Energy

1. Energy The ability to do work. • Work - cause a change or move an object. • Many types- all can be changed into the other. W = F x d

2. Types of energy • Potential- stored energy • Position, condition or composition • Kinetic Energy- energy something has because its moving • Heat- the energy that moves because of a temperature difference. • Chemical energy- energy released or absorbed in a chemical change. • Electrical energy - energy of moving charges

3. Types of Energy con’t • Radiant Energy- energy that can travel through empty space (light, UV, infrared, radio) • Nuclear Energy – Energy from changing the nucleus of atoms

4. Conservation of Energy Energy can be neither created or destroyed in ordinary changes (not nuclear), it can only change form. • Discovered by Julius Robert Mayer in 1842 • Now called: The First Law of Thermodynamics Law of Conservation of Mass - Energy The total amount of mass and energy in the universe is constant.

5. The Kinetic Molecular Theory is useful in describing thermal energy, heat, and temperature. • Some theories are based on supporting postulates. • A postulate is a statement which is agreed on by consensus among scientists. • The following are important postulates of the kinetic molecular theory:

6. All matter consists of atoms. • Atoms may join together to form molecules. • Solids usually maintain both their shape and their volume. • Liquids maintain their volume, but not their shape. • Gases do not maintain shape or volume. They will expand to fill a container of any size. • Molecular motion is random. • Molecular motion is greatest in gases, less in liquids, and least in solids. • Collisions between atoms and molecules transfers energy between them. • Molecules in motion possess kinetic energy. • Molecules in gases do not exert large forces on one another, unless they are colliding. Kinetic Molecular Theory

7. Thermal energy is the average of the potential and kinetic energies possessed by atoms and molecules experiencing random motion. • Heat is transferred by convection, conduction, or radiation. (review the definitions of these words) • Heat is the thermal energy transferred from one object to another due to differences in temperature. Heat flow from high to low temperature.

8. There is no direct method used to measure heat. Indirect methods must be used. Temperature is a measure of the average kinetic energy of the molecules of a substance. • There is a direct relationship between temperature and avg. kinetic energy! • Temperature can be measured with a thermometer.

9. One way a thermometer can be calibrated is by the amount of thermal expansionand contraction that occurs within a given type of substance. • Thermometers are limited by the physical properties of the substance from which they are made. (i.e., An alcohol thermometer is of little use above the boiling point of alcohol, and a mercury thermometer will not be of any use below the freezing point of mercury.)

10. What's the difference between the Fahrenheit and Celsius temperature scales? • Both scales are based on the freezing conditions of water, a very common and available liquid. • Since water freezes and boils at temperatures that are rather easy to generate (even before modern refrigeration), it is the most likely substance on which to base a temperature scale.

11. 100ºC = 212ºF 0ºC = 32ºF 100ºC 212ºF 32ºF 0ºC

12. How much it changes 100ºC = 212ºF 0ºC = 32ºF 100ºC = 180ºF 0ºC 100ºC 212ºF 32ºF

13. How much it changes 100ºC = 212ºF 0ºC = 32ºF 100ºC = 180ºF 1ºC = (180/100)ºF 1ºC = 9/5ºF 0ºC 100ºC 212ºF 32ºF

14. Scientists use a third scale, called the "absolute" or Kelvin scale. • This scale was invented by William Thomson, Lord Kelvin, a British scientist who made important discoveries about heat in the 1800's. • Scientists have determined that the coldest it can get (theoretically) is minus 273.15 degrees Celsius. • This temperature has never actually been reached, though scientists have come close. The value, minus 273.15 degrees Celsius, is called "absolute zero". • At this temperature scientists believe that molecular motion would stop. You can't get any colder than that. • The Kelvin scale uses this number as zero. To get other temperatures in the Kelvin scale, you add 273 degrees to the Celsius temperature.

15. The important idea is that temperature is really a measure of something, the average motion (kinetic energy, KE) of the molecules. KE = ½ mv2 • Does 0°C really mean 0 KE? nope... it simply means the freezing point of water, a convenient standard. • We have to cool things down to –273.15°C before we reach 0 KE. This is called 0 Kelvin (0 K, note: NO ° symbol.) • For phenomena that are proportional to the KE of the particles (pressure of a gas, etc.) you must use temperatures in K.

16. Temperature Conversion • K = °C + 273 • °C = K – 273 • °F = 9/5 °C + 32 • °C = 5/9 (°F – 32) Note: In Kelvin notation, the degree sign is omitted: 283K

17. Celsius to Fahrenheit: • *A mental shortcut for a rough estimate: • Double the temperature given in Celsius • Add 30 to the result to find the approximate temperature in Fahrenheit.

18. Fahrenheit to Celsius: • *A mental shortcut for a rough estimate: • Subtract 30 from the temperature given in Fahrenheit • Take half of the result to find the approximate temperature in Celsius.

19. How Do We Measure Energy? Energy is measured in many ways. BTU • One of the basic measuring blocks is called a BTU. This stands for British thermal unit and was invented by the English. • Btu is the amount of heat energy it takes to raise the temperature of one pound of water by one degree Fahrenheit, at sea level. • One Btu equals about one blue-tip kitchen match. • One thousand BTUs roughly equals: One average candy bar or 4/5 of a peanut butter and jelly sandwich. • It takes about 2,000 BTUs to make a pot of coffee.

20. Calorie • A calorie is a unit of measurement for energy. • Calorie is a French word derived from the Latin word: calor (heat). Modern definitions for calorie fall into two classes: • The small calorie or gram calorie approximates the energy needed to increase the temperature of 1gram of water by 1 °C. This is about 4.184 joules. • The large calorie or kilogram calorie approximates the energy needed to increase the temperature of 1 kg of water by 1 °C. This is about 4.184 kJ, and exactly 1000 small calories. • 1 cal = 4.184 J

21. Joule • Energy also can be measured in joules. (Joules sounds exactly like the word jewels, as in diamonds and emeralds.) • A thousand joules is equal to a British thermal unit. • 1,000 joules = 1 Btu • So, it would take 2 million joules to make a pot of coffee.

22. The term "joule" is named after an English scientist James Prescott Joule who lived from 1818 to 1889. • He discovered that heat is a type of energy. • One joule is the amount of energy needed to lift something weighing one pound to a height of nine inches. • Around the world, scientists measure energy in j.

23. Like in the metric system, you can have kilojoules -- "kilo" means 1,000. • 1,000 joules = 1 kilojoule= 1 Btu • 1 cal = 4.184 J

24. Melting Vaporization Freezing Condensation Phase Changes Solid Gas Liquid

25. Sublimation Vaporization Deposition endothermic Melting Solid Gas Liquid Freezing Condensation exothermic

26. Fusion (Melting) • The temperature at which a liquid and a solid are in equilibrium • The melting point for ice is 0ºC • The melting point of a substance is the same as its freezing point • Fusion (melting) is an endothermic process

27. Vaporization (Endothermic) • Vaporization is the changing of a liquid to a gas below its boiling point. • Vapor pressure is the collision of many gas particles with the walls of a sealed container. • A substance in its gaseous state has more kinetic energy than in its liquid state so it must absorb energy (take energy away) from its environment (reducing its kinetic energy or temperature) to change from a liquid to a gas.

28. Boiling Point (Endothermic) • The boiling point of a liquid is the temperature at which it boils at a certain pressure. • At the boiling point the vapor pressure exerted upward by the vapor particles equals the downward pressure of the atmosphere ("pressure up = pressure down"). • The normal boiling point of a liquid is the temperate at which it boils under one atmosphere of pressure. • The “normal” boiling point of water is 100 degrees C.

29. Boiling Point and Pressure Changes • Since the boiling point of a liquid is directly related to external pressure, the boiling point of a liquid will change with external pressure. • If the air pressure is low, such as at a high altitude, it would require less vapor pressure and the boiling temperature would be less than the normal boiling point. For this reason food needs to be boiled for longer at higher altitude to become fully cooked. • Conversely, if food is put in a pressure cooker, the high pressure allows water to get much hotter than 100 ºC and food will cook much quicker.

30. Evaporation (Endothermic) • Even at room temperature liquids can be converted into gases. • A small percentage of molecules are moving with relatively high kinetic energy. • If these fast moving molecules possess enough KE to overcome attractive forces within a liquid, they can escape through the surface into the gaseous state. • The loss of the higher energy molecules due to evaporation leads to a lowering of the average kinetic energy of the remaining molecules, this results in a decrease in the temperature of the liquid. • You feel cool after being in hot water because the evaporation of water from the body has drawn heat away. Also when we sweat! ( Cooling process)

31. Condensation (Exothermic) • The phase change when a gas is converted into a liquid. • When a gas condenses into its liquid state the particles lose energy which is absorbed by the environment causing it to gain energy. • This is, in part, the cause of storms such as hurricanes. when warm, moist air rises over the southern Atlantic ocean the water vapor it carries condenses and releases heat energy. • This heat energy draws up more warm, moist air and its water vapor condenses creating a positive feedback system which creates the hurricane. • Hurricanes generally die soon after moving over land because the moisture supply which generated the heat by condensation is cut off.

32. Freezing (Exothermic) • The value of a substances freezing point and melting points are the same (the change of state is the same, only in different direction. ) • Water normally freezes at 0 degrees C.

33. Sublimation • Sublimation is the process by which a solid substance changes directly into a gaseous substance without going through the liquid phase. • The most common example of sublimation is dry ice. • Dry ice is solid (frozen) carbon dioxide which sublimes into gaseous carbon dioxide. • The opposite of sublimation is deposition.

34. 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

35. 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 gas liquid Slope = Specific Heat Solid

36. 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 Solid and liquid

37. 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 liquid and gas

38. 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

39. 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 Fusion

40. 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 Plateau = phase equilibrium

41. Energy and Phase Changes

42. Heat is transferred to different materials at different rates. • The specific heat capacity (C) determines the rate at which heat will be absorbed. • Even though mass is present in the formula it is an intensive property like density and is unique for each substance. • The specific heat capacity for water is 4.18J/g • The quantity of heat absorbed (Q) can be calculated by: Q=mCT m=mass T=change in temperature

43. Heat Capacity Heat capacity is an extensive property, meaning it depends on the mass of the object. Ex: 1000g of water can hold more heat than 10 g of water.

44. Calculating Energy • Q means heat energy lost or gained. • Law of Conservation of Mass-Energy • m= mass of substance; Cp= specific heat capacity; DT = change in temperature Qlost = Qgained Three equations: • Q= mass x Cp x DT • Q= Hf x mass • Q= Hv x mass

45. 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

46. Energy and Phase Change • Heat of fusion energy required to change one gram of a substance from solid to liquid. (endothermic rxn) • Heat of solidification energy released when one gram of a substance changes from liquid to solid. (exothermic rxn) • For water 80 cal/g or 334 J/g

47. Energy and Phase Change • Heat of vaporization energy required to change one gram of a substance from liquid to gas. (endothermic rxn) • Heat of condensation energy released when one gram of a substance changes from gas to liquid. (exothermic rxn) • For water 540 cal/g or 2260 J/g

48. Three equations: Q= mass x Cp x DT (used at slopes) Q= Hf x mass (used at s/l equilibria) Q= Hv x mass (used at l/g equilibria)

49. 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

50. Regents Question As ice melts at standard pressure, its temperature remains at 0°C until it has completely melted. Its potential energy (1) decreases (2) increases (3) remains the same þ