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Chapter 3

Chapter 3. Matter and Energy. CHAPTER OUTLINE. ENERGY & HEAT. Energy is defined as the capacity of matter to do work . Work is defined as the result of a force acting on a distance. There are two types of energy:. Potential (stored). Kinetic (moving). ENERGY & HEAT.

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Chapter 3

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  1. Chapter 3 Matter and Energy

  2. CHAPTER OUTLINE

  3. ENERGY & HEAT • Energy is defined as the capacity of matter to do work. • Work is defined as the result of a force acting on a distance. • There are two types of energy: Potential (stored) Kinetic (moving)

  4. ENERGY & HEAT • Energy possesses many forms (chemical, electrical, thermal, etc.), and can be converted from one form into another. • In chemistry, energy is commonly expressed as heat. PE is converted to KE

  5. ENERGY & HEAT • Energy is conserved. • The law of conservation of energy states that energy is neither created nor destroyed. • The total amount of energy is constant. • Energy can be changed from one form to another. • Energy can be transferred from one object to another.

  6. UNITS OF ENERGY • The SI unit of energy is the joule (J), named after the English scientist James Joule (1818–1889).

  7. HEAT vs. TEMPERATURE • Heat is measured in SI units of joule or the common unit of calorie. Heat & temperature are NOT the same thing! 1 cal = 4.184 J Although the same amount of heat is added to both containers, the temperature increases more in the container with the smaller amount of water.

  8. HEAT vs. TEMPERATURE The difference between Heat and Temperature A form of energy associated with particles of matter Heat is the total energy of all particles of matter A measure of the intensity of heat or how hot or cold a substance is Temperature is the average kinetic energy of particles of matter

  9. TEMPERATURE • Temperature is a measure of how hot or cold a substance is. • Thermometer is an instrument used for measuring temperature, and is based on thermometric properties of matter (i.e. expansion of solids or liquids). • Three scales are used for measuring temperature.

  10. TEMPERATURE SCALES Fahrenheit 32 - 212 Celsius 0 - 100 Kelvin 273 - 373

  11. TEMPERATURESCALES • To convert from one scale to another the following relationships can be used: K = C + 273 F = (1.8 x C) + 32 C = (F - 32) ÷ 1.8 • Alternately, F = [(C + 40) x 1.8]-40 C = [(F + 40) ÷ 1.8]-40

  12. Example 1: The melting point of silver is 960.8 C. What is this temperature in Kelvin? TK = TC + 273 TK = 960.8 + 273 = 1233.8 K = 1234 K

  13. Example 2: Pure iron melts at 1800 K. What is this temperature in Celsius? TC = TK - 273 TC = 1800 - 273 = 1527 C

  14. Example 3: On a winter day, the temperature is 5 F. What is this temperature on the Celsius scale? TC = [(5 +40) ÷ 1.8]- 40 = -15 C

  15. Example 4: To make ice cream, rock salt is added to crushed ice to reach temperature of -11 C. What is this temperature in Fahrenheit? TF= [(-11 + 40) x 1.8]- 40 = 12 F

  16. SPECIFICHEAT • Different materials have different capacities for storing heat. • The specific heat of a substance is the amount of heat required to change the temperature of 1 g of that substance by 1C. • Units of specific heat are: s = J / g ºC s = cal / g ºC

  17. SPECIFICHEAT Specific Heat of Some Substances Most substances have substantially lower specific heats compared to water

  18. SPECIFICHEAT • When heated, substances with low specific heat get hot faster, while substances with high specific heat get hot at a slower rate. • When cooled, substances with low specific heat cool faster, while substances with high specific heat cool at a slower rate.

  19. CALCULATINGHEAT • The amount of heat lost or gained by a substance is related to three quantities: Mass of substance Specific heat of substance Change in its temperature Heat = x x Q = m x s x T

  20. Example 1: How much heat is needed to raise the temperature of 200. g of water by 10.0 C. (Specific heat of water is 4.184 J/gC) Q = m x s x T m = 200. g s = 4.184 J/gC T = 10.0 C Q = ??? Q = (200. g)(4.184 J/gºC)(10.0 ºC) Q = 8370 J or 8.37 kJ

  21. Example 2: Ethanol has a specific heat of 2.46 J/gC. When 655 J are added to a sample of ethanol, its temperature rises from 18.2 C to 32.8 C. What is the mass in grams of the ethanol sample? Q = 655 J s = 2.46 J/gC T = 14.6 C m = ??? m = 18.2 g

  22. ENERGY & NUTRITION • In the laboratory, foods are burned in a calorimeter to determine their energy. A sample of food is burned in the calorimeter, and the energy released is absorbed by water surrounding the calorimeter. • The energy of the food can be calculated from the mass of the food and the temperature increase of the water.

  23. Example 3: A 2.3-g sample of butter is placed in a calorimeter containing 1900 g of water at a temperature of 17 C. After the complete combustion of the butter, the water has a temperature of 28 C. What is the energy value of butter in Cal/g? 1. Calculate heat absorbed by water Heat absorbed by water Heat released by butter = 2. Calculate energy value of butter

  24. Example 3: 1. Calculate heat absorbed by water m = 1900 g s = 1.00 cal/gC T = 11 C Q = ??? Q = m x s x T Q = (1900 g)(1.00 cal/gºC)(11 ºC) Q = 21000 cal = 21 Cal 2. Calculate energy value of butter 21 Cal 2.3 g = 9.1 Cal/g

  25. ENERGY INCHEMICAL CHANGES • In all chemical changes, matter either absorbs or releases energy. • Higher energy systems are less stable than lower energy systems.

  26. ENERGY INCHEMICAL CHANGES • When energy is released during a chemical change, it is said to be exothermic. • When energy is gained during a chemical change, it is said to be endothermic. Exothermic reactions heat up Endothermic reactions cool down

  27. energy is absorbed higher potential energy energy is given off lower potential energy EXOTHERMIC vs.ENDOTHERMIC Which is exothermic and which is endothermic? 4.4

  28. CLASSIFICATIONOF MATTER • Matter is anything that has mass, and occupies space. • Matter can be classified by its physical state as solid, liquid or gas.

  29. SOLIDS • Solids are densely packed particles with definite shape and volume. • Solid particles have strong forces of attraction towards each other. • Solids are not very compressible. • Ice, diamond, quartz, and iron are examples of solid matter.

  30. LIQUIDS • Liquids are loosely packed particles with definite volume but indefinite shape. • Liquid particles have moderate forces of attraction towards each other and are mobile. • Liquids are slightly compressible. • Water, gasoline, alcohol, and mercury are all examples of liquid matter.

  31. GASES • Gases are very loosely packed particles with indefinite shape or volume. • Gas particles have little or no forces of attraction towards each other. • Gases are very compressible. • Oxygen, helium, and carbon dioxide are examples of gases.

  32. Since the atoms or molecules that compose gases are not in contact with one another, gases can be compressed. GASES ARE COMPRESSIBLE

  33. SUMMARY OFPROPERTIES OF MATTER

  34. CLASSIFICATIONOF MATTER MATTER Anything that has mass MIXTURE Variable composition & properties PURE SUBSTANCE Fixed composition & properties Mixtures can be converted into pure substances by simple physical processes (e.g. filtration, evaporation)

  35. MIXTURES MIXTURE Variable composition & properties HOMOGENEOUS Uniform composition & properties HETEROGENEOUS Non-uniform composition & properties Also called solutions Tea, Coke Ink Salad dressing Cement

  36. PURE SUBSTANCES PURE SUBSTANCE Fixed composition & properties COMPOUNDS 2 or more elements chemically combined ELEMENTS Composed of one type of atom Properties are unique compared to their components Smallest particle is a molecule hydrogen, copper, gold Compounds can be converted into elements by chemical processes or reactions (e.g. electrolysis) water, salt aspirin

  37. PURE SUBSTANCES separation of compound through chemical methods (electrolysis)

  38. CONCEPTCHECK Classify each substance below as element, compound or mixture. Mixture: made of two or more types of substances Element: only one type of atom Compound: composition is fixed Element: only one type of atom

  39. MIXTURES • Mixtures are 2 or more substances physically combined together. • Mixtures possess properties similar to those of their components. • Mixtures can be separated easily by a physical process. • Two types of mixtures are possible: heterogeneous homogeneous

  40. HETEROGENEOUSMIXTURES • Heterogeneous mixtures are non-uniform in their composition. Heterogeneous • Examples include vegetable soup, cement and salad dressing.

  41. HOMOGENEOUSMIXTURES • Homogeneous mixtures are uniform in their composition. • Homogeneous mixtures are called solutions. Homogeneous • Examples include gasoline, soda pop and salt solution.

  42. MIXTURES vs.COMPOUNDS List 3 differences between compounds & mixtures. Composition Compounds have fixed composition while mixtures have varied composition Properties Compounds have unique properties while mixtures have blended properties

  43. MIXTURES vs.COMPOUNDS List 3 differences between compounds & mixtures. Make-up Compounds are chemically combined (cannot be easily separated) while mixtures are physically combined (easily separated)

  44. PHYSICAL & CHEMICALPROPERTIES • The characteristics of a substance are called its properties. • Physical properties are those that describe the matter without changing its composition. Examples are density, color, melting and boiling points, and electrical conductivity.

  45. PHYSICAL & CHEMICALPROPERTIES • The characteristics of a substance are called its properties. • Chemical properties are those that describe how matter behave in combination with other matter, and involve change in its composition. Examples are flammability, corrosion, and reactivity with acids.

  46. Examples: Identify each of the following properties as physical or chemical: • Oxygen is a gas • Helium is un-reactive • Water has high specific heat • Gasoline is flammable • Sodium is soft & shiny Physical Chemical Physical Chemical Physical

  47. PHYSICALCHANGES • Changes in physical properties of matter that do not involve change in its composition are called physical changes. Examples are melting, evaporation and other phase changes. • Physical changes are easily reversible.

  48. CHEMICALCHANGES • A change that alters the chemical composition of matter, and forms new substance is called a chemical change. • Chemical changes are not easily reversible, and are commonly called chemical reactions. Examples are burning, rusting, and reaction with acids.

  49. Examples: Identify each of the following changes as physical or chemical: • Cooking food • Mixing sugar in tea • Carving wood • Burning gas • Food molding Chemical Physical Physical Chemical Chemical

  50. CONSERVATIONOF MASS • Similar to the law of conservation of energy, the law of the conservation of mass states that matter is neither created nor destroyed. • The total mass of substances does not change during a chemical reaction. Mass of Reactants Mass of Products =

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