Honors Unit 9: Liquids and Solids

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1 Name: Honors Unit 9: Liquids and Solids Objectives: 1. Students will be able to describe particles in the solid, liquid, and gas phases, and to explain what happens during phase transitions in terms of the kinetic energy of particles. 2. Students will learn the meaning of dynamic equilibrium as it relates to liquid vapor equilibria, vapor pressure curves, and boiling point variations. 3. Students will be able to identify the segments of heating and cooling curves. 4. Students will be able to predict the type of intermolecular forces (dispersion forces, dipole-dipole forces, or hydrogen bonding) present in a given sample by looking at a formula for the substance. 5. Students will be familiar with the unusual properties of water due to its molecular structure and intermolecular forces. 6. Students will be able to identify and characterize four types of solids: network covalent solids, metals, molecular solids, and ionic solids. Introductory Vocabulary Comparison of States of Matter State Shape Volume Why? Solid Liquid Gas Particles close together; strong intermolecular forces Particles still close; reasonable amount of intermolecular forces Particles are far apart; intermolecular forces are negligible Why gas laws and not solid and/or liquid laws? Gases are mainly & they have attractions between molecules Solids/liquids have particles which are and have more varied forces between particles. These forces between particles ( ) play a major role in the behavior of liquids & solids, whereas they are negligible in gases!

2 2 Intermolecular Forces Intermolecular Forces (IMF) Bonds Between Molecules o Inter- prefix = o Short range forces between in a sample. o There are 3 main types of IMF: (London) Dispersion Forces Dipole-Dipole Forces Hydrogen Bonding ***All three intermolecular forces are relative to the strength of a covalent bond (an intramolecular force)!! 1) London Dispersion Forces IMF between two formed by temporary positive and negative attractions due to the shifting of electron cloud. o intermolecular force they re temporary! Found in, but become important when they are the only IMF present. Strength as molar mass (and the number of ) increases. o Higher boiling point = Example #1: Account for the fact that chlorine is a gas, bromine is a volatile liquid, and iodine is a volatile solid at room temperature.

3 3 2) Dipole-Dipole Forces Attractions between oppositely charged regions of o Caused by attraction of δ + ( ) for δ - ( ) o Present in all substances!! Compounds with dipole-dipole forces have higher o The molecules stick together and require much more energy for the phase change Solubility: like dissolves like 3) Hydrogen Bonding A special type of force o Unusually strong intermolecular force o Must have an attached directly to an in Lewis structure Remember: Hydrogen bonding is FON!! o The δ + H from one molecule is strongly attracted to the negative end of the dipole of another Hydrogen bonding melting & boiling points because energy is required to break the forces between molecules. H-bonding is especially strong in water because: o The O H bond is very polar o There are 2 lone pairs of electrons on the O atom o Accounts for many of water s unique properties such as: high specific heat capacity, why water increases volume when freezing, high surface tension

4 4 Example #2: (a) What types of intermolecular forces are present in the following substances? N2 HF SiCl4 CH3Cl NH3 (b) Rank the substances above in order of increasing boiling point. Relative Strength of Intermolecular Forces: Phase Changes Definition: o All phase changes involve energy (enthalpy, ΔH) A. Phase changes that require energy ( )

5 Temperature ( C) 5 B. Phase changes that release energy ( ) Heating Curve D E C A B Time (minutes) Sample Questions for Heating Curves: 1. Where does evaporation occur? What is the melting point of this sample? Where is there only a liquid present? Where would the molecules have the most kinetic energy? At what time does boiling begin? When does melting end? Where would freezing happen? 7. ***Note that cooling curves are just the opposite of a heating curve.

6 6 Heat and Change of State Calculations Melting/Freezing (add or subtract heat = +/-) o Heat of Fusion (ΔHfus) Vaporization/Condensation (add or subtract heat = +/-) o Heat of Vaporization (ΔHvap) o Sublimation/Deposition phase change between solid and gas Heating Curve Equations: ***Note that temperature does not change during a phase change!! Heating & Phase Change Constants for Water: ΔHvap,water = kj/mole ΔHfus, water = 6.01 kj/mole Cice = 2.09 J/g C Csteam = 1.84 J/g C Cwater = 4.18 J/g C Example #3: How many kilojoules of heat are required to completely vaporize g of ethanol (C2H5OH)? The heat of vaporization is 43.3 kj/mol.

7 7 Example #4: How much energy in kj is required to heat g of liquid water from zero to 100 C, and then vaporize all of it? Example #5: How much energy is required to heat 75.0 grams of ice from 0.0 C to C?

8 8 Liquid Vapor Equilibrium = pressure due to the force of gas particles above a liquid colliding with the walls of a container o A: B: Equilibrium = two opposing processes occur at the same rate. Dynamic Equilibrium w/ Vapor Pressure o The liquid level in the container when a system is at equilibrium!! Molecules are constantly moving between phases so no net change Vapor Pressure Curves What does the graph at the right tell you about the relationship between vapor pressure and temperature?

9 9 Vapor Pressure Curves (cont.) As Temperature, vapor pressure. As attractive (intermolecular) forces between molecules, vapor pressure. Liquids with higher vapor pressures at a given T are more. Vapor Pressure Curve Questions: (see pg. 11 in reference book for curve) 1. What does this graph tell you about the IMF between molecules in substances a e? 2. Which substance is most volatile? 3. At a specific temperature, where would the boiling point of a substance be? A liquid boils when o At this point, the liquid molecules overcome atmospheric pressure and jump into the gas phase. Normal Boiling Point = o As pressure increases, boiling point increases; as pressure decreases, boiling point decreases.

10 10 Properties of Liquids o Viscosity The particles in a liquid are close enough together that their attractive forces slow their movement as they flow past one another The the attractive forces (intermolecular forces), the the liquid is. As temperature increases, viscosity. o Surface Tension A measure of the inward pull by particles in the interior. The stronger the IMF, the the surface tension. o Surfactant any substance that interferes with the hydrogen bonding between water molecules and reduces surface tension Types of Solids Solids can be classified into many different classes. We are going to focus on a few: Crystalline solids o Network Covalent o Metallic o Molecular o Ionic Amorphous solids ****See the solids project handouts for more information!!

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