Chapter 14 “LIQUIDS AND SOLIDS” What are the properties of the
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Chapter 14 “LIQUIDS AND SOLIDS” What are the properties of the ‘condensed states’ of matter? T H Witherup 02/06 Honors (rev 07)
Ch. 14 OBJECTIVES Show how the Kinetic-Molecular (K-M) Theory accounts for the physical properties of liquids & solids. Describe different types of intermolecular forces, and how they affect properties of liquids & solids. Learn about viscosity & surface tension, and explain their relationship to intermolecular forces. Compare crystalline & amorphous substances. Relate the structure & bonding in the four types of crystalline solids to the properties they exhibit. Describe the changes of state (vaporization, condensation, boiling, sublimation, deposition, melting, freezing). Learn how to interpret “Heating/Cooling Curves” and “Phase Diagrams.”
14-1 Condensed States of Matter: Liquids and Solids Condensed matter has much higher density (mass/volume) than gases. Unlike gases, particles of condensed matter experience different amounts and types of attractive forces. Kinetic-Molecular Theory can help explain the properties of condensed matter. The state of a substance at room temperature depends on the attractive forces between its particles.
Comparing the States of Matter Gas Liquid Total disorder Particles free to move past each other Particles far apart Disorder Particles free to move past each other Particles close together Solid Ordered arrangement Particles vibrate, but remain in a fixed position Particles close together -
Comparing Properties of Gases, Liquids & Solids PROPERTY COMPRESSIBILITY DENSITY VOLUME SHAPE DIFFUSION EXPANSION (with Heat) GAS LIQUID SOLID
Comparing Properties of Gases, Liquids & Solids PROPERTY GAS LIQUID SOLID COMPRESSIBILITY DENSITY High Slight Low VOLUME SHAPE DIFFUSION EXPANSION (with Heat)
Comparing Properties of Gases, Liquids & Solids PROPERTY GAS LIQUID SOLID COMPRESSIBILITY DENSITY High Slight Low Low High High VOLUME SHAPE DIFFUSION EXPANSION (with Heat)
Comparing Properties of Gases, Liquids & Solids PROPERTY GAS LIQUID SOLID COMPRESSIBILITY DENSITY High Slight Low Low High High Fills container Definite Rigid VOLUME SHAPE DIFFUSION EXPANSION (with Heat)
Comparing Properties of Gases, Liquids & Solids PROPERTY GAS LIQUID SOLID COMPRESSIBILITY DENSITY High Slight Low Low High High Fills container Definite Rigid Of container Of container Fixed VOLUME SHAPE DIFFUSION EXPANSION (with Heat)
Comparing Properties of Gases, Liquids & Solids PROPERTY GAS LIQUID SOLID COMPRESSIBILITY DENSITY High Slight Low Low High High Fills container Definite Rigid Of container Of container Fixed Rapid Slow Very slow (@ surface) VOLUME SHAPE DIFFUSION EXPANSION (with Heat)
Comparing Properties of Gases, Liquids & Solids PROPERTY GAS LIQUID SOLID COMPRESSIBILITY DENSITY High Slight Low Low High High Fills container Definite Rigid Of container Of container Fixed DIFFUSION Rapid Slow EXPANSION (with Heat) High Low Very slow (@ surface) Low VOLUME SHAPE
Review of Chemical Bonding Ionic: transfer of electrons from a metal to a nonmetal. Metallic: sharing of valence electrons of the metal atoms. Resulting ions have opposite charge and are attracted to each other, forming an ionic bond. Results in a network of positive ions in a “sea of electrons.” Covalent: strong intramolecular forces from the sharing of valence electrons between atoms. Results in individual molecules with specific shapes. Intermolecular forces exist between molecules.
Intermolecular Forces have a wide range of strengths. They are much weaker than ionic, covalent and metallic bonds. Basic types of Intermolecular Forces: Dispersion Forces attraction between temporary ‘induced dipoles.’ Consider noble gas boiling points and Fig. 14-8.
Dispersion Forces: Boiling Points of Noble Gases ( C) 0 Boiling Point (deg C) -50 0 20 40 -300 100 -107 Xe -150 -250 80 -67 Rn -100 -200 60 -153 Kr -186 Ar -246 Ne -269 He Atomic Number What causes the boiling point of helium to be so low and that of radon to be so high? Consider the effects shown in Figure 14-8 of text. (Data from Figure 14-7, page 462 of your textbook.)
Intermolecular Forces have a wide range of strengths. They are much weaker than ionic, covalent and metallic bonds. Basic types of Intermolecular Forces: Dispersion Forces attraction between temporary ‘induced dipoles.’ Consider noble gas boiling points and Fig. 14-8. Dipole-Dipole attraction between polar molecules (dipoles) having permanent charge separation. Consider HCl, HBr, CO2 (linear) and SO2 (bent).
Dipole-Dipole Forces: Polar Molecules Consider HCl (a dipole or polar molecule): δ δ- H Cl δ :::::::::: H :::::::::: δCl Compare CO2 and SO2 (covalent molecules with polar bonds): Polar or Nonpolar? Why? B.P. -78 C O C O B.P. -10 C S O O
Intermolecular Forces have a wide range of strengths. They are much weaker than ionic, covalent and metallic bonds. Basic types of Intermolecular Forces: Dispersion Forces attraction between temporary ‘induced dipoles.’ Consider noble gas boiling points and Fig. 14-8. Dipole-Dipole attraction between polar molecules (dipoles) having permanent charge separation. Consider HCl, HBr, CO2 (linear) and SO2 (bent). Hydrogen Bonds strong attraction between H atom of a molecule and a very electronegative atom (F,O,N) of another molecule. See data on the next slides.
Consider Boiling Points As mass increases, we expect Boiling Point to rise. CH4, SiH4, GeH4, SnH4 NH3, PH3, AsH3, SbH3 Observed: HCl HBr HI HF (anomaly) H2O, H2S, H2Se, H2Te Observed: PH3 AsH3 SbH3 NH3 (anomaly) HF, HCl, HBr, HI Observed: CH4 SiH4 GeH4 SnH4 (follow trend) Observed: H2S H2Se H2Te H2O (anomaly) See Table on next slide.
Consider Boiling Points (kelvins) Formula XH4 XH3 XH2 XH CH4 NH3 H2O Period 3 (93) SiH4 (243) PH3 (373) H2S HF (303) Period 4 (163) GeH4 (183) AsH3 (213) H2Se Period 5 (193) SnH4 (203) SbH3 (243) H2Te (223) (243) (273) Period Period 2 HCl (193) HBr (213) HI (253) What causes these anomalies (shown in yellow)? HYDROGEN BONDING! Data is plotted on page 466 of text.
Hydrogen Bonding Forces Involves F-H, N-H or O-H bonds. How different are the electronegativity values of these atom pairs? What does this do to the bond polarity? Consider water. As you will soon see, many of water’s unusual Properties result from “hydrogen bonding”! *** * * ** ** ** ** * ** * ******* ** ** *******
Comparison of Intramolecular & Intermolecular Forces FORCE ATTRACTION ENERGY, kJ/mol EXAMPLE Ionic Anion-cation 400 – 4000 NaCl Covalent Shared Electrons 150 - 1100 Cl-Cl Metallic Cations in “Sea of Electrons” 75 - 1000 Cu Ion-Dipole Ion with Dipole 40 – 600 Cl- H2O Dipole-Dipole Dipole charges 5 – 25 Br-Cl Br-Cl H-Bond H to N, O, F 10 - 40 H2O to H2O Ion-induced dipole Ion & e- cloud of neighbor 3 – 15 2 O Fe2 2 Dipole-induced dipole Dipole charge & ecloud of neighbor 2 – 10 H-Br Br2 Dispersion (London) Electron clouds of neighbors 0.05 – 40 Cl-Cl Cl-Cl
Summary: Intermolecular Forces have a wide range of strengths. They are much weaker than ionic, covalent and metallic bonds. Basic types of Intermolecular Forces: Dispersion Forces attraction between temporary ‘induced dipoles.’ Dipole-Dipole attraction between polar molecules (dipoles) having permanent charge separation. Hydrogen Bonds strong attraction between H atom of a molecule and a very electronegative atom (F,O,N) of another molecule.
14-2 Properties of Liquids Intermolecular forces generally determine the physical properties of liquids, such as. Viscosity resistance to motion between molecules of a liquid as they move past each other. Surface Tension unbalanced attractive forces at the surface of a liquid that causes the surface to act like a film. Water bugs, and the paperclip experiment! Capillarity the tendency of a liquid to flow through a small opening or tube. Syrup, motor oil, hot fudge! “Wicking” of fabric when it gets wet. Density mass/volume. Wide range, from 1g/mL (butane) to 13.6g/mL (mercury).
Video Clip (Water)
Comparing Small Molecules Compound Formula Mass (u) State @ 25C MP* BP* Methane CH4 16 Gas 90 112 Ammonia NH3 17 Gas 195 240 Water H2O 18 Liquid 273 373 Nitrogen N2 28 Gas 63 77 Oxygen O2 32 Gas 55 90 *MP Melting Point; BP Boiling Point (kelvins).
Water: A very special compound The most abundant substance on Earth’s surface. (The “Blue Planet.”) Critical to life forms as we know them. Some unusual properties of water are: Unexpectedly high boiling point for its size. Unusually high specific heat. Solid form (ice) has lower density than liquid. High surface tension. High of vaporization. Excellent solvent (“universal solvent”). Why? HYDROGEN BONDING! What are some consequences of these properties?
14-3 The Nature of Solids Crystalline Solids Highly ordered, repeating arrangement of particles. Ionic (NaCl), covalent (sugar) or metallic (Fe). Characterized by specific ‘unit cells.’ (Fig. 14-20) Generally have sharp melting points. Fracture occurs along definite planes when stressed. Amorphous Solids (“supercooled liquids”) Highly disordered, random arrangement of particles. Wax, plastics (PET, PE, PS, Nylon ). Characterized by lack of organized structure. Generally soften over a wide temperature range. Fracture occurs randomly with stress.
Bonding in Solids Metallic Solids All metallic elements. Molecular Solids Most ‘organic’ compounds (contain carbon) & many ‘inorganic’ compounds (CO2, H2O, SO2). Ionic Solids Typical salts (NaCl, KBr, CaCl2) Covalent-Network Solids Diamond, graphite, silicon. See the link: http://undergrad-ed.chemistry.ohio-state.edu/chemapplets/ Crystals/ClosestPackedStructures.html
Properties of Crystalline Solids (Fig 14-23) Type Particles Forces Between Particles Properties Metallic Atoms Metallic bond Soft to hard; variable melting points; good conductivity; malleable; ductile. Molecular Atoms or H-bond, dipolemolecules dipole, dispersion Soft; variable melting points; poor conductors. Ionic Electrostatic Hard; brittle; high melting points; poor conductors. Covalent bonds Very hard; very high melting points; poor conductivity Ions Covalent- Atoms network
14-4 Changes of State (Phase Changes) Phase Change – conversion of a substance from one physical state of matter to another. Six types of phase changes to consider: Melting (solid to liquid) Freezing (liquid to solid) Vaporization (evaporation, boiling) (liquid to gas) Condensation (gas to liquid) Sublimation (solid to gas) Deposition (gas to solid)
Relative Potential Energy (P.E.) Of Matter During Phase Changes (from Fig. 14-30) DEPOSITION SUBLIMATION CONDENSATION VAPORIZATION FREEZING MELTING POTENTIAL ENERGY GAS LIQUID SOLID
Energy Changes & Phase Changes To change a substance from one state of matter to another always involves a gain or loss of energy. Melting a solid requires energy input, but freezing a liquid removes energy from the liquid. Boiling a liquid requires energy input, but condensing a gas removes energy from the gas. Subliming a solid requires energy input, but deposition of a gas removes energy from the gas.
Vaporization & Condensation According to K-M Theory, temperature is a measure of the average kinetic energy of the particles of a substance. Fast molecules near the surface of a liquid may escape the liquid (evaporation). This results in some vapor in the space above a liquid. Volatile liquids have a tendency to easily evaporate due to the poor attraction among their molecules, resulting in high concentrations of vapor around the liquid. Gasoline, for example, poses high hazards because of this property. What happens as fast molecules leave the liquid? The remaining molecules have lower average kinetic energy, which is observed by the lower temperature of the liquid. Evaporative cooling, sweating, wind chill are examples. Liquid-Vapor Equilibrium occurs in closed vessels. Rates of evaporation and condensation are the same. Dynamic equilibrium (rates of opposing processes are equal).
Kinetic Energy Distribution http://undergrad-ed.chemistry.ohio-state.edu/chemapplets/ KineticMolecularTheory/Maxwell.htmlKinetic Energy Distribution In Chapter 13 we looked at two simulations at this website: “Gases – Kinetic Molecular Theory.” (Maxwell distribution.) “Maxwell distribution of a gas (or liquid) at different temperatures.” If you need to refresh your memory, revisit the website to look at the Maxwell distribution shapes. The general shapes of the curves apply to liquids as well.
VAPOR PRESSURE Not all molecules in a sample of a substance have the same amount of kinetic energy. Some molecules are moving faster than others, so there is a distribution of kinetic energy. The average kinetic energy of the substance is proportional to its temperature. Molecules that move very fast may achieve ‘escape velocity’ and leave the liquid entirely. (They evaporate.) This tendency of a liquid to become a gas at a given temperature is its ‘vapor pressure,’ and all liquids exert a vapor pressure against the container (or atmosphere). Vapor pressure increases with temperature. (Why?) Consider what happens to a puddle of water. Does the water evaporate? What if the water is in a closed jar? What would you see? This is an example of “liquid-vapor equilibrium.”
HEAT OF VAPORIZATION is the amount of heat needed to vaporize a given amount of liquid at its boiling point. Energy is added to the liquid, but the temperature does not change. Molar heat of vaporization of water 40.7 kJ/mol or 540. cal/mol. The added energy is expended to overcome the intermolecular attractions of molecules of the liquid, converting the substance from liquid to gas. Liquids with strong intermolecular attractions (such as water) have high heats of vaporization. Condensation is the opposite of evaporations. Heat of Condensation -(Heat of Vaporization) Why is a burn from steam more damaging to tissue than that from boiling water?
“IN CLASS” ACTIVITY Explore Heating & Cooling Curves Explore Phase Diagrams To do these, get one of the laptops from the cart, and go to the websites shown on the following slides. Read the information. Perform the simulations as directed. Take notes as appropriate. Record your observations. Answer the questions.
Heating/Cooling Curves “Phase Changes” Animation and Activity: http://www.chm.davidson.edu/chemistryapplets/Phase Changes/HeatingCurve.html In class, go to this website, read the information, and perform the Heating Curve exercise. Answer all questions. Also, Doc Brown’s website has a nice summary of the states of matter and “heating/cooling” curves. Check it out! http://www.wpbschoolhouse.btinternet.co.uk/page03/3 52states.htm
HEATING/COOLING CURVE SUMMARY Plots time (x-axis) vs. Temperature (y-axis). Identical in shape, but opposite in direction. Ramps indicate temperature changes. Plateaus indicate no change in temperature, but show absorption/release of energy. Two phases co-exist where there is a plateau. Length of plateau is proportional to the energy change that occurs during a phase change. During a phase change, all of the energy is used to induce the change, so the temperature of the mixture of phases stays constant. For example, water on a stove will continue to boil at the same temperature until all of the water has been converted to a gas. Similarly, a mixture of ice and water will remain at zero C until all of the ice melts, after which the liquid begins to heat.
Phase Diagrams In class, go to this animation to observe what happens to phases as temperature and pressure change: http://www.chm.davidson.edu/ChemistryApplets/PhaseChanges/ PhaseDiagram.html At this website (in class) you are to explore all of the following applets: Phase Diagram Phase Diagram Part 1, Part 2, Part 3, Part 4, Part 5. Read the information, perform the simulations and answer all questions. To learn more about phase diagrams, check out this website: http://www.chemistrycoach.com/Phase diagram.htm
PHASE DIAGRAM SUMMARY A plot of pressure and temperature is determined experimentally to show the changes in phase that occur under various conditions. Phases in equilibrium may be identified from the Phase Diagram. Triple Point and Critical Point may be identified. Triple Point – The temperature and pressure at which all three phases exist in equilibrium. Critical Point – The temperature and pressure beyond which the substance can only exist as a gas. http:// www.activotec.com/graphics/products/scfexpl. gif
Ch. 14 OBJECTIVES Show how the Kinetic-Molecular (K-M) Theory accounts for the physical properties of liquids & solids. Describe different types of intermolecular forces, and how they affect properties of liquids & solids. Learn about viscosity & surface tension, and explain their relationship to intermolecular forces. Compare crystalline & amorphous substances. Relate the structure & bonding in the four types of crystalline solids to the properties they exhibit. Describe the changes of state (vaporization, condensation, boiling, sublimation, deposition, melting, freezing). Learn how to interpret “Heating/Cooling Curves” and “Phase Diagrams.”
