Chapter 10: States of Matter. Concept Base: Chapter 1: Properties of Matter Chapter 2: Density Chapter 6: Covalent and Ionic Bonding

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1 Chapter 10: States of Matter Concept Base: Chapter 1: Properties of Matter Chapter 2: Density Chapter 6: Covalent and Ionic Bonding

2 Pressure standard pressure the pressure exerted at sea level in dry air 760. mmhg 29.9 inhg 760. torr kpa x 10 5 N/m 2 (Pa) 1.00 atm 14.7 psi F P = A

3 Pressure A column of air 1.00 m 2 in cross-sectional area extending from the earth s surface through the upper atmosphere has a mass of about 10,300 kg (22700 lb), producing an atmospheric pressure of approximately 101,000 Pa.

4 What do you think? Which city has higher atmospheric pressure, Pittsburgh or Denver?

5 Covalent Vs. Ionic Bonding atoms combine to form ionic bonds covalent bonds (M + NM) (NM + NM) chemical bond a mutual electrical attraction between the nuclei and valence electrons of two atoms that binds the atoms together

6 Ionic Bonding ionic bond when electrons are taken by one atom from another atom metal and a nonmetal NaCl cation and anion

7 Covalent Bonding covalent bond when electrons are shared between two atoms two nonmetals No ions formed! (no electrons are taken) H-H.... O H H

8 There is another type of bond, not purely Pure Nonpolar Covalent Polar Covalent covalent and not purely ionic. Ionic

9 The Kinetic-Molecular Theory of Matter In reality, all atoms are moving: vibrating, rotating, translating. The only time something would NOT be moving is at ABSOLUTE ZERO, (0 Kelvin or -273 oc).

10 The Kinetic-Molecular Theory of Gases ideal gas hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory. real gas a gas that does not fit all the assumptions of the kinetic-molecular theory. Although an ideal gas does not exist, many gases behave nearly ideally if pressure is not very high and temperature is not very low.

11 The Kinetic-Molecular Theory of Gases 1. Gases consist of large numbers of tiny particles that are far apart relative to their size Gases take up much more space than solids or liquids Gases are easily compressed

12 The Kinetic-Molecular Theory of Gases 2. Collisions between gas particles and between particles and container walls are elastic collisions. elastic collision a collision with no loss of energy inelastic collision a collision with some loss of energy

13 The Kinetic-Molecular Theory of Gases 3. Gas particles are in continuous, rapid, straightline, random motion. They therefore possess kinetic energy, which is energy of motion. The KE that the molecules have as a gas overcomes any attractive forces that they might have.

14 The Kinetic-Molecular Theory of Gases 4. There are no forces of attraction between gas particles.

15 The Kinetic-Molecular Theory of Gases 5. The temperature of a gas depends on the average kinetic energy (energy of motion) of the particles of the gas KE T Kelvin

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17 The Nature of Gases Gases do not have a definite shape or volume They completely fill any container because they have no attraction for the other molecules in the container (unlike liquids and solids).

18 The Nature of Gases fluid a nonsolid state of matter in which the atoms or molecules are free to move past each other, as in a gas or liquid (flows) Attractive forces are insignificant, thus gas particles glide easily past one another and are called fluids.

19 The Nature of Gases Gases have low densities. M D = V

20 The Nature of Gases Gases can be compressed Thus, many more molecules can be inside a container under great amounts of pressure.

21 The Nature of Gases Gases diffuse readily into one another. Flasks are NOT connected. Valve is opened to connect the flasks.

22 The Nature of Gases diffusion the spontaneous movement of particles of gas caused by random motion effusion process by which gas particles pass through a tiny opening due to pressure being exerted upon them

23 KE = ½mv 2 Gas molecules move at different speeds or velocities, depending on the temperature at which the molecules are the molar mass of the molecules The higher the MM, the slower they move.

24 Check for Understanding What kind of states of matter can be poured? Which molecules are moving faster, water at 50 o C or water at 20 o C? Which state(s) of matter can be compressed to great extents?

25 Intermolecular Forces pages Chapter 6 intermolecular forces - attractive forces between molecules These forces vary in strength, however are generally much weaker than covalent bonds (not intermolecular). The stronger the intermolecular force the closer the molecules get to one another, thus perhaps creating a solid versus a liquid, etc.

26 What do you think? Now that we learned about intermolecular forces, what do you think an intramolecular force is? Give two examples. Intramolecular Forces > Intermolecular Forces (stronger than)

27 Dipole-Dipole Forces dipole a molecule or a part of a molecule that contains both partial positive and partial negative regions A dipole is created when there is a large difference in electronegativity.

28 Dipole-Dipole Forces The partial positive end is attracted to the partial negative end.

29 Dipole-Dipole Forces Compare ICl to Br 2 because they have approximately the same molar mass ICl Br 2 BP = 97 o C BP = 59 o C What makes ICl have a higher BP?

30 Dipole-Induced Dipole Attraction

31 Ion-Induced Dipole Attraction

32 What do you think? Looking at the previous intermolecular forces, do you think it would be possible to have an ion-dipole attraction? Explain.

33 Hydrogen Bonds F-H (HF) O-H (H 2 O) N-H (NH 3 ) All of these bonds have a large difference in electronegativity, thus creating a large dipole, or a highly polar bond. These highly polar bonds have a very strong attraction. These very strong attractions are called hydrogen bonds.

34 Hydrogen Bonds Hydrogen Bonding in Water

35 Hydrogen Bonding in Water Because of the hydrogen bonding in water, an open, rigid structure is formed when freezing. As a solid, there is more hydrogen bonding than as a liquid. D ice at 0oC = g/ml D water at 0oC = g/ml

36 Hydrogen Bonds Hydrogen Bonding in Acetic Acid HC 2 H 3 O 2

37 Hydrogen Bonds Compare H 2 O to H 2 S H 2 O BP = 100 o C H 2 S BP = -61 o C Compare NH 3 to PH 3 NH 3 BP = -33 o C PH 3 BP = -88 o C What makes H 2 O and NH 3 have higher boiling points?

38 Hydrogen Bonds Snowflakes are large ice crystals that have a unique shape. The shape reflects the rigid position of the hydrogen bonding of the solid.

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41 London Dispersion Forces Nonpolar molecules will also exhibit a weak attraction for one another. The constant motion of electrons within a molecule can create a temporary dipole, or London dispersion force, that attracts to another temporary dipole.

42 London Dispersion Forces London forces are dependent upon the motion of electrons. Therefore, the more electrons, the greater the London forces. molar mass London forces

43 Check for Understanding Compare Cl 2, Br 2, and I 2 and arrange them according to the strength of their London forces. What are their states of matter at room temperature? WHY?

44 Intermolecular Forces Strength of Intermolecular Forces hydrogen bonding from Highest to Lowest dipole-dipole attraction dipole-induced dipole attraction London dispersion forces van der Waals forces - any dipole forces and London forces are groups as these vdw forces

45 Check for Understanding What physical property directly correlates with the strength of the London dispersion forces? What types of molecules have the strongest intermolecular forces? What do you think accounts for NH 3, ammonia, having a boiling point 130 o C higher than CH 4, methane?

46 Comparison of Boiling Points

47 Liquids How do you know that this contains a liquid without opening it?

48 Liquids Liquid nitrogen changing to gaseous nitrogen

49 Liquids LIQUIDS have a definite volume and take on the shape of their container (unlike gases) have a high density are not compressed well (brake fluid) diffuse (like food coloring in water) have surface tension are fluids (but fluid liquid)

50 Liquids surface tension a force that tends to pull adjacent parts of a liquid s surface together, thereby decreasing surface area to the smallest possible size. Surface tension results from attractive forces between the particles in the liquid. The stronger the attractive force, the higher the surface tension.

51 Liquids Notice that the pull on the mercury atoms at the top is not symmetrical. That is what gives the characteristic spherical shape to drops of liquid.

52 Liquids viscosity the resistance of a liquid to flow high viscosity = thick liquid low viscosity = thin liquid Liquids with stronger intermolecular forces have higher viscosity. An increase in temperature will decrease the viscosity. KE T Kelvin

53 Liquids volatile liquid - a liquid that evaporates readily at low temperatures The higher the volatility of a liquid, the weaker the intermolecular forces of attraction between their particles. An increase in T will increase evaporation. KE T Kelvin

54 Liquids In each cylinder, can you describe which forces are greater: cohesive or adhesive? cohesive forces forces of attraction between like molecules (H 2 O to H 2 O) adhesive forces forces of attraction between unlike molecules (H 2 O to glass)

55 Liquids capillary action - the attraction of the surface of a liquid to the surface of a solid (adhesive forces) Many liquids will creep along a solid, like water does to paper or cloth fibers until the pull of gravity is too much for it to overcome.

56 Check for Understanding Order the following liquids from highest surface tension to lowest? Explain. Br 2, H 2 O, H 2 S Which has a higher viscosity, water or carbon tetrachloride, CCl 4? How is this waterstrider able to walk on water?

57 Evaporative Cooling Why do you feel cold when you get out of the shower? Why does fanning yourself cool you down? Why do you sweat when you have a fever?

58 Solids Solids are in a relatively fixed position. Solids have only vibrational movements around fixed points. Solids have definite shape and volume. Solids are almost incompressible. Solids do not diffuse (practically).

59 Solids Solids are either crystalline or amorphous. crystalline consist of crystals, particles arranged in an orderly, geometric repeating pattern amorphous Greek for without shape ; consist of particles, randomly arranged

60 Crystalline Solids Crystalline solids break into orderly pieces. After breaking salt, the cubic structure is still visible. NaCl is cubic.

61 Crystalline Solids crystal structure three-dimensional arrangement of particles of a crystal, represented by a lattice unit cell The smallest portion of a crystal lattice that shows the 3-D pattern Be familiar with the seven basic crystalline systems: Figure 11 - p.339

62 Crystalline Systems

63 Amorphous Solids Amorphous solids break into random pieces. They usually shatter into irregular shapes. Most plastics are amorphous.

64 Amorphous Solids The freezing point of amorphous solids can vary according to how slowly the material cools. (Ex: butter) /General/Glass/glass.html

65 Crystal Types ionic covalent covalent network each atom is covalently bonding to its neighbors covalent molecular each molecule is held together by intermolecular forces metallic metallic bonding

66 Ionic Crystals ionic bonding high melting points positive and negative ions Ex: NaCl, MgF 2

67 Held together by covalent bonds. Covalent Network Crystals Diamond (C) and SiO 2 are examples of solids that form these giant molecules.

68 Covalent Network Crystals

69 Held together by intermolecular forces. Covalent Molecular Crystals Water is an example of covalent molecular crystals due to its hydrogen bonding.

70 Metallic Crystals metallic bonding the chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons Metallic bonding allows metals to be good conductors of electricity, malleable, and ductile. Ex: Hg, Cu, Fe, W

71 Checking for Understanding Name at least two common examples of amorphous substances. What type of crystal would ammonia, NH 3, as a solid be likely classified as? Why do you think ionic crystals have such high melting points, thus are usually found as solids at room temperature?

72 Checking for Understanding What is glass: crystalline or amorphous? Support your answer.

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75 The Kinetic Theory of Heat and Temperature When a phase change is occurring, the temperature does not change, only the position of the particles. (PE) When something is being heated and it is not changing phase, the temperature will rise. (KE)

76 Potential Energy Differences PE energy of position

77 PE vs. KE Only differences in kinetic energy are reflected by temperature differences. Difference in potential energy are NOT reflected in temperature differences. THERE IS NO TEMPERATURE CHANGE DURING A PHASE CHANGE.

78 Heating Curve for Water

79 K v = heat of vaporization, the amount of heat needed to boil/condense 1g of a substance Phase Changes During solidification or melting H = K f m K f = heat of fusion, the amount of heat needed to melt/freeze 1g of a substance During boiling or condensing H = K v m

80 Phase Changes K f for water = 333 J/g K v for water = 2260 J/g

81 Heating Curve for Water

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83 Frost represents Deposition

84 Vapor Pressure vapor pressure the pressure due to a vapor above a liquid How does the vapor get above the liquid? How do you increase the vapor pressure?

85 Open System:

86 Closed System

87 Pressure How can you increase the pressure, P, of a gas inside a container? increase the temperature, T (KE) of the gas decrease the size (volume, V) of the container add more molecules to the container (increase the number of moles, n)

88 Dynamic equilibrium occurs when the rate of evaporation = rate of condensation.

89 Equilibrium dynamic equilibrium Although there are two opposing processes going on, they are occurring at the same rate. vapor vapor liquid solid

90 Vapor Pressure

91 Boiling P vapor = P external boiling the conversion of a liquid to a vapor not only at its surface, but within the liquid as well boiling point The temperature at which the vapor pressure is equal to the external pressure (usually atmospheric pressure) normal boiling point The temperature at which the vapor pressure is equal to standard pressure (1.00 atm)

92 Boiling When P vapor P atmosphere, the liquid will boil.

93 What do you think? vacuum pump How is it possible that this beaker of water is boiling at room temperature?

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95 Phase Diagrams CRITICAL PRESSURE (the lowest pressure at which the substance can still be a liquid at the critical T) CRITICAL POINT TRIPLE POINT CRITICAL TEMPERATURE (no more liquids above this)

96 Phase Diagram for Water

97 Carbon Phase Diagram

98 What do you think? Why is it that a closed bottle of water, of which part has been consumed, has condensation all over the inside of the container after time?

99 What do you think? If you can boil water in a vacuum pump, can you hard boil an egg in the boiling water in the vacuum pump?

100 What do you think? Why is it that the snow in your driveway melts when you back over it with the car?

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