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States of Matter & Gas Laws

States of Matter & Gas Laws. Review of the Properties of Solids, Liquids, and Gases Kinetic Molecular Theory of Gases Boyle’s Law, Charles’s Law, Gay-Lussac’s Law and the Combined Gas Law Dalton’s Law of Partial Pressure, Avogadro’s Law and the Ideal Gas Law

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States of Matter & Gas Laws

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  1. States of Matter & Gas Laws • Review of the Properties of Solids, Liquids, and Gases • Kinetic Molecular Theory of Gases • Boyle’s Law, Charles’s Law, Gay-Lussac’s Law and the Combined Gas Law • Dalton’s Law of Partial Pressure, Avogadro’s Law and the Ideal Gas Law • Phase Changes and Heating and Cooling Curves • Phase Diagrams

  2. Properties of Solids 1. The particles in a solid are packed closely together. 2. The particles in a solid are held together by strong intermolecular forces of attraction. 3. The particles in a solid vibrate slowly in place.

  3. Properties of Solids 4. Solids have a definite shape and a definite volume. 5. Solids are incompressible. 6. Solids are relatively dense.

  4. Types of Solids Crystalline solids have particles which are arranged in an orderly, geometric, repeating pattern. Example: NaCl Solids may be crystalline or amorphous. Amorphous solids contain particles which are arranged randomly. Examples: wax, glass, plastics

  5. Properties of Liquids 1. The particles in a liquid are not as close together or arranged as orderly as the particles in a solid. • The intermolecular forces of attraction between the particles of a liquid are not as strong as those found in solids. • 3. Liquids are fluid. The particles in a liquid are able to slide past each other.

  6. Properties of Liquids 4.Liquids have a definite volume but they do not have a definite shape. 5. Liquids are virtually incompressible. 6. Liquids tend to be more dense than gases but less dense than solids.

  7. Properties of Gases 1. The particles in a gas are far apart . 2. The particles in a gas have very weak intermolecular forces of attraction between them. 3. Gases are fluid. The particles in a gas are free to move in all directions.

  8. Properties of Gases Diffusion is the movement of gas particles from an area of high concentration to an area of low concentration. Ex. Opening a bottle of perfume, air fresheners. Effusion is the escape of gas particles through the pores or tiny pinholes in a container. Ex. A balloon deflating over time.

  9. Properties of Gases 4.Gases do not have a definite shape or volume. Gases expand to completely fill any container in which they are enclosed, and they take its shape. 5. Gases are compressible. The volume of a gas may be greatly decreased. 6. Gases have low densities. This is because there are not as many particles in a large space.

  10. The Kinetic Molecular Theory The kinetic molecular theory is based on the idea that particles of matter are always in motion. It can be used to explain the properties of solids, liquids, and gases in terms of the energy of particles and the forces that act between them.

  11. The Kinetic Molecular Theory of Gases The kinetic molecular theory provides a model of what is called an ideal gas. An ideal gas is an imaginary gas that perfectly fits all the assumptions of the kinetic molecular theory.

  12. Assumptions of The Kinetic – Molecular Theory of Gases 1. Gases consist of large numbers of tiny particles that are in continuous, rapid, random, straight-line motion and are far apart relative to their size. 2. The particles of an ideal gas are said to occupy zero volume and are dimensionless points. 3.Collisions between gas particles and between the particles and container walls are elastic. An elastic collision is one in which there is no net loss of kinetic energy.

  13. Assumptions of The Kinetic – Molecular Theory of Gases 4. There are no forces of attraction or repulsion between the particles of an ideal gas.

  14. Assumptions of The Kinetic Molecular Theory of Gases 5. The Kelvin temperature of a substance is directly proportion to the average kinetic energy of the particles of the substance. What happens to the kinetic energy of a gas particle as temperature increases? Kinetic energy increases decreases? Kinetic energy decreases

  15. Comparison of Ideal Gases to Real Gases Ideal gas particles have zero volume. Real gas particles occupy a small volume. The particles of an ideal gas do not exert attractive or repulsive forces on each other. Real gas particles experience weak forces of attraction and repulsion between the particles. 3. Real gases can be liquefied and sometimes solidified by cooling and by applying pressure. Ideal gases cannot be liquefied or solidified.

  16. Deviation from Ideal Gas Behavior Under certain conditions real gases can behave like ideal gases. These conditions include: High temperatures Low Pressures

  17. Deviation from Ideal Gas Behavior Polar molecules deviate more from ideal gas behavior than nonpolar molecules. This is because the intermolecular forces of attraction are usually stronger in polar molecules. Large molecules deviate more from ideal gas behavior than small diatomic molecules. This is because the particles take up more volume.

  18. Measuring Gases There are four main properties or variables of gases that are commonly measured: Volume, Temperature, Pressure, and Amount. Liters (L) Amount of space taken up by the substance 1000 cm3 = 1000 mL = 1 L Celsius;K = °C + 273 (Reminder: Standard Temperature is 0°C) Kelvin (K) Average KE of the particles 1 atm = 760 mm Hg 1 atm = 760 torr 1 atm = 101.3 kPa (Reminder: Standard Pressure is 1 atm) The force per unit area exerted by the collisions of gas particles. Gas pressure is measured with a barometer. Kilopascals (kPa) No other unit; can be converted to grams using molar mass Moles (mol) The number of particles present

  19. Boyle’s Law Robert Boyle (1627-1691), a British physicist and chemist, carefully investigated the compressibility of gases. Boyle’s Law states that at constant temperature, the volume of a gas varies inversely with pressure.

  20. Charles’s Law Jacques Charles (1746-1823), a French physicist, established by experimental measurement that the volume of a gas under constant pressure changes by 1/273 of its volume at 0°C for each degree change in temperature. Charles’s Law states that at constant pressure, the volume of a given mass of gas is directly proportional to its Kelvin temperature.

  21. Gay-Lussac’s Law J.L. Gay-Lussac (1778-1850), a French chemist and physicist, performed experiments to determine the relationship between the temperature and pressure exerted by a confined gas when volume is held constant. Gay-Lussac’s Law states that the pressure of a fixed mass of gas is directly proportional to the Kelvin temperature if the volume is kept constant.

  22. Combined Gas Law Boyle’s Law, Charles’s Law and Gay-Lussac’s Law are often combined into one mathematical statement.

  23. 1. A given mass of gas occupies a volume of 12 L when under a pressure of 2.0 atm. Assuming no change in temperature, what will the volume be if the pressure is changed to 3.0 atm?

  24. 2. A given mass of gas occupies a volume of 640 mL at 47°C and 650 mm Hg. What will the new temperature be if the volume is decreased to 210 mL and the pressure remains unchanged?

  25. 3. A quantity of acetylene in a steel cylinder is under a pressure of 1.2 atm at 27°C. What will be the pressure of the gas if the temperature is increased to 87°C?

  26. 4. A given mass of gas occupies a volume of 1.2 L at a temperature of 27°C and a pressure of 1.0 atm. What will the volume be if the temperature is increased to 77°C and the pressure remains unchanged?

  27. 5. The volume of a sample of gas is 200. mL at 275 K and 92.1 kPa. What will the new volume of the gas be at 350. K and 98.5 kPa?

  28. 6. A sample of nitrogen gas has a pressure of 2.5 atm at a temperature of 25°C. What temperature is required to increase the pressure to 4.0 atm, assuming that the volume is fixed and the amount of gas does not change?

  29. 7. 400. mL of a gas at -23°C and a pressure of 750. mm Hg will occupy what volume if the temperature is increased to 77°C and the pressure is decreased to 700. mm Hg?

  30. Dalton’ Law of Partial Pressure John Dalton (1766-1844), an English chemist and physicist, established that when two or more gases occupy the same container, each acts independently of the other(s) and exerts the same pressure that it would if it were alone in the container.

  31. Dalton’s Law of Partial Pressure Dalton’s law of partial pressure states that the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the individual gases in the mixture (at constant temperature). Ptotal = P1 + P2 + P3 …

  32. 1. Air contains oxygen, nitrogen, carbon dioxide, and trace amounts of other gases. What is the partial pressure of oxygen (PO2) at 760 mm Hg of pressure if PN2 = 593.4 mm Hg, PCO2= 0.3 mmHg, Pother= 7.1 mm Hg? Ptotal = PO2 + PN2 + PCO2 + Pother 760 mm Hg = PO2 + 593.4 mm Hg + 0.3 mm Hg + 7.1 mm Hg PO2= 760 mm Hg - (593.4 mm Hg + 0.3 mm Hg + 7.1 mm Hg) PO2= 159.2 mm Hg

  33. Collecting Gases over Water One of the most common ways of collecting a gas sample experimentally is by having it bubble into a jar filled with water. The collected gas is a mixture of the gas plus water vapor. You can use Dalton’s Law of Partial Pressures to correct for the partial pressure of water vapor, by subtracting the pressure of the water vapor from the total pressure.

  34. Collecting Gases over Water The amount of water vapor present is very sensitive to the temperature. The warmer the water is, the more water vapor present. In order to determine the vapor pressure of the dry gas, you will need to know the vapor pressure of the water. This will either be given to you or you will have to use a reference table to find it.

  35. 2. 0.750 L of a gas is collected over water at 25.0°C with a total pressure of 747.8 mm Hg. What is the pressure of the dry gas? Pgas =747.8 mm Hg–23.8 mm Hg Pgas= 724.0 mm Hg

  36. Avogadro’s Law Avogadro’s Law states that at the same conditions of temperature and pressure, equal volumes of gases contain the same number of particles. This means that the volume of a gas is directly proportional to the number of moles of the gas. Remember: At STP, one mole of any gas contains 6.02×1023 particles and occupies a volume of 22.4 L.

  37. Avogadro’s Law - Example Two sealed flasks of equal volume are at the same temperature and pressure. The first flask contains hydrogen gas. The second flask contains oxygen gas. 1. Which flask contains the most particles? 2. Which flask contains the most mass? The flask with the oxygen has the most mass. They contain the same number of particles.

  38. The Ideal Gas Law The ideal gas law equation allows for the quantity of the gas to be considered. The ideal gas law is as follows: PV = nRT Where: V (Volume) is expressed in Liters (1 L = 1000 mL) T (Temperature) is expressed in Kelvin (K = °C + 273) P (Pressure) can be expressed in atm, kPa, mm Hg or torr R = Ideal or Universal Gas Constant

  39. The Ideal Gas Law There are 3 different values for R, depending upon the units for pressure. If the pressure is given in atm. If the pressure is given in mm Hg. If the pressure is given in kPa.

  40. 1. What volume will 1.2 mol of N2 occupy at 17°C and 2.0 atm?

  41. 2. What pressure in atmospheres, will 0.80 mol of gas in a volume of 2.0 L exert if the temperature is 27°C?

  42. 3. At what Celsius temperature will 0.20 mol of gas in a 1.5 L cylinder exert a pressure of 2.8 atm? T = 256 - 273 = -17°C

  43. 4. At 27°C the gas in a 920 mL flask exerts a pressure of 730 mm Hg. How many grams of nitrogen gas are in the flask?

  44. Review Concept 1. Nitrogen gas and hydrogen gas react to produce ammonia. N2+ 3H2 → 2NH3 What volume of ammonia will be produced at STP if 20.0 g of hydrogen gas react with excess nitrogen gas?

  45. Phase Changes melting freezing liquid  solid solid  liquid melting freezing when a substance reaches its freezing point when a substance reaches its melting point water freezing Ice cube melting

  46. Phase Changes vaporization condensation liquid  gas gas  liquid vaporization condensation when the liquid vapor pressure is equal to the atmospheric pressure (boiling) through vaporization at the surfaceof a liquid (evaporation) when a gaseous substance comes into contact with acoolsurface blowing your hair dry water on the outside of a cup water boiling

  47. Phase Changes sublimation deposition gas  solid solid  gas sublimation deposition when a substance goes directlyfrom a solid to a gas when a substance goes directlyfrom a gas to a solid dry ice and solid air fresheners frost in the freezer It is important to note that the temperature of a substance does not change during a phase change.

  48. Phase Changes Phase changes (or transitions) occur when a material changes from one phase or state to another. The six most common phase changes are: Which of the phase changes represent endothermic processes? Melting, vaporization and sublimation Which of the phase changes represent exothermic processes? Freezing, condensing, and deposition

  49. Vapor Pressure Curves A vapor pressure curve can be used to represent the boiling point of a substance at various vapor pressures. Any point along a vapor pressure curve for a compound represents the boiling point of the substance. Thenormal boiling pointof a substance occurs when the vapor pressure is equal to 1 atm (760 mmHg, 101.3 kPa).

  50. Vapor Pressure Curves At what temperature will ethanol boil when the atmospheric pressure is 200 mm Hg? What vapor pressure would be needed to make water boil at 80°C? Approximately 48°C 333 torr

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