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Energy And States of Matter

Energy And States of Matter. Unit 10. Energy is… the capacity to do work Work is moving an object over a distance by applying a force. Energy is generally given off or absorbed during a chemical reaction.

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Energy And States of Matter

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  1. Energy And States of Matter Unit 10

  2. Energy is… • the capacity to do work • Work is moving an object over a distance by applying a force. • Energy is generally given off or absorbed during a chemical reaction. • Energy can change form, but it can not be created or destroyed in any chemical reaction or physical change (Law of Conservation of Energy) Energy and Energy Changes

  3. This theory (KMT) explains the effects of temp. and pressure on matter through 3 basic assumptions: • All matter is composed of small particles. • These particles are in constant random motion. • All collisions are perfectly elastic. • there is no change in the total kinetic energy of the 2 particles before and after the collision. • assumed that the gas particles attraction for each other is negligible • The motion of the particles vary w/ changes in temp., # of particles, and particle mass. • See it here. Kinetic-Molecular Theory

  4. Kinetic Energy – energy of motion, doing work Potential Energy – stored energy, not doing work yet FORMS Chemical Nuclear Electrical Heat Radiation (light) Sound Mechanical Electrical Types of Energy

  5. When dealing with changes in energy, there are only 2 parts to the universe: • The SYSTEM – reactants and products in a chemical process. 2. The SURROUNDINGS – everything else. Endothermic vs. Exothermic

  6. When the system loses energy to the surroundings during the change – EXOTHERMIC Ex. – combustion, Zn + S, hand warmers • When the system needs to draw energy from the surroundings for the change to take place – ENDOTHERMIC Ex. – decomposition of CaCO3(s), ice cube melting (a physical change)

  7. Temperature is the measure of how hot or how cold an object is. • determined by the avg. KE of molecules in sample area • temperature measured in K or oC • Heat is energy transferred from 1 object to another as a result of a difference in temperature. • temperature determines the direction of heat flow • object with high KE (higher temp) will lose energy to object with lower KE (lower temp) Heat vs. Temperature

  8. Energy is measured in Joules (J) or calories (cal). • One calorie is the amount of energy needed to raise the temperature of 1 ml of water by 1°C • kcal = energy needed to raise the temperature of 1000 ml of water 1°C • joule • 4.184 J = 1 cal • In nutrition, calories are capitalized • 1 Cal = 1 kcal = 1000 cal Units of Energy

  9. 1. How many J are in 553 cal? 2. How many calories are in 690.3 J? How many Calories? 3. A package of M&M’s (1.74 oz.) has 250. Cal, how many J of potential energy does that equal? Examples 4. It takes 1.64 x 106 J to run 1.5 miles. How many Calories would you need to consume to run that far?

  10. 1. 553 cal x 4.184 J/1 cal = 2310 J • 2. 690.3 J x 1 cal/4.184 J = 165.0 cal; 165.0 cal x 1 Cal/1000 cal = 0.1650 Cal • 3. 250 Cal x 1000 cal/1 cal x 4.184 J/1 cal = 1.05 x 106 J • 4. 1.64 x 106 J x 1 cal/4.184 J x 1 Cal/1000 cal = 392 Cal

  11. The amount the temperature of an object increases depends on the amount of heat added. • If you double the added heat energy the temperature will increase twice as much. • The amount the temperature of an object increases also depends on its mass. • If you double the mass it will take twice as much heat energy to raise the temperature the same amount. Energy and the Temperature of Matter

  12. To change the state of matter, energy must be added or removed. ENDOTHERMIC PROCESSES • Solid to a Liquid  Melting (Fusion) • particles overcome IMF and move around & past other particles • Solid to a Gas  Sublimation • occurs only at conditions far from normal MP • See it happen HERE! Energy Requirements for State Changes

  13. Liquid to a Gas  Vaporization • particles are very spread out – requires a lot of energy • evaporation – vaporization at the surface of a liquid EXOTHERMIC PROCESSES • Gas to a Liquid  Condensation (equal and opposite of vaporization) • Liquid to a solid  Solidification (equal and opposite of melting) Energy Requirements for State Changes

  14. Vapor Pressure & Dynamic Equilibrium • Vapor pressure is the pressure exerted by a vapor above the liquid • Vapor pressure increases with temperature • Dynamic equilibrium refers to the point at which the rate of evaporation and rate of condensation are equal. Equilibrium Liquid just poured into open container, little vapor Evaporation faster than Condensation Evaporation as fast as Condensation

  15. Boiling occurs when a liquid turns to a gas inside the liquid • bubbles are produced • Liquid boils when its Vapor Pressure = Atmospheric Pressure • Normal boiling point (1 atm. of pressure) • Large IMF = lower vapor pressure = high BP • Weak IMF = high vapor pressure = lower BP Boiling Point

  16. We use these diagrams to relate the process that occur when a substance changes from one phase to another. • Substances are in the following states when in certain locations on the diagram: • Solid – left side of diagram • Liquid – middle of diagram • Gas/Vapor – right side of diagram • When either the temp or pressure is changed, you can identify the process that is taking place and identify the phase change. • Ex (from diagram on last slide) – At 1 atm if you increase the temperature from 90oC to 200oC, the process you are undergoing is vaporization or boiling (liquid to gas). Phase Change Diagram

  17. The change of state occurs right on the equilibrium line. • Triple point identifies the conditions when you have all 3 states in dynamic equilibrium with one another. • Tm  normal melting point • The point at 1 atm or 101.3 kPa when solid turns to liquid. • Tb  normal boiling point • The point at 1 atm or 101.3 kPa when a liquid turns to a vapor • Critical point – you are no longer able to distinguish between gas and liquid phases past this point. Phase Change Diagram

  18. Gases

  19. Gases are composed of tiny particles • The particles are small compared to the average space between them – ie they are far apart • Compressible, fluid, expandable, and take the shape of their container • Particles constantly and rapidly moving in a straight line until they bump into each other or the wall – MEAN FREE PATH • Average kinetic energy proportional to the temperature • Results in gas pressure Describing a Gas

  20. Pressure = total force applied to a certain area • larger force = larger pressure • smaller area = larger pressure • Gas pressure caused by gas particles colliding with container or surface (the motion of the particles) • More forceful collisions or more frequent collisions mean higher gas pressure Pressure

  21. P = F/A  N/m2  1 Pa • 1 atm = 760. mm Hg = 760. torr = 29.92 in Hg = 101,325 Pa = 101.325 kPa = 14.7 psi • Calculate the number of mmHg in 1.42 atm • Calculate the number of kPa in 1110 Torr. Units of Gas Pressure

  22. Measuring Pressure of a Trapped Gas • Manometer – device used to measure gas pressure • Open-armed manometer • if gas end lower than open end, Pgas = Pair + diff. in height of Hg • if gas end higher than open end, Pgas = Pair – diff. in height of Hg • Closed-armed manometer • Pgas = difference in height of mercury • Barometer – special closed-armed manometer designed to measure air pressure. • Developed by Evangelista Torricelli (1640’s)

  23. Gas Relationships • #1 Pressure is inversely proportional to Volume • P1 x V1 = P2 x V2 • Boyles Law • #2 Volume is directly proportional to Temperature • V1 = V2 T1 T2 • Charles’ Law • #3 Pressure is directly proportional to the temperature. • P1 = P2 T1 T2 • Gay-Lussac’s Law • Temp. must be in K! • Gas properties

  24. It’s kind of a pain to remember three separate laws…and why should we when we can combine them together and only remember one? • That’s exactly what the combined gas law is…It can be used to solve simple gas law problems, as well as some more complex problems as well. • If a particular pair of variables are not used in the problem drop them out. This leaves one of the original 3 formulas behind. • Temperature must still be in K Next Up: Combine them!

  25. If nitrogen occupies a volume of 563 mL at 701 torr and 25oC, what volume will it have at 857 torr and 15oC? • A weather balloon contains 325 dm3 of helium at 15oC and 90.0 kPa. As the balloon rises into the atmosphere, what volume will it occupy at 70.0 kPa and -30.oC? • If the balloon from the previous problem continues to rise, it will eventually burst when it reaches a volume of 425 dm3. If its pressure at this point is 50.0 kPa, what will its temperature be? Combined Gas Law - Examples

  26. Volume directly proportional to the number of gas molecules • V = constant x n (moles) • Constant P and T • More gas molecules = larger volume • Count number of gas molecules by moles • One mole of any ideal gas occupies 22.4 L at standard conditions - molar volume • Equal volumes of gases contain equal numbers of molecules • It doesn’t matter what the gas is! Avogadro’s Law

  27. PV = nRT • P = pressure in atm • V = volume in liters • T = temperature in Kelvins • n = moles • R = a constant = 0.0821 L atm/ mol·K Holds closely at high temp (above 0oC) and low pressure (below 1 atm). Ideal Gas Law

  28. 1. What volume will 3.15 mol of hydrogen gas occupy at STP? 2. What pressure is exerted by 1.24 moles of a gas stored in a 15.0 L container at 22oC? 3. If you wish to collect 10.0 L of nitrogen gas at 15oC and 919 mm Hg, what mass of gas will you need to have? 4. What temperature (in oC) will you need to produce 515 cm3 of ammonia gas, if you had 14.3g of gas at a pressure of 117.2 kPa? Examples – Ideal Gas Law

  29. 5. What is the volume of a gas bottle that holds 0.225 moles of a gas stored at -45oC and 3.88 atm pressure? 6. What volume of chlorine is collected at 65oC and 74.5 cm Hg pressure when 93.8 grams of aluminum chloride is decomposed by heating? 2AlCl3 2Al + 3Cl2 Examples – Ideal Gas Law

  30. Diffusion and Effusion • As a result of the constant and rapid motion of gas particles, 2 things can occur: • Diffusion – particles move from an area of higher particle concentration to an area of lower particle concentration. 2. Effusion – same as diffusion except the gas particles escape through small holes (pores) in a container.

  31. Think About It 1. Rank the following gases in terms of their velocities, assuming they are at the same temperature (1-highest, 4-lowest). CO2, O2, N2, H2 2. Explain the similarities and differences between effusion and diffusion. 3. Rank the following gases based on the appropriateness for inflating a balloon you want to keep inflated (1-best to use, 4-worst to use). CO2, O2, N2, H2

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