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Chapter 6 States of Matter: Gases, Liquids, and Solids Denniston Topping Caret 4 th Edition Copyright  The McGraw-Hill Companies, Inc. Permission required.

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Presentation on theme: "Chapter 6 States of Matter: Gases, Liquids, and Solids Denniston Topping Caret 4 th Edition Copyright  The McGraw-Hill Companies, Inc. Permission required."— Presentation transcript:

1 Chapter 6 States of Matter: Gases, Liquids, and Solids Denniston Topping Caret 4 th Edition Copyright  The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2 6.1 The Gaseous State Ideal Gas Concept Ideal gas - a model of the way that particles of a gas behave at the microscopic level. We can measure the following of a gas: –temperature, –volume, –pressure and –quantity (mass) We can systematically change one of the properties and see the effect on each of the others.

3 6.1 The Gaseous State Measurement of Gases The most important gas laws involve the relationship between –number of moles (n) of gas –volume (V) –temperature (T) –pressure (P) Pressure - force per unit area. Gas pressure is a result of force exerted by the collision of particles with the walls of the container.

4 6.1 The Gaseous State Barometer - measures atmospheric pressure. –Invented by Evangelista Torricelli A commonly used unit of pressure is the atmosphere (atm). 1 atm is equal to:  760 mmHg  760 torr  76 cmHg

5 6.1 The Gaseous State Boyle’s Law Boyle’s Law - volume of a gas is inversely proportional to pressure if the temperature and number of moles is held constant. PV = k 1 or P i V i = P f V f 1

6 6.1 The Gaseous State

7 1.A 5.0 L sample of a gas at 25 o C and 3.0 atm is compressed at constant temperature to a volume of 1.0 L. What is the new pressure? 2.A 3.5 L sample of a gas at 1.0 atm is expanded at constant temperature until the pressure is 0.10 atm. What is the volume of the gas? 2

8 1 6.1 The Gaseous State Charles’ Law Charles’ Law - volume of a gas varies directly with the absolute temperature (K) if pressure and number of moles of gas are constant. or

9 6.1 The Gaseous State

10 1.A 2.5 L sample of gas at 25 o C is heated to 50 o C at constant pressure. Will the volume double? 2. What would be the volume in question 1? 3. What temperature would be required to double the volume in question 1? 2

11 6.1 The Gaseous State Combined Gas Law This law is used when a sample of gas undergoes change involving volume, pressure, and temperature simultaneously. 1

12 2 6.1 The Gaseous State Calculate the temperature when a 0.50 L sample of gas at 1.0 atm and 25 o C is compressed to 0.05 L of gas at 5.0 atm.

13 6.1 The Gaseous State Avogadro’s Law Avogadro’s Law - equal volumes of an ideal gas contain the same number of moles if measured under the same conditions of temperature and pressure. or 1

14 6.1 The Gaseous State Assuming no change in temperature and pressure, how many moles of gas would be needed to double the volume occupied by 0.50 moles of gas? 2

15 6.1 The Gaseous State Molar Volume of a Gas Molar Volume - the volume occupied by 1 mol of any gas. STP - Standard Temperature and Pressure  T = 273 K (or 0 o C)  P = 1 atm At STP the molar volume of a gas is 22.4 L –We will learn to calculate the volume later.

16 6.1 The Gaseous State Gas Densities We know: density = mass/volume Let’s calculate the density of H 2. What is the mass of 1 mol of H 2 ? 2.0 g What is the volume of 1 mol of H 2 ? 22.4 L Density = 2.0 g/22.4 L = g/L

17 6.1 The Gaseous State The Ideal Gas Law 1 Combining Boyle’s Law, Charles’ Law and Avogadro’s Law gives the Ideal Gas Law. PV=nRT R (ideal gas constant) = L. Atm/mol. K

18 6.1 The Gaseous State Let’s calculate the molar volume at STP using the ideal gas law: PV = nRT What would be the pressure?  1 atm What would be the temperature?  273 K What would be the number of moles?  1 mol 22.4 L 2

19 6.1 The Gaseous State 1.What is the volume of gas occupied by 5.0 g CH 4 at 25 o C and 1 atm? 2.What is the mass of N 2 required to occupy 3.0 L at 100 o C and 700 mmHg?

20 6.1 The Gaseous State Dalton’s Law of Partial Pressures 1 Dalton’s Law - a mixture of gases exerts a pressure that is the sum of the pressures that each gas would exert if it were present alone under the same conditions. P t =p 1 +p 2 +p For example, the total pressure of our atmosphere is equal to the sum of the pressures of N 2 and O 2.

21 6.1 The Gaseous State Kinetic Molecular Theory of Gases 3 1. Gases are made up of small atoms or molecules that are in constant and random motion. 2.The distance of separation is very large compared to the size of the atoms or molecules. –the gas is mostly empty space. Provides an explanation of the behavior of gases that we have studied in this chapter. Summary follows:

22 6.1 The Gaseous State 3.All gas particles behave independently. –No attractive or repulsive forces exist between them. 4.Gas particles collide with each other and with the walls of the container without losing energy. –The energy is transferred from one atom or molecule to another. 5.The average kinetic energy of the atoms or molecules is proportional to absolute temperature. –K.E. = 1/2mv 2 so as temperature goes up, the speed of the particles goes up.

23 6.1 The Gaseous State How does the Kinetic Molecular Theory of Gases explain the following statements? Gases are easily compressible. Gases will expand to fill any available volume. Gases have low density. 4

24 6.1 The Gaseous State Remember: pressure is a force per unit area resulting from collision of gas particles with the walls of the container. If pressure remains constant why does volume increase with temperature?

25 6.1 The Gaseous State Gases behave most ideally at low pressure and high temperatures. Ideal Gases Vs. Real Gases In reality there is no such thing as an ideal gas. Instead this is a useful model to explain gas behavior. Non-polar gases behave more ideally than polar gases because attractive forces are present in polar gases.

26 6.2 The Liquid State Liquids are practically incompressible. –Enables brake fluid to work in your car Viscosity - a measure of a liquids resistance to flow. –Flow occurs because the molecules can easily slide past each other. –Glycerol - example of a very viscous liquid. –Viscosity decreases with increased temperature. 5

27 6.2 The Liquid State Surface tension - a measure of the attractive forces exerted among molecules at the surface of a liquid. Surface molecules are surrounded and attracted by fewer liquid molecules than those below. Net attractive forces on surface molecules pull them downward. –Results in “beading” Surfactant - substance added which decreases the surface tension –example: soap

28 6.2 The Liquid State Vapor Pressure of a Liquid 6 What happens when you put water in a sealed container? Both liquid water and water vapor will exist in the container. How does this happen below the boiling point? Kinetic Theory - Liquid molecules are in continuous motion, with their average kinetic energy directly proportional to the Kelvin temperature.

29 6.2 The Liquid State energy + H 2 O(l)  H 2 O(g) The green line represents the minimum energy required to break the intermolecular attractions. Even at the cold temp, some molecules can be converted.

30 6.2 The Liquid State Once there are molecules in the vapor phase, they can be converted back to the liquid phase H 2 O(g)  H 2 O(l) + energy evaporation - the process of conversion of liquid to gas, at a temperature too low to boil condensation - conversion of the gas to the liquid state.

31 6.2 The Liquid State When the rate of evaporation equals the rate of condensation, the system is at equilibrium. Vapor pressure of a liquid - the pressure exerted by the vapor at equilibrium H 2 O(g) H 2 O(l)

32 6.2 The Liquid State Boiling point - the temperature at which the vapor pressure of the liquid becomes equal to the atmospheric pressure. Normal boiling point - temperature at which the vapor pressure of the liquid is equal to 1 atm. What happens when you go to a mountain where the atmospheric pressure is lower than 1 atm? The boiling point lowers.

33 6.2 The Liquid State Boiling point is dependant on the intermolecular forces –Polar molecules have higher b.p. than nonpolar molecules.

34 6.2 The Liquid State Van der Waals Forces 7 Van der Waals Forces are types of intermolecular forces. Consists of: 1Dipole-dipole interactions (section 4.5) 2London forces

35 London forces: –exist between all molecules, –is the only attractive force between nonpolar atoms or molecules, –and dipole-dipole attractions occur between polar molecules. Electrons are in constant motion. Electrons can be, in an instant, arranged in such a way that they have a dipole. (Instantaneous dipole) The temporary dipole interacts with other temporary dipoles to cause attraction. 6.2 The Liquid State

36 Hydrogen Bonding 8 Hydrogen bonding: –not considered a Van der Waals Force –is a special type of dipole-dipole attraction –is a very strong intermolecular attraction causing higher than expected b.p. and m.p. Requirement for hydrogen bonding: –molecules have hydrogen directly bonded to O, N, or F

37 6.2 The Liquid State Examples of hydrogen bonding: H 2 O NH 3 HF

38 6.3 The Solid State Particles highly organized and well defined fashion Fixed shape and volume Properties of Solids: –incompressible –m.p. depends on strength of attractive force between particles –Crystalline solid - regular repeating structure –Amorphous solid - no organized structure. 9

39 6.3 The Solid State Types of Crystalline Solids 1.Ionic Solids held together by electrostatic forces high m.p. and b.p. hard and brittle if dissolves in water, electrolytes NaCl 2.Covalent Solid Held together entirely by covalent bonds high m.p. and b.p. extremely hard Diamond

40 6.3 The Solid State 3.Molecular solids molecules are held together with intermolecular forces often soft low m.p. often volatile ice 4.Metallic solids metal atoms held together with metal bonds metal bonds –overlap of orbitals of metal atoms –overlap causes regions of high electron density where electrons are extremely mobile - conducts electricity

41 6.3 The Solid State

42 The End Chapter 6


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