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Chapter 5 Gas Laws

by: Gabrielle Herman

Chapter 5 Gas Laws Chem 101

Gabrielle Herman

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These notes go over the gas laws of chemistry. They include all the formulas and the explanation of all of them.
General, Organic, and Biochemistry for the Health Sciences
Class Notes
General Chemistry, dalton's law kinetic molecular theory graham's law chemistry real gases ideal gas law partial pressures degrazzia
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This 5 page Class Notes was uploaded by Gabrielle Herman on Sunday October 2, 2016. The Class Notes belongs to Chem 101 at Ball State University taught by Khisamutdinov in Fall 2016. Since its upload, it has received 13 views. For similar materials see General, Organic, and Biochemistry for the Health Sciences in Chemistry at Ball State University.


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Date Created: 10/02/16
CHEM 101_Dr. Emil Lecture Notes and In-class problems. CHAPTER 5 NOTES OUTLINE: 5.1 Changes of State 5.2 The Gas Laws 5.3 Gas Mixtures and Partial Pressures 5.1 Changes of State Changes of state are the processes of going from one physical state to another physical state. A change of state is a physical change because covalent bonds are not formed or broken in the process. There are no intermolecular forces in the gas phase but intermolecular forces exist in the liquid phase, and are even more abundant in the solid phase.  A change of state from the solid to the liquid phase is melting; the reverse is freezing.  A change of state from liquid to gas is vaporization or evaporation; the reverse is condensation.  A solid can go directly to the gas phase, the process is known as sublimation; the reverse is deposition. During melting, vaporization, and sublimation, intermolecular forces are disrupted as kinetic energy increases. During condensation, freezing, and deposition, intermolecular forces are created as kinetic energy decreases. Energy must be added for melting, vaporization, and sublimation to occur. Energy must be removed for freezing, condensing and deposition to occur. A heating curve shows the temperature change of a substance as energy is added. Temperature increases linearly except during phase changes: at these points, the energy is being used to disrupt intermolecular forces. The first plateau is the freezing point. The second plateau is the boiling point. The amount of energy required to melt a solid at its melting point is the heat of fusion or enthalpy of fusion: H .fus The amount of energy required to vaporize a liquid at its boiling point is the heat of vapori- zation or enthalpy of vaporization: H . vap A cooling curve is an inverted heating curve. Steam burns are so severe because the H vap is transferred into the skin. Vapor pressure is the pressure exerted by a gas in the liquid-gas equilibrium. Pressure (P) is the force per unit area exerted by gas particles colliding against the walls of their container. A liquid is volatile if it has a high vapor pressure; its molecules enter the gas phase readily. The vapor pressure of a substance increases with temperature. The normal boiling point of a liquid is the temperature at which the vapor pressure is 1 atm. Molecules in the air exert a downward pressure as a result of gravity. This is atmospheric pressure or barometric pressure. It is measured with a barometer. At higher altitudes, atmospheric pressure decreases. Hikers must acclimate to lower air pressure and less oxygen per breath. Deep-sea divers experience increases in pressure, and also must acclimate to lower pressures as they return to the surface. 5.2 The Gas Laws The kinetic-molecular view of a gas states: Particles of a gas are in constant, random motion. The volume of the gas particles is negligible. The attractive forces between the particles of a gas are negligible. The temperature of a gas depends on the average kinetic energy of the gas particles. The macroscopic properties of a gas are described by four interrelated variables: • pressure (P) • volume (V) • temperature (T) • number of moles (n) Changing one affects the others in a predictable way. As the volume of a gas decreases, the pressure of the gas increases, and as the volume of a gas increases, the pressure decreases if temperature and amount remain constant. Boyle’s law : Pressure and volume are inversely proportional: as one goes up the other goes down. It follows that it is possible to determine pressure or volume variables after a change with the following equation: Charles’s law: As the temperature of a gas increases, the kinetic energy of the particles increases, and the volume increases, if pressure and amount are constant. Volume and temperature of a gas are directly proportional. This law can be demonstrated by heating a balloon and observing its volume increase or cooling the balloon and observing its volume decrease. Avogadro’s law: As the amount of a gas increases, the volume increases, if pressure and temperature are constant. Moles and volume of a gas are directly proportional. Ideal gas law: Volume is affected by changes in temperature, pressure, or number of moles: The proportionality can be converted to an equality with the use of a constant, R: R is the universal gas constant, and is the same for all ideal gases: From the ideal gas law, it is possible to calculate the molar volume of a gas at standard temperature and pressure (0°C and 1 atm). The molar volume of any gas at STP is 22.4 L. If the molar volume of a gas is 22.4 L, the density of any gas is: A gas lighter than air will cause a balloon to float Dalton’s law of partial pressures: in a mixture of gases, each gas exerts a pressure independent of the other gases, and each gas will behave as if it alone occupied the total volume. Henry’s law: the number of gas molecules dissolved in a liquid is directly proportional to the partial pressure of the gas. The smaller the Henry’s constant (k) is, the higher the concentration in the water.


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