1 / 64

Energy

Energy. Energy. The capacity to do work Many types of energy and one type can be transformed into other types. Example: Law of Conservation of Energy: Energy cannot be created nor destroyed, only transformed. Energy. Potential Energy: Energy stored in an object due to its position

regina
Télécharger la présentation

Energy

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Energy

  2. Energy • The capacity to do work • Many types of energy and one type can be transformed into other types. • Example: • Law of Conservation of Energy: • Energy cannot be created nor destroyed, only transformed

  3. Energy • Potential Energy: • Energy stored in an object due to its position • Kinetic Energy: • Energy of motion

  4. Potential Energy • Chemical Potential Energy • Gravitational Potential Energy • Elastic Potential Energy

  5. Kinetic Energy • Thermal Kinetic Energy • Sound Kinetic Energy • Mechanical Kinetic Energy • Electrical Kinetic Energy • Radiant Kinetic Energy

  6. Heat & Temperature • Heat is not the same thing as temperature!! • The addition of heat to a system often causes an increase in temperature – and vice versa – but the two are not synonymous

  7. Heat & Temperature • Temperature • A measure of the average kinetic energy of the particles in a sample • Heat • A measure of the total kinetic energy of the particles in a sample

  8. Heat & Temperature • Both cups of water are at the same temperature… • but which has a bigger heat content?

  9. Heat & Temperature • Compare a match and a warm bath… • Which has higher temperature? • Which has a bigger heat content?

  10. Heat Transfer – BIG IDEAS • Energy (heat) transfers occur during both physical and chemical changes

  11. Heat Transfer – BIG IDEAS • Heat will always be spontaneously transferred from matter at a higher temperature to matter at a lower temperature • When you place your hand on a cold glass, it feels cold – but what is really happening in terms of heat transfer? • Heat is being transferred from your hand to the glass.

  12. Heat Transfer • Endothermic • If a system absorbs heat, the system undergoes an endothermic change • Heat is being absorbed from the surroundings, which makes it feel cold!

  13. Heat Transfer • Exothermic • If a system releases/loses heat, the system undergoes an exothermic change • Heat is being released into the surroundings, which makes it feel warm!

  14. Enthalpy • Thermochemistry is the study of heat exchange during physical and chemical changes • Enthalpy (H) is a measure of the heat contained by a system • Remember, though, that we cannot measure heat directly, so instead we measure changes in enthalpy (ΔH) • An enthalpy change is the amount of heat absorbed/lost by a system as heat

  15. Enthalpy Changes • If a system is gaining energy (+ΔH) in an endothermic change, why does the system feel cold to the touch? • Because the system is absorbing heat from the environment (i.e., your hand) • Similarly, when a system undergoes an exothermic change, you feel the heat being released by the system as warmth

  16. Enthalpy of Reaction

  17. Enthalpy of Reaction • The quantity of energy transferred during a chemical reaction is referred to as the enthalpy of reaction (or “heat of reaction”) • The reaction system includes all reactants and products • The enthalpy of reaction is calculated by:

  18. Enthalpy of Reaction • If the amount of energy stored in the bonds of the reactants is greater than the amount of energy stored in the bonds of the products, then the system has lost energy • - ΔH • Exothermic

  19. Exothermic Reaction Pathway DH = negative

  20. Enthalpy of Reaction • If the amount of energy stored in the bonds of the reactants is less than the amount of energy stored in the bonds of the products, then the system has gained/absorbed energy • + ΔH • Endothermic

  21. Endothermic Reaction Pathway DH = positive

  22. Calculating Enthalpy of Reaction • The enthalpy of reaction can be calculated using the individuals heats of formation of each product/reactant in a chemical reaction • Example: • Calculate the standard reaction enthalpy, DHoRXN for the following equation: • Is this reaction exothermic or endothermic? • = -890.2 kJ • Exothermic

  23. Thermochemical Equations • A thermochemical equation includes the ΔH of the reaction EXOTHERMIC ENDOTHERMIC

  24. Thermochemical Equations EXOTHERMIC ENDOTHERMIC

  25. ThermochemicalStoichiometry • Sample Problem: How much heat will be released if 1.00 g of hydrogen peroxide decomposes in a bombardier beetle to produce a steam spray?

  26. Chemistry Question • Calculate the standard enthalpy of reaction for the combustion of ethyne (C2H6) • Answer = -3121.2 kJ • Compare this answer to the heat of combustion of ethane in Table A-5. What do you notice? • Why are these values different?

  27. Terms • Enthalpy of reaction • Quantity of energy transferred during a chemical reaction • Molar enthalpy of formation • Enthalpy change that occurs when one mole of a compound is formed • Defined as one mole of product • Molar enthalpy of combustion • Enthalpy change that occurs during the complete combustion of one mole of a substance • Defined as one mole of reactant

  28. Standard States • The degree sign is used to indicate a measurement made at the standard state of a substance • Typically 1 atm and room temperature (~298 K) • ΔH°f for water is the standard enthalpy of formation for liquid water (since water would be liquid under standard conditions)

  29. Bonus Assignment for Today • A) ΔH°RXN for combustion of ethyne (C2H2) • B) The absolute value of the amount of energy released when 1.000 pound of ethyne is completely combusted (1 lb = 453.6 g) • C) The mass of ethyne combusted if ΔH = -1200. kJ • D) The absolute value of the amount of energy released when 14 g of ethyne are mixed with 25 g of O2

  30. Heat Transfer& Specific Heat

  31. Heat Transfer • Heat is always transferred from a “warm” object to a “cold” object – until an equilibrium is reached • Say you place some ice at into a beaker of hot water at. What factors will affect the final equilibrium temperature of the system?

  32. Energy • calorie • the amount of energy required to raise the temperature of one gram of water by one degree Celsius. • Joule • SI unit of energy • 1 cal = 4.184 J

  33. Energy • Express 60.1 cal of energy in units of joules.

  34. Calculating Energy Requirements • How does the amount of substance heated affect the energy required? • Determine the amount of energy in joules required to raise the temperature of 7.40 g of water from 29 °C to 46 °C.

  35. Specific Heat • The amount of heat required to raise the temperature of 1 g of any substance by 1oC • For liquid water, s = 4.184

  36. Specific Heat • Intensive Property! • Specific heat describes how well an object retains heat • A substance with a low specific heat is quickly heated, but also quickly cools • A substance with a high specific heat takes a long time to warm up, but will also retain that heat for a longer period

  37. Specific Heat • Styrofoam is a very poor conductor of heat; it is a good insulator. It has a high specific heat. • Metals are good conductors of heat. They have low specific heats.

  38. Heat Transfer • Say you place some ice at into a beaker of hot water at. What factors will affect the final equilibrium temperature of the system? • Mass of ice and water • Initial temperature of the water/ice • The specific heat of the ice/water

  39. Specific Heat • What quantity of energy (in joules) is required to heat a piece of iron weighing 1.3 g from 25 °C to 46 ° C? • The specific heat of iron is 0.45

  40. Specific Heat • Q = s x m xΔT • Q = energy required • s = specific heat capacity • m = mass of sample in grams • ΔT = change in temperature in degrees Celsius

  41. Specific Heat • A 1.6 g sample of a metal that has the appearance of gold requires 5.8 J of energy to change its temperature from 23 °C to 41 °C. Is the metal pure gold? • Specific heat capacity of gold = 0.13

  42. Heat Transfer • The Law of Conservation of Energy says that energy can never be created or destroyed • So, if one object is absorbing heat, another object must be losing it, and vice versa • We can use this idea to solve problems involving two objects.

  43. Example • A 5.0 g sample of an unknown metal is placed into a beaker of water containing 500. g of water. The metal has an initial temperature of 45.0°C and the water has an initial temperature of 20.0°C. When equilibrium is reached, the final temperature is 33.5°C. What is the specific heat of the metal?

  44. Example • A 20.0 g piece of aluminum (cp = 0.902 J/g·°C) is placed into a beaker containing 50.0 mL of water at 10.0°C. After a few minutes, an equilibrium temperature of 25.0°C is reached. What was the initial temperature of the aluminum metal?

More Related