Gases, Liquids, and Solids. Chapter 5

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1 Gases, Liquids, and Solids Chapter 5

2 Educational Goals 1. Define, compare, contrast the terms specific heat, heat of fusion, and heat of vaporization. Know the equations that involve these concepts and be able to use them in calculations. 2. Describe the meaning of the terms enthalpy change, entropy change, and free energy change. Explain if a process is spontaneous or not based on the free energy change. 3. Convert between pressure units of atm, torr, and psi. 4. List the variables that describe the condition of a gas and give the equations for the various gas laws.

3 Educational Goals (cont.) 5. Explain Dalton s law of partial pressures. 6. Define the terms density and specific gravity. Given the density, and either the mass or volume, be able to determine the volume or mass (respectively). 7. Know that a liquid will boil when its vapor pressure is equal to the atmospheric pressure. 8. Describe, compare, and contrast amorphous solids and crystalline solids. 9. Describe the makeup of the four classes of crystalline solids.

4 Phases of Matter and Energy Example: Three phases of water solid ice H 2 O(s) liquid water H 2 O(l) gas steam H 2 O(g)

5 Why are some molecular compounds solid while others are gaseous and others are liquid at room temperature? Competing Powers Intermolecular forces working to hold particles together as liquids or solids Kinetic Energy = Motion = Temperature, work to separate particles Kinetic Energy = Temperature

6 One major factor that is responsible for the varied behavior of solids, liquids, and gases is the nature of the interaction that attracts one particle (atom, ion, or molecule) to another.

7 What forces hold matter together to make liquids and solids?

8 What forces hold matter together to make liquids and solids? The attractive forces that hold molecules together are called Intermolecular forces.

9 3 Types of Intermolecular Forces 1) Dipole-Dipole Forces 2) Hydrogen Bonding 3) London Forces also called Induced Dipole Forces

10 Other Noncovalent Interactions Noncovalent interactions are interactions that do not involve the sharing of valence electrons (covalent bonding). Other noncovalent interactions due to the attraction of permanent charges. 1) Salt bridges 2) Ion-dipole interactions

11 A salt bridge is another name for ionic bond. Ion-dipole interactions occur between ions with a full charge and atoms with a partial charge.

12 Energy meets Matter Adding energy to liquids will overcome the forces holding the molecules together boiling Adding energy to solids will overcome the forces holding the molecules together melting

13 Phase Changes Language

14 To reduce a fever, rubbing alcohol (2- propanol) can be applied to the skin. As the alcohol evaporates (liquid becomes a gas), the skin cools. Explain the changes in heat energy as this process takes place. Note: 2-propanol vapors are flammable, so care must be taken when using this technique.

15 Units of Energy One calorie is the amount of energy needed to raise the temperature of one gram of water by 1 C joule J = 1 cal In nutrition, calories are capitalized 1 Cal = 1 kcal = 1000 cal

16 Converting Between Calories and Joules Example: Convert 60.1 cal to joules Equivalence statement: 1 cal = J 1 cal Conversion J J Factors 1 cal 60.1 cal J 1 cal = 251 J

17 Calculations Involving Heat Energy One of two things will happen if energy is added or removed from matter (assuming no chemical change takes place). 1) Change the phase of the substance Example: melt, freeze, vaporize (boil) 2) Change the temperature of the substance You can only do one of these at a time!!! See graph on the next slide

19 1) Phase Changes Energy calculations for phase changes may be carried out using the tabulated values for: heat of fusion (symbol = H fus ) for a substance (Table 5.2). Energy required to melt one gram of a solid Change sign to negative for freezing (liquid to solid)

20 1) Phase Changes Energy calculations for phase changes may be carried out using the tabulated values for: heat of vaporization (symbol H vap ) of a substance Energy required to vaporize one gram of a liquid Change sign to negative for gas going to liquid

21 Delta ( ) means change in Calculations Involving Changing the Phase Energy Change = (mass) x (heat of fusion or vaporization) ΔE = m x (H fus or vap ) Get from a table

22 Calculations Involving Heat Energy Example: Determine the amount of energy needed to melt 155 g of ice at 0 C, we use the heat of fusion of water (79.7 cal/g) as a conversion factor. ΔE = (mass) x (heat of fusion) 155 g x 79.7 cal = 1.24 x 10 4 cal g Note: No Temperature Change! Ice (0 o C) Water (0 o C)

23 Group Work A patient with a fever is sponged with 50.0 g of 2-propanol. How much energy is drawn from the patient when 2-propanol vaporizes? (heat of vaporization for 2-propanol is 159 cal/g)

24 2) Changing Temperature of Matter The amount the temperature of an object increases depends on the amount of energy added (Q). If you double the added heat energy the temperature will increase twice as much. The amount the temperature of an object increases depends on its mass. If you double the mass it will take twice as much heat energy to raise the temperature the same amount.

25 Calculations Involving Changing the Temperature Energy calculations may be carried out using the values for the specific heat of a substance. Specific heat is the amount of energy required to raise the temperature of one gram of a substance by one Celsius degree.

26 Energy required = Specific Heat x Mass x Temperature Change Q = S x m x (ΔT) is always: (final) (initial) ( T) = T final -T initial By definition, the specificheat of water is g J C

27 Example: Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0 C to 46.0 C Mass = 7.40 g Q = S x m x ΔT Specific heat of water: J g o C Temperature Change (ΔT) = 46.0 C 29.0 C = 17.0 C Q = J g o C (7.40 g) (17.0 o C) =526 J

28 Group Work How much energy needs to be removed from 175 g of water to lower the temperature from 23.0 o C to 15.0 o C?

29 How much energy is required to convert 25g of ice at -7.0 o C to water at 50.0 o C Ice Water -7.0 o C 50.0 o C Ice Water Step o C 0.0 o C Step 2 Step 3 Temperature Change Q 1 =S ice xmx( T) Phase Change E 2 =mxh fus Temperature Change Q 3 =S water xmx( T) Energy Total = Q 1 + E 2 + Q 3

30 New Topic: Will a change occur? Spontaneous vs. Nonspontaneous Changes An important question to ask is why some changes are: spontaneous (continue to occur once they are started) OR nonspontaneous (will not run by themselves unless something keeps them going). Energy is the key factor in determining this.

31 Spontaneous vs. Nonspontaneous Changes Spontaneous vs. Nonspontaneous

32 Energy vs. Free Energy The energy (E) of matter depends on the position (potential energy) and velocity (kinetic energy) of every molecule in the system. E = E potential + E kinetic This is not practical to measure in the lab or to model in calculations! When working at constant temperature and pressure, it is mathematically convenient and experimentally practical to look at the: Free Energy (G)

33 Energy vs. Free Energy Just like the energy (E), in nature, given the chance, everything proceeds to the lowest possible free energy (G)!

34 Free Energy (G) The free energy (ΔG) of a process can be thought of as the potential for change. G = G f - G i A spontaneous process has a negative G and a nonspontaneous process has a positive G.

35 Gases and Pressures Properties of Gases How to visualize a gas: Gas molecules or atoms are very far apart from one another. different from liquids and solids!! Gas particles move in a straight line until they collide with another particle or the container wall.

36 Properties of Gases Gases Have Low Density Because of the relatively large distances between gas particles, most of the volume occupied by a gas is empty space.

37 Properties of Gases Gases completely fill their container. Except for a few very heavy gases, most gasses will completely fill their container.

38 Properties of Gases Gases Are Highly Compressible Gases are compressible Liquids and Solids are not

39 Properties of Gases Gases Are Highly Compressible Compressibility is the ability to make the space a substance takes up become smaller.

40 Properties of Gases Gases can diffuse. Gaseous molecules travel at high speeds in all directions and mix quickly with molecules of gases in the air in a process called diffusion. Diffusion is the movement of one substance within another substance until it is evenly distributed.

41 Ammonia Skunk in da house Properties of Gases Examples of diffusion.

42 Gas Pressure Pressure = total force applied to a certain area larger force = larger pressure Gas pressure is caused by gas molecules colliding with container walls or surfaces.

43 Air Pressure Constantly present when air present Decreases with altitude less air

45 Air Pressure Measured using a barometer Column of mercury supported by air pressure Force of the air on the surface of the mercury balanced by the pull of gravity on the column of mercury

46 Various Units for Gas Pressure 1) atmosphere (atm) 2) height of a column of mercury (mm Hg, in Hg) 3) Torr 4) Pascal (Pa) 6) pounds per square inch (psi, lbs./in 2 )

47 Units we will use for pressure: Atmospheres (atm) Pounds per square inch (psi) Millimeters of mercury (mm Hg) also called torr (1mm Hg = 1 Torr) 760. mm Hg 1 atm Relationships: 1 atm = 760. mmhg 1 atm 760. mm Hg 760. Torr 1 atm 14.7 psi 1 atm 1 atm = 760. Torr 1 atm = 14.7 psi 1 atm 760. Torr 1 atm 14.7 psi

48 Pressure Unit Conversions A pressure of 690. Torr is how many atmospheres? 1 atm = 760 Torr 690. Torr 1 atm 760 Torr =.908 atm

49 Group Work A pressure of 35.0 psi is how many atm? 1 atm = 14.7 psi A pressure of 812 mm Hg is how many atmospheres? 1 atm = 760. mm Hg

50 Gas Laws

51 Instructional Goals Understand and be able to use the following gas laws in calculations: Boyles Law (relationship between pressure and volume) Charles Law (relationship between volume and temperature) Gay-Lussac s Law (relationship between pressure and temperature) Avogadro s Law (relationship between moles and volume) Combined Gas Law (relationship between pressure, volume and temperature) Ideal Gas Law (relationship between pressure, volume, number of moles, and temperature)

52 The Gas Laws The gas laws are the mathematical equations that show the relationship between volume, temperature, pressure, and amount of gas. As with all laws, they were discovered by experiments.

53 Boyle s Law Boyle studied the relationship between volume and pressure. The inverse relationship between pressure and volume is known as Boyle s law.

54 Boyle s Law When the volume decreases, the pressure increases Pressure Gage low high

55 Boyle s Law When the volume increases, the pressure decreases Pressure Gage low high

56 Boyle s Law Boyle also noticed that when the pressure and/ or volume of a gas is changed the product of the pressure and volume remains the same. P x V = Constant

57 Initial Initial Final Final pressure volume pressure volume P 1 V 1 = P 2 V 2 Boyle s Law

58 Remember that when using Boyle s Law, that the temperature is never changing. Only the pressure and volume change.

59 Example The initial volume of the gas in the piston below is 3.00 liters and the initial pressure is 1.00 atm. The piston compressed (at constant temperature) to a new final volume of 1.00 L. What is the final pressure? Pressure Gage low high

60 Solution P 1 V P 1 = 2 V 2 P 1 V 1 = P 2 V 2 P 1 = 1.00 atm P 2 =? V 2 V 2 V 1 = 3.00 L V 2 = 1.00 L P P 1 2 = V 1 V 2 = (1.00 atm) (1.00 L) (3.00 L) = 3.00 atm

61 Group Work If the syringe shown has an initial volume of 0.50 ml and the gas in the syringe is at a pressure of 1.0 atm, what is the pressure inside the syringe if your finger is placed over the opening and the plunger is pulled back to give a final volume of 3.0 ml?

62 Charles Law Charles observed that as the temperature increases, the volume increases and vice versa. Jacques Charles ( ) The direct relationship between temperature and volume is known as Charles law.

63 Charles Law When the temperature increases, the volume increases

64 Charles Law When the temperature decreases, the volume decreases

65 Charles Law Charles also noticed that ratio of volume to temperature of a gas is always the same. V T = Constant

66 Charles Law

67 Charles Law Initial volume Initial temperature V 1 V 2 = T 1 T 2 Final volume Final temperature

68 Remember that when using Charles Law, that the pressure is never changing. Only the temperature and volume change. Temperature must be Kelvin (K). Kelvin temperature scale is always positive K = o C

69 Example The initial volume of the gas in the piston below is 1.35 liters and the initial pressure is 1.00 atm. The temperature is lowered from 373 K to 250. K (at constant pressure). What is the final volume? Pressure Gage low high

70 Solution V 2 V 2 T 2 V 1 = T 1 T 2 T 2 V 1 = T 1 T 2 V 1 = 1.35 L V 2 =? T 1 = 373 K T 2 = 250. K T V 2 2 = V 1 T 1 (1.35 L) = (250.K) (373K) = L

71 Group Work A balloon is inflated to 665 ml volume at 27 C. It is immersed in a dry-ice bath at 79 C. What is its volume, assuming the pressure remains constant? V 1 = T 1 V 2 T 2 Remember to convert to Kelvin (K) K = o C

72 Gay-Lussac s Law Gay-Lussac s observed that as the temperature increases, the pressure increases and vice versa. Joseph Gay-Lussac ( ) The direct relationship between temperature and pressure is known as Gay-Lussac s Law.

73 Gay-Lussac s Law When the temperature decreases, the pressure decreases. Pressure Gage low high

74 Gay-Lussac s Law When the temperature increases, the pressure increases. Pressure Gage low high

75 Gay-Lussac s Law Gay-Lussac also noticed that ratio of pressure to temperature of a gas is always the same. P T = Constant

76 Gay-Lussac s Law Initial pressure Initial temperature P 1 P 2 = T 1 T 2 Final pressure Final temperature

77 Remember that when using Gay-Lussac s Law, that the volume is never changing. Only the temperature and pressure change. Temperature must be Kelvin (K).

78 Example The initial pressure of the gas in the container below is.870 torr and the initial temperature is. 300.K. The temperature is raised from 300. K to K (at constant volume). What is the final pressure? Pressure Gage low high

79 Solution P 1 = P 2 T 2 P 1 = P 2 T 2 T 1 T 2 T 1 T 2 P 1 = torr P 2 =? T 1 = 300. K T 2 = K T P 2 2 = P 1 T 1 (1250. K) (0.870 tor) = (300.K) = 3.63 torr

80 Group Work An aerosol can containing gas at 25 atm and 22 C is heated to 55 C. Calculate the pressure in the heated can. P 1 = T 1 P 2 T 2 Remember to convert to Kelvin (K) K = o C

81 The Combined Gas Law Boyles s, Charles s, and Gay-Lussac s Laws can be combined mathematically. The relationship between temperature, volume, and pressure is known as the Combined Gas Law.

82 The Combined Gas Law P V P V = 2 T T 1 2

83 Example At an ocean depth of 33 ft, where the pressure is 2.0 atm and the temperature is 285K, a scuba diver releases a bubble of air with a volume of 6.0 ml. What is the volume of the air bubble when it reaches the surface, where the pressure is 1.0 atm and the temperature is 298 K?

84 Solution P 1 = 2.0 atm P 1 V 1 = T 1 P 2 V 2 T 2 P 2 = 1.0 atm T V 2 P 1 V 2 = 1 P 2 T 1 V 1 = 6.0 ml V 2 =? T 1 = 285 K T 2 = 298 K (298 K) (2.0 atm) (6.0 ml) = = 13 ml (1.0 atm) (285K)

85 Avogadro s Law Avogadro s observed that the volume of a gas is directly proportional to the number of gas molecules. Amedeo Avogadro ( ) The direct relationship between moles of gas molecules and volume is known as Avogadro s Law.

86 Avogadro s Law When the number of moles of gas decreases, the volume decreases. Pressure Gage low high

87 Avogadro s Law When the number of moles of gas increases, the volume increases. Pressure Gage low high

88 Avogadro s Law Avogadro noticed that ratio of volume to the number of moles of a gas is always the same. V n = Constant

89 Avogadro s Law Initial volume V 1 V 2 = n 1 Initial # moles n 2 Final volume Final # moles

90 Remember that when using Avogadro s Law, that the pressure and temperature are never changing. Only the number of particles and volume change.

91 Example The initial volume of the 3.5 moles of gas in a container is 1.5 L. Amadeo adds 2.0 moles of gas. (at constant temperature and pressure). What is the final volume?

92 Solution V 1 = V 2 n 2 V 1 = V 2 n 2 n 1 n 2 n 1 n 2 V 1 = 1.5 L V 2 =? n 1 = 3.5 mol n 2 = 5.5 mol n V 2 2 = n 1 V 1 = (5.5 mol) (1.5 L) (3.5 mol) = 2.4 L

93 Group Work A balloon has a volume of 2.4 L and contains 0.12 moles of air. A child blows more air into the balloon until it has a final volume of 3.5 L. How many moles of gas are in the balloon?

94 Gas Law Summary V 1 V 2 n = 1 n 2 Combined Gas Law P 1 V 1 = T 1 P 2 V 2 T 2

95 The Ideal Gas Law No gas perfectly obeys all four of these laws under all conditions. These assumptions work well for most gases and most conditions. One way to model a gas s behavior is to assume that the gas is an ideal gas that perfectly follows these laws.

96 The Ideal Gas Law If we combine all these equations, we get the Ideal Gas Law. P x V = C b V n = C a V T = C c P T = C g PV nt = R Gas Constant

97 The Ideal Gas Law The gas constant (R) is a mathematical combination of all the individual gas law constants (C b, C c, C g, C a ) PV nt = R The Ideal Gas Law is more commonly written as: PV = nrt

98 The Ideal Gas Law The previous gas laws we studied involved a change in either P, V, T, or n. P 1 = P 2 P 1 V 1 = P 2 V 2 T 1 T 2 V 1 = V 2 V 1 = V 2 n 1 n 2 T 1 T 2

99 The Ideal Gas Law The ideal gas law is used for any gas system, any time. No changes are involved in the equation PV = nrt

100 Temperature = K The Ideal Gas Law The value of R is:.0821 L. atm K. mol When using this equation you must have the following units: Pressure = atm Volume = liters PV = nrt

101 The Ideal Gas Law There are 4 variables in this equation: PV = nrt Pressure Volume # Moles Temperature In problems, we will always be given 3 of the 4 variables, then solve for the unknown variable.

102 Example: The Ideal Gas Law How many moles of gas are contained in 11.2 liters at 1.00 atm and 0.0 C? PV = n RT P = 1.00 atm V = 11.2 L n =??? T = K

103 Partial Pressure

104 Partial pressure of gas A Partial pressure of gas C Dalton s law of partial pressure states that the total pressure of a mixture of gases is the sum of the partial pressures of its components. The partial pressure of a gas in a mixture is the pressure that the gas would exert if alone. Partial pressure of gas B Total pressure P T = P A + P B + P C

105 When two gases are present, the total pressure is the sum of the partial pressures of the gases.

106 P Partial Pressures The partial pressure of each gas in a mixture can be calculated using the Ideal Gas Law for gases na x R x T n PA = PB = V the temperature and volume of in the mixture are the same total = P n A and B in a total A + = P B n A = + n n B total mixture B x R x T V everything x R x T V

107 Let s Try It! A 1.00L flask contains 5.00 x 10-2 mol of neon and 5.00 x 10-3 mol of argon. At 30.0 C, what is the partial pressure of each gas in atmospheres and what is the total pressure?

108 Liquids Viscosity is the resistance to flow. - It is related to the strength of the noncovalent interactions between the molecules that make up the liquid - the stronger the attractions, the thicker the liquid. - Temperature has an effect on viscosity. - As temperature rises, the increase in the kinetic energy of the molecules in the liquid helps the molecules pull away from one another - higher temperature produces lower viscosity.

109 Glycerol is able to form more hydrogen bonds than 2-propanol. That is why glycerol is thicker (more viscous) than 2-propanol.

110 Density of a liquid (or any other substance) is the amount of mass contained in a given volume. d = m V Density is the relationship between the mass of a substance and its volume.

111

112 Vapor Pressure Due to collisions that take place between particles (atoms or molecules) that make up a liquid, particles at the surface are continually evaporating - being bounced off into the gas phase. At the same time gas phase molecules are being trapped and converted to liquid.

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114 Vapor Pressure The boiling point of a liquid is the temperature at which the vapor pressure of the liquid equals the atmospheric pressure. Liquids boil when their vapor pressure equals the pressure of the air above them.

115 Solids The atoms, ions, or molecules that make up a solid are held close to one another and have a limited ability to move around. Solids can be classified based on whether or not the arrangement of these particles is ordered (in crystalline solids) or not (in amorphous solids).

116 Ionic Crystalline Solids consist of oppositely charged ions held to one another by ionic bonds Molecular consist of an ordered arrangement of molecules attracted to one another by noncovalent interactions Covalent Networks atoms are held to one another by an arrangement of covalent bonds that extends through the solids. Metallic An array of metal cations immersed in a cloud of electrons that spans the entire crystalline structure.

117 Metallic Solids Metallic Bonding The valence electrons in metals are free to move about the entire crystal of metal nuclei and core electrons. We can imagine it like a sea of electrons that are bonding the positive nuclei together

118 Properties of Metallic Substances Metallic substances are solid at room temperature. Except for : Metallic substances are malleable (they can be hammered or beaten in thin sheets) Metallic substances are ductile (they can be drawn, pulled, or extruded through a small opening to produce wire. Metallic substances are good conductors of electricity.

119 A few substances exist as: Covalent Networks Atoms are covalently bonded as if it was a huge molecule Not too many covalent network substances exist Examples: Diamond (carbon)

120 Amorphous Solids- no regular repeating pattern of ions or molecules. Example: rubber

121 Solids: Summary a) Ionic b) Covalent c) Molecular d) Metallic e) Amorphous

States of Matter. The Solid State. Particles are tightly packed, very close together (strong cohesive forces) Low kinetic energy (energy of motion)

States of Matter. The Solid State. Particles are tightly packed, very close together (strong cohesive forces) Low kinetic energy (energy of motion) States of Matter The Solid State Particles are tightly packed, very close together (strong cohesive forces) Low kinetic energy (energy of motion) Fixed shape and volume Crystalline or amorphous structure