1. Chapter 2 Matter and Energy

2. 2.1 Classification of Matter Matter is anything that has mass and occupies space. Classification of matters are ◦ Pure substance ◦ Mixture

3. Pure substance Has fixed or definite composition ◦ An element, the simplest type of pure substance and is composed of atoms  E.g. H, Na, O, C etc… ◦ A compound consists of atoms of two or more elements  Chemically combined in proportion and held together by a bond  E.g. H 2 O, NaCl, CO 2 etc…

4. Mixtures Two or more substances are physically mixed, not chemically combined. ◦ Homogeneous mixture or solution  Has a uniform composition  E.g.air contains oxygen and nitrogen ◦ Heterogeneous mixture  Does not have a uniform composition  E.g.raisins in cookie

6. 2.2 States and Properties of Matters Solids have a definite shape. a definite volume. particles that are close together in a fixed arrangement. particles that move very slowly. 6

7. 2.2 States and Properties of Matters Liquids have an indefinite shape, but a definite volume. the same shape as their container. particles that are close together, but mobile. particles that move slowly. 7

8. 2.2 States and Properties of Matters Gases have an indefinite shape. an indefinite volume. the same shape and volume as their container. particles that are far apart. particles that move very fast. 8

9. 2.2 States and Properties of Matters Physical Properties: Characteristics that do not involve a change in a sample’s chemical makeup. Chemical Properties: Characteristics that do involve a change in a sample’s chemical makeup.

10. 2.2 States and Properties of Matters Physical change occurs when matter changes its appearance but the composition stay the same Chemical change takes place when the original substance is converted into one or more new substances ◦ Have different physical and chemical properties

11. Example What type of change, physical or chemical, takes place in the each of the following ◦ Water vapor condenses to form rain ◦ Cesium metal reacts explosively with water ◦ Gold melts at 1064 oC ◦ Food is digested

12. Energy Energy makes objects move. makes things stop. is needed to “do work.” 12

13. Work Work is done when you climb. you lift a bag of groceries. you ride a bicycle. you breathe. your heart pumps blood. water goes over a dam. 13

14. Potential Energy Potential energy is stored energy. Examples are water behind a dam. a compressed spring. chemical bonds in gasoline, coal, or food. 14

15. Kinetic Energy Kinetic energy is the energy of motion. Examples are swimming. water flowing over a dam. working out. burning gasoline. 15

16. Units for Measuring Energy or Heat Heat is measured in joules or calories. 4.184 Joules (J) = 1 calorie (cal) 1 kJ = 1000 J 1 kilocalorie (kcal) = 1000 calories (cal) 16

17. Examples Identify the energy as potential or kinetic. A. Rollerblading B. a peanut butter and jelly sandwich C. mowing the lawn D. gasoline in the gas tank 17

18. 2.3Temperature Conversion Temperature is a measure of how hot or cold an object is compared to another object. indicates that heat flows from the object with a higher temperature to the object with a lower temperature. is measured using a thermometer. 18

19. Temperature conversion T F = 1.8 T C + 32  T C is obtained by rearranging the equation for T F. T F - 32 ° = T C 1.8 T K = T C + 273 contains the lowest possible temperature, absolute zero (0 K). 0 K = –273 °C 19

20. Solving A Temperature Problem A person with hypothermia has a body temperature of 34.8 °C. What is that temperature in °F? T F = 1.8 T C + 32  T F = 1.8 (34.8 °C) + 32 ° exact 3 SFs exact = 62.6 + 32 ° (addition) = 94.6 °F tenth’s 20

21. Example A person with hypothermia has a body temperature of 34.8 °C. What is that temperature in °F?

22. 2.5 Specific Heat Specific heat is different for different substances. is the amount of heat that raises the temperature of 1 g of a substance by 1 °C. in the SI system has units of J/g °C. in the metric system has units of cal/g °C. 22

23. Examples of Specific Heats 23 TABLE 2.7

24. Examples A. When ocean water cools, the surrounding air 1) cools. 2) warms.3) stays the same. B. Sand in the desert is hot in the day, and cool at night. Sand must have a 1) high specific heat. 2) low specific heat. 24

25. Heat = mass x specific heat x ΔT Heat Equation The amount of heat lost or gained by a substance is calculated from the mass of substance (g). temperature change (ΔT). specific heat of the substance (J/g °C). This is expressed as the heat equation. 25 Heat gained = - Heat lost

26. Example How many kJ are needed to raise the temperature of 325 g of water from 15.0 °C to 77.0 °C? 1) 20.4 kJ 2) 77.7 kJ 3) 84.3 kJ 26

27. 2.6 Energy and Nutrition Energy Values in Nutrition ◦ 1 Cal = 1 kcal = 1000 cal ◦ 1 Cal = 4.184 kJ = 4184 J The number of Calories in a food is determined by using an apparatus called a calorimeter

28. Energy Value for Foods The energy (caloric) value of food are the kilocalories or kilojoules obtained from burning 1g of carbohydrate, fat or protein. Table 2.8 Typical Energy (caloric) values of the three food types

29. Examples At a fast-food restaurant, a hamburger contains 37g of carbohydrate, 19g of fat, and 24g of protein. What is the total energy content in kilocalories? Round off the answer to the tenth place

30. Change in state Water evaporates when molecules on the surface gain sufficient energy to form a gas. condenses when gas molecules lose energy and form a liquid.

31. Calculations Using Heat of Fusion The heat of fusion is the amount of heat released when 1 gram of liquid freezes (at its freezing point). is the amount of heat needed to melt 1 gram of a solid (at its melting point). for water (at 0 °C) is 80. cal 1 g water 31

32. Calculation Using Heat of Fusion The heat needed to freeze (or melt) a specific mass of water (or ice) is calculated using the heat of fusion. Heat = g water (ice) x 80. cal 1 g water (ice) Example: How much heat in cal is needed to melt 15. g of ice? 32

33. Sublimation Sublimation occurs when particles change directly from solid to a gas. is typical of dry ice, which sublimes at -78  C. takes place in frost-free refrigerators. is used to prepare freeze- dried foods for long-term storage. 33 Copyright © 2009 by Pearson Education, Inc.

34. Heat of Vaporization The heat of vaporization is the amount of heat absorbed to vaporize 1 g of a liquid to gas at the boiling point. released when 1 g of a gas condenses to liquid at the boiling point. Boiling Point of Water = 100 °C Heat of Vaporization or condense (water) = 540 cal 1 g water 34

35. Summary of Changes of State 35

36. Heating Curve A heating curve illustrates the changes of state as a solid is heated. uses sloped lines to show an increase in temperature. uses plateaus (flat lines) to indicate a change of state. 36

37. Cooling Curve A cooling curve illustrates the changes of state as a gas is cooled. uses sloped lines to indicate a decrease in temperature. uses plateaus (flat lines) to indicate a change of state. 37

38. Combined Heat Calculations To reduce a fever, an infant is packed in 250. g of ice. If the ice (at 0 °C) melts and warms to body temperature (37.0 °C), how many calories are removed from the body? 38