when in the liquid state? : Attractive forces existing between molecules, holding them into condensed phases (l & s) Three types of intermolecular forces: - Dipole-dipole forces - Hydrogen bonding - London dispersion forces 10-18 6
10.2 Intermolecular Forces molecules. : the attractive forces felt between polar A polar molecule has a positive and negative end, so it is considered a dipole. The δ+ of one molecule and the δ- of another molecule will attract one another. SO 2 is polar so it experiences dipole-dipole forces in the liquid state. 10-19 : Dipole-Dipole Forces Which of the following molecules experience dipole-dipole forces SO 2 CH 2 Cl 2 PCl 3 BF3 Which of the following molecules experience dipole-dipole forces? 1. SCl 2 2. CO 3. NH 3 4. CCl 4 10-20 : Dipole-Dipole Forces 1. Which is more polar? HCl or HBr? 2. Which has the stronger dipole-dipole forces? 10-21 7
10.2 Intermolecular Forces The irregularity in boiling point trends in this figure is due to a particularly strong dipole-dipole force called hydrogen bonding. 10-22 10.2 Intermolecular Forces : occurs between molecules that contain hydrogen covalently bonded to either nitrogen, oxygen, or fluorine. N-H, O-H, or F-H Hydrogen bonding causes boiling points to be much higher than expected. An important force in living systems, where it stabilizes molecular Copyshapes 10-23 right Mc Graw-Hill Educ ation. Permis sion required for reproduction or dis play. 10.2 Intermolecular Forces Hydrogen bonding is the strong force that holds together the DNA double helix. 10-24 8
10.2 Intermolecular Forces Hydrogen bonding in H 2 O causes ice to be less dense than the liquid state. Ice being less dense than liquid water, causes it to expand as freezing, resulting in a great deal of damage caused by ice. 10-25 : Hydrogen Bonding Which of the following experience hydrogen-bonding? 1. CH 3 NH 2 2. NF 3 3. CH 3 -O-CH 3 4. CH 3 CH 2 OH 5. HCl 10-26 10.2 Intermolecular Forces : are a result of temporary electron cloud distortions and temporary polarity (dipoles). exist in all substances. most important in large molecules is the only intermolecular force present in nonpolar substances. 10-27 9
: London Dispersion Forces 1. Which has stronger London dispersion forces, CH 4 or SiH 4? 2. For each pair, determine which has the stronger London dispersion forces. He or Kr HCl or HBr CH 4 or C 2 H 6 10-28 10.2 Intermolecular Forces Relative boiling points tell us about intermolecular force strength. The stronger the forces, the higher the temperature needed to overcome them. 10-29 10.2 Intermolecular Forces The larger the molecule, the more London Dispersion forces, the higher the boiling point. 10-30 10
Summary of Intermolecular Forces 10-31 10.2 Intermolecular Forces : To determine what has the stronger total intermolecular forces o the higher boiling point: Similar molar masses? Consider strength of polarity - H-bonding > dipole-dipole > London dispersion forces - Example: HCl > CO > N 2 Very different molar masses? London dispersion forces are more important - Large MM > smaller MM - Example: HI (s) > HBr (l) > HCl (g) 10-32 : Trends in Intermolecular Forces Consider the substances CO and HF. 1. Which has the stronger intermolecular forces 2. Which has the higher melting point & boiling point 10-33 11
: Trends in Boiling Point Which of the following should have the highest boiling point? 1. CH 3 OH or CH 3 SH 2. I 2 or SO 2 3. CO or N 2 10-34 10.3 Properties of Liquids Properties of liquids are related to their intermolecular forces. Particles in a liquid are much closer together than the particles in a gas. Liquid particles can slide past each other and are not fixed, as they are in a solid. 10-35 10.3 Properties of Liquids : How is vapor pressure related to intermolecular force strength? The stronger the intermolecular forces: the more strongly the molecules are held in the liquid phase the lower the vapor pressure 10-36 12
: Trends in Vapor Pressure Which of the following should have the highest vapor pressure at a give temperature? 1. CH 3 OH or CH 3 SH 2. I 2 or SO 2 3. CO or N 2 10-37 10.3 Properties of Liquids : resistance to flow which is depends on molecular size and intermolecular force strength. Large molecules get tangled in each other to increase viscosity. Viscosity decreases as the temperature increases. 10-38 10.3 Properties of Liquids : amount of work required to increase the surface area of a liquid by a unit amount. The greater the IMF in a liquid, the greater the surface tension. Surface tension decreases as temperature increases. Surface molecules have fewer molecules to interact with than the inner molecules. 10-39 13
10.3 Properties of Liquids : a curved surface of a liquid produced by intermolecular forces. A concave surface occurs when the IMF between the liquid and the glass are greater than the IMF among the liquid molecules. A convex surface occurs when the IMF between the liquid and the glass are less than the IMF among the liquid molecules. 10-40 10.4 Properties of Solids The particles in the solid state are held together by strong intermolecular forces in specific positions No translational motion. Vibrational and rotational motion Crystalline solids are well ordered. Amorphous solids lack regular form. 10-41 10.4 Properties of Solids : Occurs when the temperature of a liquid drops rapidly, resulting in the particles solidifying in a partially disordered state. Particles are randomly arranged. Lacks regular form : Occurs when a liquid solidifies slowly, allowing the array of particles to become well ordered. Most solids are of this type. Results in the symmetrical arrangement of atoms/ions called a crystal lattice. Atoms, molecules, or ions pack together efficiently to maximize attractive forces and bonding interactions. 10-42 14
10.4 Properties of Solids Metallic Pure metals and alloys; composed of atoms Metallic bonding Ionic Ionic compounds; composed of ions Ionic bonding Molecular Molecular compounds or nonmetal elements; composed of molecules or atoms Connected by intermolecular forces Network Molecular compounds or nonmetal elements All atoms are connected by covalent bonds 10-43 10.4 Properties of Solids: Metals Are ductile and malleable Valence electrons move freely through all parts of a metal. Attractions between atoms of a metal are delocalized, and therefore it is easy to move atoms by applying force. Sea of electrons model 10-44 10.4 Properties of Solids: Metals : homogeneous solutions of a metal mixed with one or more additional metallic or nonmetallic elements. Properties differ from those of their parent elements replace some of the metal atoms with another element. - Stainless steel have some smaller atoms in the space between the metal atoms. - Cast iron 10-45 15
10.4 Properties of Solids: Ionic : contain cations and anions arranged in crystalline solids. Electrostatic forces (ionic bonds) hold together ionic crystals. High melting points Hard and brittle Solids are not electrical conductors, but liquids are. 10-46 10.4 Properties of Solids: Molecular : composed of molecules held together by intermolecular forces. Nonpolar solids are held together by London dispersion forces. Form soft crystals with low melting points Electrical insulators (no ions) Poor conductors of heat Polar solids are held together by London dispersion forces and dipole-dipole forces. Typically harder than nonpolar solids Low to moderate melting points Electrical insulators (no ions) Dry ice, CO 2 10-47 10.4 Properties of Solids Network : consists of a giant molecule that forms the entire crystal. Formed by metalloids or carbon Strong covalent bonds connect the atoms in a network solid Non-conductors of electricity High melting points Very hard Have very stable, three dimensional structures Many have a diamond structure, or some derivative of it. 10-48 16
: Types of Solids Predict the type of solid for each: 1. CO 2 sublimes at -78.4ºC 2. Br 2 mp = -7.2ºC 3. C (diamond) mp = 3550ºC 4. MgO mp = 2800ºC 5. Identify the type of solid shown by the following molecularlevel image. 6. What types of forces hold the particles together? 10-49 10.4 Properties of Solids 10-50 : Types of Crystalline Solids A particular crystalline solid is very hard, does not conduct electricity when solid or when melted, and has a high melting point. What type of solid is it? 10-51 17