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Solids, Liquids, & Gases Chapter 7

This chapter explores the structure and properties of gases, including their composition, movement, compressibility, density, and pressure. It also covers gas laws such as Boyle's Law and Charles' Law.

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Solids, Liquids, & Gases Chapter 7

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  1. Solids, Liquids, & GasesChapter 7

  2. The Structure of a Gas • Gases are composed of particles that are flying around very fast in their container(s). • They move in straight lines until they encounter either the container wall or another particle, then they bounce off. • In gases, there is a lot of empty space. Tro's Introductory Chemistry, Chapter 11

  3. Pressure • Is caused by the collisions of molecules with the walls of a container • Is equal to force/unit area • P=F/A • SI units = Newton/meter2 = 1 Pascal (Pa) • 101,325 Pa = 1 standard atmosphere (1atm) • 1 atm = 760mmHg = 760 torr = 14.7 psi

  4. Properties—Indefinite Shape and Indefinite Volume Because the gas molecules have enough kinetic energy to overcome attractions, they keep moving around and spreading out until they fill the container. As a result, gases take the shape and the volume of the container they are in. Tro's Introductory Chemistry, Chapter 11

  5. Properties—Compressibility Because there is a lot of unoccupied space in the structure of a gas, the gas molecules can be squeezed closer together. Tro's Introductory Chemistry, Chapter 11

  6. Gas Properties Explained—Low Density Because there is a lot of unoccupied space in the structure of a gas, gases do not have a lot of mass in a given volume, the result is that they have low density. Tro's Introductory Chemistry, Chapter 11

  7. The Pressure of a Gas • Pressure is the result of the constant movement of the gas molecules and their collisions with the surfaces around them. • The pressure of a gas depends on several factors: • Number of gas particles in a given volume. • Volume of the container. • Average speed of the gas particles. Tro's Introductory Chemistry, Chapter 11

  8. Density and Pressure • Pressure is the result of the constant movement of the gas molecules and their collisions with the surfaces around them. • When more molecules are added, more molecules hit the container at any one instant, resulting in higher pressure. • Also higher density. Tro's Introductory Chemistry, Chapter 11

  9. Air Pressure • The atmosphere exerts a pressure on everything it contacts. • On average 14.7 psi. • The atmosphere goes up about 370 miles, but 80% is in the first 10 miles from Earth’s surface. • This is the same pressure that a column of water would exert if it were about 10.3 m high. Tro's Introductory Chemistry, Chapter 11

  10. Measuring Pressure • The first device for measuring atmospheric pressure was developed by Evangelista Torricelli during the 17th century. • The device was called a “barometer” • Baro = weight • Meter = measure

  11. An Early Barometer • The normal pressure due to the atmosphere at sea level can support a column of mercury that is 760 mm high.

  12. The Aneroid Barometer

  13. The Digital Barometer

  14. Common Units of Pressure Tro's Introductory Chemistry, Chapter 11

  15. GAS LAWS Boyle’s Law Charles’ Law Gay-Lussac’s Law

  16. The Ideal Gas Law is: PV=nRT

  17. Boyle’s Law • Pressure of a gas is inversely proportional to its volume. • Constant T and amount of gas. • Graph P vs. V is curved. • Graph P vs. 1/V is in a straight line. • As P increases, V decreases by the same factor. • P x V = constant. • P1 x V1 = P2 x V2. Tro's Introductory Chemistry, Chapter 11

  18. When you double the pressure on a gas, the volume is cut in half (as long as the temperature and amount of gas do not change). Tro's Introductory Chemistry, Chapter 11

  19. Gas Laws Explained— Boyle’s Law • Boyle’s law says that the volume of a gas is inversely proportional to the pressure. • Decreasing the volume forces the molecules into a smaller space. • More molecules will collide with the container at any one instant, increasing the pressure. Tro's Introductory Chemistry, Chapter 11

  20. Boyle’s Law and Diving Scuba tanks have a regulator so that the air from the tank is delivered at the same pressure as the water surrounding you. This allows you to take in air even when the outside pressure is large. • Since water is more dense than air, for each 10 m you dive below the surface, the pressure on your lungs increases 1 atm. • At 20 m the total pressure is 3 atm. • If your tank contained air at 1 atm of pressure, you would not be able to inhale it into your lungs. • You can only generate enough force to overcome about 1.06 atm. Tro's Introductory Chemistry, Chapter 11

  21. Boyle’s Law and Diving • If a diver holds her breath and rises to the surface quickly, the outside pressure drops to 1 atm. • According to Boyle’s law, what should happen to the volume of air in the lungs? • Since the pressure is decreasing by a factor of 3, the volume will expand by a factor of 3, causing damage to internal organs. Always Exhale When Rising!! Tro's Introductory Chemistry, Chapter 11

  22. The Relationship Between Temperature and Volume

  23. Jaques Charles (1746-1823) • Charles studied the compressibility of gases nearly a century after Boyle • French Physicist • Conducted the first scientific balloon flight in 1783

  24. Volume and Temperature As a gas is heated, it expands. This causes the density of the gas to decrease. Because the hot air in the balloon is less dense than the surrounding air, it rises. Tro's Introductory Chemistry, Chapter 11

  25. Charles’ Law • At a fixed pressure, the volume of a gas is proportional to the temperature of the gas. • As the temperature increases, the volume increases. • As the temperature decreases, the volume decreases.

  26. Charles’s Law • Volume is directly proportional to temperature. • Constant P and amount of gas. • Graph of V vs. T is a straight line. • As T increases, V also increases. • Kelvin T = Celsius T + 273. • V = constant x T. • If T is measured in kelvin. Tro's Introductory Chemistry, Chapter 11

  27. Charles’ Equation • Rearrange the equation to solve for V1, T1, V2, or T2 • V1=(V2T1)/T2 • T1=(V1T2)/V2 • V2=(V1T2)/T1 • T2=(V2T1)/V1 • All temperatures must be in KELVIN!!!

  28. Gay-Lussac’s Law The Relationship Between Pressure and Temperature

  29. Joseph Louis Gay-Lussac1778 - 1850 • French chemist and physicist • Known for his studies on the physical properties of gases. • In 1804 he made balloon ascensions to study magnetic forces and to observe the composition and temperature of the air at different altitudes

  30. Gay-Lussac’s Law • The pressure of a fixed amount of gas at fixed volume is directly proportional to its temperature in Kelvin. • As the temperature increases, the pressure also increases • As the temperature decreases, the pressure also decreases

  31. Gay-Lussac’s Law • Expressed Mathematically as: • Rearranging this equation to solve for the variables gives us: • P1=(P2T1)/T2 • P2=(P1T2)/T1 • T1=(P1T2)/P2 • T2=(P2T1)/P1

  32. 7.3 Gas LawsD. The Combined Gas Law • All three gas laws can be combined into one equation: P1V1 P2V2 = T1 T2 initial conditions new conditions • This equation is used for determining the effect of changing two factors (e.g., P and T) on the third factor (V).

  33. The Combined Gas Law • Boyle’s law shows the relationship between pressure and volume. • At constant temperature. • Charles’s law shows the relationship between volume and absolute temperature. • At constant pressure. • The two laws can be combined together to give a law that predicts what happens to the volume of a sample of gas when both the pressure and temperature change. • As long as the amount of gas stays constant.

  34. The Combined Gas Law • A combination of Boyle’s, Charles’, and Gay-Lussac’s laws. • Written mathematically as: • Temperature must be in KELVIN.

  35. Avogadro’s LawThe Relationship Between Pressure, Temperature, Volume and Moles

  36. Amedeo Avogadro(1776-1856) • Italian physicist and mathematician • Born in a noble ancient family of Piedmont • “Avogadro’s Number” is named after him • 6.022 x 1023 • The number of things in a mole

  37. Avogadro’s Law • One mole of any gas occupies exactly 22.4 liters (dm3) at STP. • STP = Standard Temperature and Pressure • Temp = 0ºC = 273K • Pressure = 1atm = 760 mmHg = 760 torr etc. • This is often referred to as the “molar volume” of a gas.

  38. Equal volumes of gases, at the same temperature and pressure, contain the same number of particles, or molecules. • Thus, the number of molecules in a specific volume of gas is independent of the size or mass of the gas molecules. • Therefore, IT DOESN’T MATTER WHICH GAS YOU ARE TALKING ABOUT.

  39. Avogadro’s Law • Written mathematically: • Where “n” is the number of moles of gas present • Mini Version:

  40. The Ideal Gas Law

  41. The Ideal Gas Law is most often written as: PV=nRT • Rearranging the equation gives us: • P = (nRT)/V • V = (nRT)/P • n = (PV)/(RT) • T = (PV)/(nR)

  42. R is the Universal Gas Constant • R = 0.082 L x atm x K-1 x mol-1 • R = 62.36 L x torr x K-1 x mol-1 • R = 62.36 L x mmHg x K-1 x mol-1 • R = 8.315 L x kPa x K-1 x mol-1

  43. Molar Mass of a Gas • One of the methods chemists use to determine the molar mass of an unknown substance is to heat a weighed sample until it becomes a gas, measure the temperature, pressure, and volume, and use the ideal gas law. Tro's Introductory Chemistry, Chapter 11

  44. Dalton’s LawThe Law of Partial Pressures

  45. John Dalton(1766-1844) • English Chemist and Physicist • Did much research on color blindness • He was colorblind • Famous for his atomic theory • Worked with gases

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