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Unit 3: Phases of Matter-Key Regents Chemistry ’14-‘15 Mr Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Unit 3: Phases of Matter Student Name: _______________Key_________________ Class Period: ________ Page 1 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Page intentionally blank Page 2 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Unit 3 Vocabulary: 1. Avogadro’s Hypothesis: Equal volumes of two ideal gases under the same conditions of temperature and pressure will contain equal number of molecules. 2. Boiling: the transition of a liquid into a gas at the vaporization point. 3. Condensing: the transition of a gas into a liquid at the vaporization point. 4. Equilibrium: the condition that exists when the rates of two opposing phase changes are equal. 5. Evaporating: the transition of the surface molecules of a liquid into a gas below the boiling point. 6. Freezing: the transition of a liquid into a solid at the solidification point. 7. Gas: the phase of matter with complete dissociation of particles of matter with a great average distance between the particles compared to the particle size. Negligible attractive forces between particles. 8. Heat of Fusion: energy required to liquefy a gram of solid at melting point. 9. Heat of vaporization: energy required to gasify a gram of liquid at its vaporization point. 10. Ideal Gas: a gas in which the molecules are infinitely small and far apart, the molecules travel with straight-line motion, all collisions are elastic (no net energy loss), no attractive forces between gas molecules and speed of molecules is directly proportional to Kelvin. Gasses act most ideal at high temperature and low pressures. 11. Liquid: a phase of matter characterized as loosely organized yet held together by intermolecular or ionic attractive forces. 12. Melting: the transition of a solid into a liquid at the liquefaction point. 13. Pressure: Force exerted over an area. Page 3 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Unit 2 Vocabulary (Cont’d) 14. Solid: a phase with matter arranged in regular geometric patterns (lattices) with only vibrational motion, and no relative motion. 15. Specific Heat: the energy required to heat one gram of a substance one Kelvin. 16. Sublimation: the transition of a solid directly into a gas, skipping liquefaction. 17. Vapor Pressure: the pressure exerted by vapor in a vapor-liquid mixture closed system at equilibrium. 18. Vapor-Liquid Equilibrium: a system where the rate of evaporation equals the rate of condensing. Page 4 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Unit 3 Homework Assignments: Assignment: Date: Due: Page 5 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Phases and Phase Change Objective: How does closeness of matter affect their properties? Positive and negative charges attract each other. Attractive forces between molecules are called intermolecular attractive forces (IMAF). The strength of these forces determines what phase of matter a substance is in at a given temperature. Substances having weak attractive forces (London dispersion) between their molecules tend to be gasses at room temperature, and substances with strong attractive forces (ionic) tend to be solids at room temperature. Phases of matter are simply stages of attraction. Gases are composed of molecules with little or no attractive forces, allowing the molecules to fly freely past each other. Liquids are made of molecules with stronger attractive forces, allowing molecules to flow past each other, but still stay together. Solids are made of molecules or ions with strong attractive forces, which lock the molecules into a crystal lattice where the particles are free to vibrate, but they cannot move relative to each other. Page 6 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Properties of Solids Objective: How does closeness of solids affect their properties? Properties of Solids: 1. Molecules, atoms, or ions arranged into a regular, geometric pattern called a crystal lattice. 2. Molecules, atoms, or ions vibrate in place. They do NOT move relative to each other. 3. Solids have a definite shape and definite volume. Page 7 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Properties of Liquids Objective: How does closeness of liquids affect their properties? Properties of Liquids: 1. Molecules, atoms, or ions may flow past each other. 2. Viscosity is a resistance to flow due to intermolecular attractive forces (IMAF). 3. Viscosity increases as temperature decreases and IMAF strength increases. 4. Liquid molecules near a gaseous surface may escape IMAF and enter the vapor phase at temperatures below the boiling point in a process called evaporation. (This is the source for vapor pressure) 5. Liquids take on the shape of the container they are in and have a definite volume. (In the absence of gravity or a container, liquids form a perfect sphere.) Page 8 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Properties of Gases Objective: How does closeness of gases affect their properties? Properties of Gases: 1. Gas molecules are extremely far apart compared to the size of the molecules. 2. Gas molecules travel in a straight line until they collide with something. 3. Collisions are elastic, meaning they don’t lose kinetic energy in the collision. 4. Gas molecules move faster when it is hotter (higher Kelvin). 5. The gas phase is the only phase that is affected by changes in pressure. 6. Gases spread out to take the shape and entire volume of whatever container they are placed in. Page 9 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Phase Change Diagram Objective: How do the changes of phase relate to each other? Phase Change Diagram: Phase Symbols: Formula(s) = solid phase Formula(l) = liquid phase Formula(g) = gaseous phase Page 10 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Phase Equilibrium Objective: What processes are at work during phase equilibrium? Phase equilibrium: Boiling and condensing both occur at the boiling point (100.0˚C for pure water), freezing and melting both occur at the melting point (0.0 ˚C for pure water). During the phase change, both phases exist at equilibrium. Equilibrium: a condition where the rates of opposing changes are equal. At equilibrium temperature a substance at the melting/freezing point is melting at the SAME rate that it is freezing. Equilibrium is the reason why water melts ABOVE 0°C and freezes BELOW 0°C. AT 0°C a mixture of water and ice will NOT change either direction. If you have a sealed flask containing 20.0 g of water ice and 20.0 g of liquid water and keep it at EXACTLY 0.0°C, there will still 20.0 g of water ice and 20.0 g of liquid water for as long as the temperature remains 0.0°C. Page 11 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Sublimation Objective: What occurs during the solid to gaseous phase change? Sublimation: During sublimation the attractive forces between solid molecules are so weak that heating a solid causes it to go directly into the gas phase. There are two common substances that undergo this change; CO2(s), known as dry ice, and I2(s). Water also undergoes sublimation. Ice cubes ‘shrink’ if left alone in a freezer. The opposite of sublimation is the process of DEPOSITION. This can be witnessed by the intricate patterns of ice found on window surfaces on those cold, dry upstate New York mornings when water vapor in the air contacts the cold glass and the gaseous water phase changes directly into solid water. Page 12 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Notes page: Page 13 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Heating Curve Diagrams Objective: What steps occur during constant heating of a substance? Page 14 of 63 Website upload 2014 Unit 3: Phases of Matter-key Regents Chemistry ’14-‘15 Mr. Murdoch Topic: Heating Curve Diagrams Objective: What steps occur during constant heating of a substance? Phase Change Diagrams (Heating/Cooling Curve) Page 14: When applying CONSTANT HEAT to a solid (point A), the temperature of the solid increases (segment AB) until the melting point (point B) is reached, for example of water ice at 0˚C or 273 K. The solid then absorbs potential energy as it transitions from solid to liquid, or melts (segment BC). During the melting process, the temperatures (under CONSTANT HEAT) of both solid and liquid water BOTH remain constant until the ENTIRE solid has become liquid (point C). Once the substance is completely in the liquid phase, the temperature will increase (segment CD). Once the boiling temperature (point D) has been reached, for liquid water at 100˚C or 373 K, the temperatures (under CONSTANT HEAT) of both liquid water and water vapor will remain constant in a process called boiling, or vaporization (segment DE). Once ALL the water is in the gas phase (point E), the temperature (under CONSTANT HEAT) of the water vapor will increase (segment EF). Both phase changes are ENDOTHERMIC here, because heat is a CONSTANT application at the same RATE.
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