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Unit 11: Liquids, Solids, and Solutions

Unit 11: Liquids, Solids, and Solutions. Mr. Gower Presenting The “GowerHour”. I. Introduction A. Liquids and solids are considered to be the _________ states of matter.

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Unit 11: Liquids, Solids, and Solutions

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  1. Unit 11: Liquids, Solids, and Solutions Mr. Gower Presenting The “GowerHour”

  2. I. Introduction A. Liquids and solids are considered to be the _________ states of matter. 1. Unlike ideal gases which have no attractive forces between molecules, substances in the liquid and solid states are held close together by ________ forces. 2. Due to the attractive forces between molecules, the molecules in the condensed states are relatively ____ together. B. Characteristic Properties of Gases, Liquids, and Solids: condensed attractive close

  3. Intermolecular force •• •• •• •• •• •• O O O H H H H H H Intramolecular force between II. Intermolecular Forces A. Types of forces 1. Intermolecular forces are attractive forces _______ molecules. Example: attractive forces between water molecules 2. Intramolecular forces are attractive forces _____ molecules (chemical _____). Example: attractive forces between H and O in a water molecule within bonds

  4. Na Cl– weaker 3. In general, intermolecular forces are much _______ than intramolecular forces. 4. Boiling points and melting points of substances depend upon the strength of _____________ forces. The stronger the intermolecular force, the ______ the boiling point and melting point of the substance. B. Types of Intermolecular Forces 1.Ion – Ion Force: Strong force between __ and __ ions. (ionic bond) Example: sodium and chloride ions intermolecular higher + –

  5. H2O molecule polar 2.Ion – Dipole Force: Attractive force between an ion and a _____ molecule. Example: sodium ion and water molecule / magnesium ion and water molecule. H - O H Stronger interaction w/Mg2+ b/c it has a higher charge & the charge is more concentrated (smaller atom than Na). H - O H

  6. + + + + H H H H Cl Cl Cl Cl - - - - 3.Dipole – Dipole Force: Attractive force between polar molecules. Partial positive (__) end of molecule aligns with partial negative (__) end of another molecule. Example: Hydrogen chloride + – Dipole Dipole Dipole Dipole

  7. Hydrogen Bonds •• •• •• •• •• •• O O O H H H H H H F, O, N 4.Hydrogen Bond: A strong attractive force between hydrogen of one molecule and a small, highly electronegative atom (_______) of another molecule. This is a strong type of ___________ force. Example: Water dipole-dipole

  8. London instantaneous induced all non-polar atoms 5. Dispersion Forces (________): Short-range attractive force for ___ molecules. It can be _______ or ____________. This is the only attractive force for _____ (like Helium) and ________ substances. a. Types of dispersion forces: 1. Induced dipole Helium atom (spherical) Distortion of e- cloud due to a cation Distortion due to a polar molecule.

  9. H H H H H H H H H H H H + + + + - - - - Na H H • Instantaneous dipole At one instant the e- are closer to one H atom. b. Strength of dispersion force depends on: 1. The number of ________ in the atoms of the molecules. 2. Polarizability: The ease with which the electron cloud can be distorted. As the # electrons increases, the polarizability __. Which substance is more polarizable: Helium or Xenon? 3. As the molar mass increases, the dispersion force ________. electrons  more e- increases Summary: Dispersion < Dipole-Dipole < H-bond < Ion-Dipole < Ion-Ion

  10. Decreasing molar mass Decreasing boiling point C. Effects of Intermolecular Forces on Melting Point and Boiling Point 1. In general, the order of increasing strength of intermolecular forces is: Dispersion < Dipole-Dipole < Hydrogen bonding < Ion-Dipole < Ion-Ion 2. As the strength of the intermolecular forces increase, the melting point and boiling point ________. 3. Normally we expect the boiling point of a substance to increase as we move down a group (due to increased __________). But in the figure to the right, H2O, HF, and NH3 have the highest boiling points of Groups VA, VIA, and VIIA. Why? increases molar mass They all have H-bonding. Intermolecular forces  BP 

  11. •• •• O O •• •• •• •• N •• O •• N •• N • D. Examples: 1. Explain in terms of intermolecular forces why: (a) the boiling point of O2 (-183°C) is higher than that of N2 (-196°C). (b) the boiling point of NO is higher than either N2 or O2. dispersion As molar mass  BP  Molar mass of O2 > N2 dispersion Polar: dispersion & dipole-dipole + - Strength of int. forces  BP 

  12. H H Cl Cl Cl C Cl O S •• •• •• C N •• •• H H H dispersion 2. (a) What types of intermolecular forces are present in H2? CCl4? OCS? NH3? CaCl2? dispersion OCS = dispersion, dipole-dipole dispersion, dipole-dipole, hydrogen bonding NH3 = CaCl2 = Ion-ion, dispersion (b) Place the compounds in order of increasing boiling points. H2 < CCl4 < OCS < NH3 < CaCl2

  13. H •• •• •• •• S O H H C •• •• H H H H H O H C H C H H 3. Which of the following are capable of hydrogen bonding: H2S, CH3OH, CH3OCH3? No Yes No “O” must be bonded to the “H”

  14. 4. Explain in terms of intermolecular forces why: (a) ICl has a higher melting point than Br2. (b) C2H6 (Ethane) has a higher boiling point than CH4 (Methane). (c) H2O2 has a higher boiling oint than C2H6. (d) C2H5OH (Ethanol) has a lower boiling point than NaF. (e) C2H5OH has a higher boiling point than butane, C4H10. (a) I – Cl; dispersion, dipole-dipole ☻ Br – Br; dispersion (b) Both non-polar (dispersion)  MW  BP  ☻ (c) H2O2; dispersion, dipole-dipole, H-bond C2H6; dispersion (d) C2H5OH; dispersion, dipole-dipole, H-bond NaF; Ion-Ion, dispersion ☻ ☻ (e) C2H5OH; dispersion, dipole-dipole, H-bond C4H10; dispersion

  15. upward III. Properties of Liquids A. Surface Tension: The contractive force of a surface. 1. The intermolecular attractive forces pull molecules of a surface downward and sideways, but not _______. This causes the molecules at the surface to tighten like an elastic film.  2. The stronger the intermolecular forces, the _______ the surface tension of a liquid. stronger

  16. polar non-polar 3. When water (_____) is placed on a waxed surface (_________), the water molecules are pulled inwards into droplets of water (which minimizes the surface area to volume ratio).   waxed surface H2O C6H6 (Benzene)  4. Water striders can “walk” on water due to the ____ surface tension. 5. As the temperature increases, the surface tension _________.  6. _____ reduces surface tension. (Surfactant) high decreases Soap

  17. B. Cohesion and Adhesion 1. Cohesion: The attractive forces that exist between like molecules. Example: Water molecules attracted to other water molecules. 2. Adhesion: The attractive forces that exist between unlike molecules. This attractive force is what allows things to adhere to other things. 3. Examples: In each case, which is stronger: cohesive or adhesive forces? a. water on wax?   b. (i) water in glass cylinder?   (ii) water in plastic cylinder?   (iii) Hg in glass cylinder?   (i) (ii) (iii)  4. Application: Blood banks use ______ containers for blood so that the blood cells do not adhere to the ____ and become damaged. ad. > coh. ad.  coh. coh. > ad. plastic glass

  18. diameter adheres gravity C. Capillary Action 1. The rise (or depression) of a liquid surface in a small ________ tube called a capillary tube. 2. Why? The water _______ to the glass tube, pulling it upward. Why does it stop? ______. Would the water rise higher or lower in a larger diameter tube? _____. 3. Capillary action in (a) is that of ____. In (b) it is ___. 4. Applications: a. blood tests b. roots in plants • absorption of water in paper d. sponges, towels, etc. lower H2O Hg

  19. resistance greater D. Viscosity 1. Viscosity: The _________ to flow. 2. The more cohesion of the molecules, the _______ the viscosity. 3. Viscosity also depends on _________ features. 4. Measured by timing how long it takes a known amount of liquid to ____ through a thin tube under the force of gravity. 5. As temperature increases, the viscosity ________. 6. Examples: a. hot syrup vs. cold syrup b. Motor oil: rated by _______. The higher the rating, the more viscous the oil. i. Summer: Use ____ viscosity oil. ii. Winter: Use ___ viscosity oil. structural flow decreases viscosity high (It will thin w/high T) low (It will thicken w/ low T)

  20. Hydrogen gas E. The Structure and Properties of Water 1. Water has many unique properties due to the _________ bonding. If water did not have hydrogen bonding it would be a ___ at room temperature!!! 2. Known as the ________ solvent. It dissolves ionic, and polar substances. 3. High specific heat. a. In order to raise the temperature (average ___) of water, the many intermolecular hydrogen bonds must first be broken. Therefore, water can ______ a large amount of heat, or give off a large amount of heat, depending on the temperature of the surroundings. b. The ocean (or any large body of water) acts as a climate _________. In summer, when the air temperature is high, the water absorbs the heat. In winter, when the air temperature is low, the water releases heat. Thus, the temperature in coastal areas remains relatively ________. 4. Solid water (ice) is ___ dense than liquid water. This is due to the extensive 3-D structure of ice which leaves large spaces between molecules in the solid form. universal KE absorb moderator constant less

  21. a. Each oxygen has a ___________-like structure. b. As water nears its MP, hydrogen bonds begin to break, allowing molecules to move into the empty space. Therefore, as water becomes a liquid, it becomes more ____. tetrahedron dense

  22. The density of liquid water is greatest at ____. d. Liquid water expands upon heating, therefore the density generally _________. 4 oC decreases

  23. top sinks convection 5. Ice forms at the ___ of a body of water. a. As the winter approaches and the water temperature decreases, the water _____ to the bottom of the lake, and the warmer water at the bottom rises. This is known as _________. b. The water continues to mix in this way until the water temperature becomes lower than ____. At this point, as the water gets colder, it becomes ____ dense, therefore stays at the ___ of the lake. c. The water temperature reaches the freezing point, and the less dense ice forms at the _______. d. This ice layer at the top of the lake can even as act as an _______, keeping the water temperature below it relatively high. T Lake 4 oC less top surface insulator

  24. IV. Solids A. Types of Solids 1. Crystalline: A solid in which the particles are arranged into an orderly ___ repeating structural pattern. This is also called a crystal ______. 2. Amorphous: “Without _____.” A non-crystalline solid in which the particles are in random arrangement. Generally formed ______, so the atoms do not have time to align themselves in a regular pattern. Example: glass: SiO2 with compounds such as Na2O, B2O3 3-D lattice shape quickly

  25. van der Waal’s forces Types of Crystals 11.6

  26. heat physical V. Phase Changes: Transformations from one phase to another. Phase changes occur when energy (_____) is added or removed. They are _______ changes. A. Liquid-Vapor Equilibrium 1. Evaporation (____________): The process in which a ______ is turned into a ____. 2. Molecules within a liquid have varying amounts of ___. When the KE of a molecule is enough for it to escape from the liquid phase, it becomes a ___. 3. Evaporation can occur at any ___________, but as the temperature increases, the rate of evaporation _________. This is because as the temperature of the liquid increases, the number of molecules with the minimum energy needed to escape _________. vaporization liquid gas KE gas temperature increases increases

  27. minimum 4. In the figure to the right, T1 is __ T2. The shaded area represents the number of molecules with the _________ energy needed to escape into the gas phase. At the higher temperature, more molecules are able to escape.

  28. changing weak Hydrocarbon  non-polar  dispersion only. 5. Volatile: Liquids that have _____ intermolecular forces, therefore they can evaporate easily. Which is more volatile: water or gasoline (C8H18)? Why? 6. Evaporation is a ______ process. A person sweats in order to release excess heat. As the water molecules on the skin evaporate, the high KE molecules turn into a gas, leaving behind the low KE molecules (Cool). 7. Vapor Pressure: When a liquid evaporates, its ________ molecules exert a _______ over the liquid. 8. Equilibrium vapor pressure: The point at which the rate of evaporation and the rate of _____________ are equal. The system reaches a point of dynamic equilibrium, or constant ________. cooling gaseous pressure condensation

  29. increases 9. As temperature increases, the vapor pressure ________. 10. Boiling point: The temperature at which the vapor pressure of a liquid is equal to the _______ pressure. 11. Boiling is a cooling process! As the substance boils, the high KE molecules escape into the gas phase, leaving behind the low KE molecules. If there is a heat source, it will keep the temperature at a _______. external constant

  30. 100 oC 101.4 oC 70 oC 100 oC 101.4 oC 70 oC 760 1 760 mmHg 800 mmHg 240 mmHg 12. Normal boiling point: The boiling point at standard pressure (___ atm or ____ mmHg). (a) (b) (c) (a) External pressure = 760 mmHg; BP of water is _______. (b) External pressure = 800 mmHg; BP of water is _______. (c) External pressure = 240 mmHg; BP of water is _______. (Mt. Everest)

  31. 13. Does it take more or less time to hard boil an egg on Mt. Everest vs. at sea level? Why? 14. Can you cook spaghetti faster if you turn up the flame under a pan of boiling water? 15. Does food cook faster in a pan with a lid on it?  16. How does a pressure cooker work? More time on Mt. Everest because the temperature of the boiling water is on 70 oC!! No, the temperature remains constant at the BP. Yes, the lid traps the high energy molecules. High P  Higher BP  Food cooks faster at a higher T.

  32. evaporation temperature KE B. Critical Temperature and Pressure 1. Condensation: The opposite of ___________. A substance can be made to liquefy by two different processes: a. Lower the ___________, which decreases the ___ of the molecules of the gas phase. b. Raise the ________ on the gas. As the distance between molecules decreases, the strength of the intermolecular forces _________. 2. Critical Temperature: Above the critical temperature, a gas cannot be ________. In other words, this is the maximum temperature at which a substance can exist as a ______. a. Above the Tc, molecules have too much ___ and the intermolecular forces are no longer strong enough to liquefy them. b. As the strength of intermolecular forces increases, the Tc _________.  3. Critical Pressure: The minimum pressure that must be applied to bring about ___________ at the critical temperature. pressure increases liquefied liquid KE increases liquefication

  33. C. Liquid – Solid Equilibrium 1. Freezing: The transformation of a liquid to a _____. The reverse process is ______ or fusion. 2. Melting point (or freezing point) of a substance is the temperature at which the ____ and ______ phases coexist at equilibrium. D. Solid – Vapor Equilibrium 1. Sublimation: The transformation of a solid to a ____. The reverse process is called _________. One example is ______. solid melting solid liquid gas deposition dry ice

  34. Temp, pressure E. Phase Diagrams 1. Phase diagrams summarize the conditions (______________) at which a substance exists as a solid, liquid or gas. 2. A line separating any two phases represents the conditions at which two phases can exist in ___________. 3. Triple point: The conditions under which all three phases can be in equilibrium with one another. 4. The figure to the right shows the phase diagram of water. At 1 atm, water melts at ____ and boils at ______. As the pressure increases, the MP ________ and the BP ________. equilibrium 0 oC 100 oC increases decreases

  35. 5. Explain ice skating. The entire weight of a skater is placed on a thin blade  high pressure. As P , MP , so a layer of liquid H2O allows the skater to glide.

  36. sublimes 5.2 atm 6. The phase diagram of carbon dioxide at right shows that at 1 atm pressure, the substance ________. In order to liquefy carbon dioxide, the pressure must be raised to _____.

  37. 7. Look at the phase diagram for water to the right and predict the results of the following changes: (a) Starting at A, we raise the T at constant P. (b) Starting at C, we lower the T at constant P. (c) Starting at B, we lower the P at constant T. (a) Solid H2O melts (b) Gaseous H2O becomes ice (deposition) L S (c) Liquid H2O vaporizes G

  38. mcT F. Warming and Cooling Curves 1. Up to point A, when heat is added, energy is used to raise the temperature of the solid (Q = _____) 2. From point A to B, the heat is used to change the substance from a solid to a liquid (Q = ____) 3. From point B to C, the heat is used to raise the temperature of the liquid (Q = _____) 4. From point C to D, the heat is used to change the substance from a liquid to a gas (Q = ____)  5. After point D, when heat is added, energy is used to raise the temperature of the gas (Q = _____) mf mcT mv mcT

  39. x x phase 1 g VI. Phase Change Calculations A. Latent heat (): The heat involved in a ______ change. 1. Latent heat of fusion (f): The heat required to melt ___ of a pure substance at its melting point. Q = m f or mHfus • f or Hfus of H2O at 0°C = 334 J/g b. Molar heat of fusion: The heat required to melt _____ of a pure substance at its melting point. Example: water Enthalpy of fusion = Hfus 1 mol 18.02 g 334 J 1 mol H2O = 6012 J 1 mol g = 6.012 kJ

  40. Enthalpy of vaporization = Hvap 2. Latent heat of vaporization (lv): The heat required to vaporize ____ of a pure substance at its boiling point. Q = m v or mHvap • v or Hvap of H2O at 100°C = 2260 J/g B. Specific heat (c): The energy required to raise the temperature of 1 g of a substance by 1°C. 1. c water = 4.184 J/g°C c ice = 2.06 J/g°C c steam = 2.02 J/g°C 1 g

  41. C. Example Calculations: 1. How much heat (kJ) is needed to melt 250 g of ice at 0°C? S  L 0°C Q = mf + Q = (250 g)(334 J/g) = 84 kJ endothermic

  42. - = 10.3 kJ S  L mcT + mf + mcT + mv L/G L 100.0 S/L S 0 -10.0 Q = mf - Q = (50.0 g)(205 J/g) = 10250 J 2. How much heat (kJ) must be removed from 50.0 g of molten copper at its MP to solidify it? 3. How much heat (kJ) must be added to change 100.0 g of ice at –10.0°C into 100.0°C steam? exothermic Q = (100.0 g)(2.06 J/goC)(10.0 oC) + (100.0 g)(334 J/g) + (100.0 g)(4.184 J/goC)(100.0 oC) + (100.0 g)(2260J/g) + endothermic Q = 303,300 J = 303 kJ

  43. 108 (5) (4) 100.0 (3) (2) 0 (1) (mcT)(1) + (mHfus)(2) + (mcT)(3) + (mHvap)(4) + (mcT)(5) 2.      Determine the amount of heat needed (kJ) to completely vaporize 25.4 g of solid ice at –26.0C and heat the vapor to 108C. -26.0 Q = Q = (25.4 g)(2.06 J/goC)(26.0 oC) + (25.4 g)(334 J/g) + (25.4 g)(4.184 J/goC)(100.0 oC) + (25.4 g)(2260 J/g) + (25.4 g)(2.02 J/goC)(8 oC) Q = 78286 J = 80 kJ

  44. Board problemsPhase Change HW Draw heating curve & calculate Q for the following: • 45 g of water @ -15.0 oC to 42.0 oC • 462 g of water @ 15.1 oC to 136.0 oC • 26 g of water @ 2.0 oC to 101.0 oC • 138 g of water @ 0 oC ice to 136 oC.

  45. homogeneous Chapter 16: Solutions I. Introduction A. Solution: A _____________ mixture of two or more substances. Examples of solutions:

  46. B. Parts of Solutions: 1. Solute: The substance being dissolved. In general, the substance in ______ amounts. 2. Solvent: The substance doing the dissolving. The substance in ________ amounts. C. Qualitative Concentrations of Solutions 1. A solution that contains relatively large amounts of solute is considered to be ____________. 2. A solution that contains relatively small amounts of solute is considered to be ______. 3. If solute is added to a solution, it becomes more _____________. 4. If solvent is added to a solution, it becomes more _____. smaller larger concentrated dilute concentrated dilute

  47. solute solution D. Quantitative Concentrations of Solutions 1. Molarity (M): The concentration of a solution expressed in moles _______ per liter of _______. 2. Molality (m): The concentration of a solution expressed in moles _____ per ____ of solvent. 3. Percentage (%): The concentration of a solution expressed as a percentage of solute by _____ or by _______. kg solute mass volume **Used for solutions of liquids in liquids

  48. 4. Mole Fraction (X): Ratio of the number of moles of a component in a solution to the total number of ______ of all of the components in the solution. moles

  49. x KOH E. Sample Problems: 1. A solution contains 10.0 g KOH and 900.0 mL of water. The volume of the solution is 904.7 mL. (a) Which component of the solution is the solute? _____ The solvent?_____ (b) What is the molarity of the solution? (c) What is the % by mass of KOH? (d) What is the mol fraction of KOH? (e) What is the molality of the solution? H2O 10.0 g 1 mol = 0.17822 mol 56.11 g 10.0 g x 100 = 1.10 % KOH 910.0 g 0.17822 mol = 0.00356 0.17822 mol + 49.945 mol 0.17822 mol = 0.198 m 0.9000 kg H2O

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