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Energy & Chemical Change

Energy & Chemical Change. Chapter 16. 16.1 Energy. What is Energy? What are some types of energy that you are familiar with? Kinetic Potential Thermal Electrical Chemical Nuclear. A. The Nature of Energy. Energy: ability to do work or produce heat

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Energy & Chemical Change

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  1. Energy & Chemical Change Chapter 16

  2. 16.1 Energy • What is Energy? • What are some types of energy that you are familiar with? • Kinetic • Potential • Thermal • Electrical • Chemical • Nuclear

  3. A. The Nature of Energy • Energy: ability to do work or produce heat • Potential Energy: energy due to composition (chemical) or position of an object (gravitational) • Kinetic Energy: energy of motion

  4. Kinetic energy of a substance is directly related to the constant random motion of it particles and is proportional to temperature.

  5. Chemical Potential Energy of a substance depends upon its composition • Type of atoms • # & type of chemical bonds • How atoms are arranged

  6. Law of conservation of energy • States that in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed

  7. Chemical potential energy • Energy stored in a substance because of its composition • Ex. Gasoline – when burned chemical potential energy is converted to useful mechanical energy. • Heat (q): process of energy flowing from a warmer object to a cooler object

  8. How is Heat Transferred? • Conduction – transfer of heat within solid objects by direct contact • Convection – transfer of heat within fluids (liquids & gases) • Radiation – transfer of heat by electromagnetic radiation (like the sun)

  9. Measuring Heat - Units • calorie: amount of heat required to raise temperature of one gram of pure water one degree Celsius • Joule: SI unit of heat & energy • 1 calorie = 4.184 joules • 1000 calories = 1 Kcal or 1 nutritional Calorie • Practice Problems p.492

  10. Specific Heat • Amount of heat required to raise the temperature of one gram of that substance by one degree Celsius • Remember water has a high specific heat – it takes lots of energy to change it’s temperature

  11. Calculating heat evolved & absorbed q = m x c x T q = heat absorbed or released (the value is positive if heat is absorbed and negative if heat is released) m = mass of sample in grams c = specific heat of substance (can be determined or looked up in a table) T = difference between final temperature & initial temperature

  12. Phase Changes – know for your TEST! • Which phase changes require energy? (endothermic) • Melting, evaporation, sublimation • Which phase changes release energy? (exothermic) • Freezing, condensation, deposition

  13. Energy and Phase Change • Heat of vaporization - energy required to change one gram of a substance from liquid to gas. • Heat of condensation - energy released when one gram of a substance changes from gas to liquid. • For water 540 cal/g

  14. Energy and Phase Change • Heat of fusion - energy required to change one gram of a substance from solid to liquid. • Heat of solidification - energy released when one gram of a substance changes from liquid to solid. • For water 80 cal/g

  15. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Temperature in Celsius Pressure in mmHg or torr

  16. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Heat of Vaporization Temperature in Celsius Pressure in mmHg or torr

  17. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Temperature in Celsius Heat of Fusion Pressure in mmHg or torr

  18. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Steam Water Temperature in Celsius Slope = Specific Heat Ice Pressure in mmHg or torr

  19. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Both Water and Steam Temperature in Celsius Pressure in mmHg or torr

  20. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Temperature in Celsius Ice and Water Pressure in mmHg or torr

  21. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Temperature in Celsius Pressure in mmHg or torr

  22. Phase Diagram • Phase diagram – a graph of pressure vs. temperature that shows in which phase a substance exist under different conditions of temperature and pressure. • Triple Point – the point on the phase diagram that represents the temperature and pressure at which all three phases can coexist • Critical Point – the point that indicates the critical temperature and pressure. Critical temperature is the temp. above which the sub. can’t exist in the liquid state. Critical pressure is the lowest pressure at which the substance can exist at the critical temperature.

  23. Phase Diagram of Water

  24. 16.2 Heat in Chemical Reactions & Processes

  25. Measuring Heat • Calorimeter is an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process

  26. Chemical Energy & Universe • System: specific part of universe that contains reaction or process • Surroundings: everything in the universe other than the system • Universe = system + surroundings

  27. 1. Enthalpy & Enthalpy changes • Enthalpy: (H) heat content of a system at constant pressure • Can’t measure actual energy or enthalpy of a substance, you can measure change in enthalpy, which is heat absorbed or released in a chemical reaction

  28. Enthalpy • ∆Hrxn = Hproducts – Hreactants • When ∆Hrxn is negative the reaction is exothermic – Hproducts < Hreactants • When ∆Hrxn is positive the reaction is endothermic - Hproducts > Hreactants

  29. 2. Sign of enthalpy reaction • Exothermic Reactions: heat pack

  30. Endothermic Reactions: cold pack

  31. Practice Problem Thermodynamic Heats of Formation for one mole at 298K and 1 atmosphere pressure • Substance (form) Enthalpy of • Formation ΔHf (kJ) • NaCl(s) -411.15 • Na+(aq) -240.12 • Cl-(aq) -167.16 • Na+(aq) + Cl- (aq) NaCl(s) What is the change in enthalpy (ΔH) for this reaction? • ∆Hrxn = Hproducts – Hreactants • ∆Hrxn = [-411.15] – [(-240.12)+(-167.16)] = -3.87 kJ • So is the reactionExothermic or Endothermic?

  32. Endothermic & Exothermic Graphs

  33. Hess’ Law States that if you can add two or more thermochemical equations to produce a final equation then the sum of the enthalpy changes of the individual reactions is the enthalpy change for the final reaction Example: Calculate the change in enthalpy for the following reaction: 2S(s) + 3O2(g)  2SO3(g) Given: S(s) + O2(g)  SO2(g) ëH = -297 kJ 2SO3(g)  2SO2(g) + O2(g) ëH = 198 kJ

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