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Temperature, Heat, & Combustion

Temperature, Heat, & Combustion. EGR 1301: Introduction to Engineering. Models. “A system of postulates, data, and inferences presented as a mathematical description of an entity or state of affairs” Quantitative approximation of reality Mathematical equations Computer simulations

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Temperature, Heat, & Combustion

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  1. Temperature, Heat,& Combustion EGR 1301: Introduction to Engineering

  2. Models • “A system of postulates, data, and inferences presented as a mathematical description of an entity or state of affairs” • Quantitative approximation of reality • Mathematical equations • Computer simulations • Physical scale models • Why do we use them? • Reality is too complex!!! Source: Merriam-Webster .com, 2010

  3. Energy Conversion Tables “For those who want some proof that physicists are human, the proof is in the idiocy of all the different units which they use for measuring energy.” Richard Feynman Source: Foundations of Engineering, Holtzapple & Reece, 2003

  4. Incandescent bulb Resistance heating in filament  Light When filament reaches sufficiently high temperature  Light is radiated 60 Watts of electricity 800 lumens of light ~10 cals of heat Fluorescent bulb Stream of electrons collide with Hg electronsLight Generates very little heat 23 Watts of electricity 800 lumens of light ~1 cal of heat Utility of Energy for Analysis Source: http://www.fullspectrumsolutions.com/26w_powercompact_65_prd1.htm

  5. Temperature – What Is It? • NOT the same as HEAT • A quantitative measure of “hotness” • More accurately described on an atomic scale • Measures vibrational kinetic energy

  6. Temperature vs. Heat • Temperature • A measure of the intensity of internal energy in a system (gas, liquid, or solid) • Heat • A measure of the total quantity of thermal energy flow into or out of a system

  7. Temperature vs. Heat • Example: • A cup of water at 60°C has much less energy than a hot water heater full of water at 60°C. • BUT, the intensity of heat is the same.

  8. Heat Capacity • Energy required to raise temperature of matter by one degree (at constant pressure or constant volume) • Q = energy in calories • m = mass in grams • ΔT = temperature change in degrees (C or K)

  9. Constant Pressure Heat Capacities Source: Foundations of Engineering, Holtzapple & Reece, 2003

  10. Converting Work into Heat:Joule’s Experiment Source: Foundations of Engineering, Holtzapple & Reece, 2003

  11. Heat CapacityExample Problem • In Joule’s experiment, • Beaker contains 5 kg of water • Mass spinning the stirrer is 90 kg (g=9.81m/s2) • The water increases in temperature by 0.1°C • How far did the mass travel?

  12. Heat Capacity

  13. States of Matter Source: Foundations of Engineering, Holtzapple & Reece, 2003

  14. Phase Diagram Source: Foundations of Engineering, Holtzapple & Reece, 2003

  15. Phase Change • Constant temperature process of transition between phases • Melting / Solidification • Boiling (vaporization) / Condensation

  16. Phase (or State) Change Energy • Where • m = mass (kg) • ΔHvap = latent heat of vaporization (kJ/kg) • ΔHfus = latent heat of fusion (kJ/kg)

  17. Phase-Change Energy Source: Foundations of Engineering, Holtzapple & Reece, 2003

  18. Combustion • Similar to phase change • Where • Qcomb= energy released (MJ) • m = mass (kg) • ΔHcomb = specific heat of combustion (MJ/kg) • Table 22.4

  19. Example 1:Phase-Change Energy • When water changes from solid to liquid, it must absorb 333.56 kJ/kg from the surroundings • What is the energy absorbed to melt ice in units of cal/g?

  20. Example 2a:1st Law of Thermodynamics • If you have 100 g of water at 22°C and add 20 g of ice at 0°C, what will be the temperature of the 100 g of water once all the ice has melted to form 20 g of water at 0°C?

  21. Example 2a:1st Law of Thermodynamics

  22. Example 2b:1st Law of Thermodynamics • What will be the final temperature when the system temperature is uniform (i.e., water from melted ice has warmed and surrounding water has further cooled so that all water is at one temperature)?

  23. Example 3:Latent Heat • If the latent heat of vaporization for water is 2,256.7 kJ/kg, what is the latent heat for water in cal/g?

  24. Example 4a:1st Law of Thermodynamics • If by sweat and evaporation, you lose 0.031 slugs of water during exercising, how many calories of energy in the form of heat is removed from your body? • Weight (mg) of 1 lb is associated with mass of 0.031 slugs • Note on p.688 the conversion from slugs to grams

  25. Example 4b:1st Law of Thermodynamics • If your body mass is 68,100 g (i.e., 150 lbs), how much would your body temperature rise if you did not sweat and evaporate the sweat in order to cool yourself? • Assume the heat capacity for your body is that for water, because your body is ~ 75% water.

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