Chapter 3

1. Chapter 3 Matter and Energy

2. Assigned Problems Recommended • Exercises: 1-27 (odd) Required • Problems: 29-85 (odd) • Cumulative Problems: 87-105 (odd) Optional • Highlight Problems: 107-113 (odd)

3. What Is Matter? • Matter is any material that has mass and occupies space • Matter is made up of small particles • Atoms • Molecules • Includes all things (living and nonliving) such as plants, soil, and rocks and any material we use such as water, wood, clothing, etc. • Classifications (of a sample of matter) is based on whether its shape and volume are definite or indefinite

4. Classifying Matter According to Its State • Solid • Has a rigid, definite shape and definite volume • Crystalline solids have a regular, internal long-range order of atoms, ions, or molecules • Amorphous solids have no long-range order of atoms, ions, or molecules in their lattice structure • Liquid • Has an indefinite shape and a definite volume. • It will take the shape of the container it fills • Gas • Has an indefinite shape and an indefinite volume. • It will take the shape and completely fill the volume of the container it fills • Gases are compressible

5. Water is one of the few substances commonly found in all three physical states

6. Classifying Matter by Its Composition • Matter can also be classified in terms of its chemical composition • Pure Substances: Composed of only one atom or molecule • Mixtures: Composed of two or more different atoms or molecules combined in various proportions

7. Pure Substances • Matter that has a definite and constant composition is a pure substance • Composed of the same substance; no variation • 6 million pure substances have been isolated: 112 are elements, the rest are compounds • The two classifications of pure substances: • Elements: e.g., a pure sample of copper or a pure sample of gold (one type of atom) • Compounds: e.g., example, a pure sample of water or a pure sample of sucrose (one type of molecule)

8. Pure Substances • Elements • Substances which can not be broken down into simpler substances by chemical reactions • Fundamental substances • Compounds • Two or more elements combined chemically in a definite and constant ratio • Can be broken down into simpler substances • Most of matter is in the compound form

9. Pure Substances • Compounds • Results from a chemical combination of two or more elements • Can be broken down into elements by chemical processes • Properties of the compound not related to the properties of the elements that compose it • Water is composed of hydrogen and oxygen gases (combined in a 2:1 ratio)

10. Mixtures • Something of variable composition • Result from the physical combination of two or more substances (elements or compounds) • Made up of two or more types of substances physically mixed • Each substance retains its identity because the substances are not chemically mixed • Mixtures of the same components can vary in composition

11. Mixtures • Mixtures can be classified by the (visual) uniformity of the mixture’s components • Homogeneous mixture: • Same uniform composition throughout • Not possible to see the two substances present • Heterogeneous mixture: • Composition is not uniform throughout the sample. • It contains visibly different parts or phases

12. Mixtures • Homogenous mixtures • A sugar solution • 14 karat gold, a mixture of copper and gold • Air, a mixture of gases (oxygen, nitrogen) • Heterogeneous mixture • Oil and vinegar • Raisin cookies • Sand • Pure substance • e.g. copper (all elements are pure substances)

13. Compounds vs. Mixtures • Compounds are not mixtures • Cannot be separated by a physical process • Can be subdivided by a chemical process into two or more simpler substances • Simpler substances have different properties from the compound • Mixtures • Unlike compounds, mixtures can be separated by a physical process • Each substance in a mixture retain its own individual properties

14. Classification of Matter Physical Methods Chemical Methods

15. Physical and Chemical Properties • Various kinds of matter are differentiated by their properties • Properties are the characteristics of a substance used to identify and describe it • Two general categories: • Physical Properties • Chemical Properties • Properties can be: • Directly observable (physical) • The interaction of the matter with other substances (chemical)

16. Physical and Chemical Properties: Physical Properties • A physical property is a characteristic of a substance that can be observed without changing a substance into another substance • Characteristics of matter that can be directly observed or measured without changing its identity or composition • Color, odor, physical state, density, melting point, boiling point

17. Physical and Chemical Properties: Chemical Properties • A chemical property describes the way a substance undergoes a change or resists change to form a new substance • Properties that matter exhibits as it undergoes changes in chemical composition: • Objects made from copper will turn green when exposed to moist air for long periods • Gold objects will resist change when exposed to moist air for long periods • Sodium metal will react strongly with water and produce hydrogen gas.

18. Classifying Properties • The boiling point of ethyl alcohol is 78 °C • Physical property – describes an inherent characteristic of alcohol, its boiling point • Diamond is very hard • Physical property – describes inherent characteristic of diamond – hardness • Sugar ferments to form ethyl alcohol • Chemical property – describes behavior of sugar, ability to form a new substance (ethyl alcohol)

19. Physical and Chemical Changes • Changes in matter are regular occurrences: • Food is cooked • Paper is burned • Iron rusts • Matter undergoes changes as a result of the application of energy • Changes in matter are also categorized as two types: • Physical • Chemical

20. Physical and Chemical Changes • A physical change is a process that alters the appearance of a substance but does not change its chemical identity or composition • Folding aluminum foil sheets • Crushing ice cubes • No new substance is formed • Most common is a change of a substance’s physical state • The freezing of liquid water • Evaporation of liquid water to steam

21. Physical and Chemical Changes • A chemical change is a process that changes the chemical composition of a substance • Also called a chemical reaction • (At least) one new substance is produced • Wood burning, iron rusting, alka-seltzer tablet reacting with water • During a chemical change, the original substance is converted into one or more new substances with different chemical and physical properties

22. Classifying Changes • Melting of snow • Physical change – a change of state but not a change in composition • Burning of gasoline • Chemical change – combines with oxygen to form new compounds • Rusting of iron • Chemical change – combines with oxygen to form a new reddish-colored substance (ferric oxide)

23. Classifying Changes • Iron metal is melted • Physical change – describes a state change, but the material is still iron • Iron combines with oxygen to form rust • Chemical change – describes how iron and oxygen combine to make a new substance, rust (ferric oxide) • Sugar ferments to form ethyl alcohol • Chemical change – describes how sugar forms a new substance (ethyl alcohol)

24. Conservation of Mass • During a physical change: No new substance is formed • During a chemical change: At least one new substance is formed • Whether it is a physical or chemical change, the amount of matter remains the same • The law of conservation of mass states that the total mass of materials present after a chemical reaction is the same as the total mass before the reaction • Matter is never created or destroyed

25. Energy • Two major components of the universe: • Matter • Energy • Energy is the capacity to do work or produce heat • Electrical, radiant, mechanical, thermal, chemical, nuclear • Nearly all changes that matter undergoes involves the release or absorption of energy • Chemistry is the study of matter • The properties of different types of matter • The way matter behaves when influenced by other matter and/or energy

26. Energy • Energyis the part of the universe that has the ability to do work • Energy can be converted from one form to another but it is neither created nor destroyed (the law of conservation of energy) • Energy has two classifications • Potential: Stored energy • Kinetic: Motion energy • All physical changes and chemical changes involve energy changes

27. Energy • Potential energy: • Determined by an objects position (or composition) • Chemical energy (also potential energy) is stored in the bonds contained within a molecule. It is released in a chemical reaction • Kinetic energy • Energy that matter acquires due to motion • Converted from the potential energy • All forms of energy can be quantified in the same units

28. Units of Energy • The joule (J) is the SI unit of heat energy • The calorie (cal) is an older unit used for measuring heat energy (not an SI unit) • The amount of energy needed to raise the temperature of one gram of water by 1°C • The Cal is the unit of heat energy in nutrition 4.184 J = 1 cal 1 kcal = 1000 cal 1 Cal = 1000 cal = 1 kcal

29. Energy: Chemical and Physical Change • All physical changes and chemical changes involve energy changes which convert energy from one form to another • In terms of a chemical reaction the universe is divided into two parts: • The system (chemical reaction) • The surroundings (everything else) • The potential energy differences between the reactants and products determine whether heat flows into or out of a chemical system • Whether a reaction is exothermic or endothermic depends on how the potential energy of the products compares to the PE of the reactants

30. Energy: Chemical and Physical Change • Chemical systems with high potential energy tend to change in order to lower their potential energy by the release of heat • Chemical reactions that release heat are called exothermic • Chemical systems with low potential energy tend to change in order to increase their potential energy by the absorption of heat • Chemical reactions that absorb heat are called endothermic

31. Temperature • Temperature is a number related to the average kinetic energy of the molecules of a substance • In a substance, the temperature: • measures the hotness or coldness of an object • measures the average molecular motions in a system • relates (directly) to the kinetic energy of the molecules

32. Temperature • Fahrenheit Scale, °F • Used in USA • Water’s freezing point = 32°F, boiling point = 212°F • Celsius Scale, °C • Used in science (USA) and everyday use in most of the world • Temperature unit larger than the Fahrenheit • Water’s freezing point = 0°C, boiling point = 100°C

33. Temperature • Kelvin Scale, K • SI Unit • Used in science • Temperature unit same size as Celsius • Water’s freezing point = 273 K (0 ºC), boiling point = 373 K (100 ºC) • Absolute zero is the lowest temperature theoretically possible • No negative temperatures

34. Converting °C to °F • Units are different sizes • Fahrenheit scale: 180 degree intervals between freezing and boiling • Celsius scale: 100 degree intervals between freezing and boiling

36. Converting °C to °F • To convert from °C to °F • Different values for the freezing points • Different size of the degree intervals in each scale 32 °F 0 °C add 32 to the °F value

37. Converting °C to K • Temperature units are the same size • Differ only in the value assigned to their reference points • 25°C is room temperature, what is the equivalent temperature on the Kelvin scale? 0 °C = 273 K add 273 to the °C value K = °C + 273 25 ºC + 273 = 298 K 25 ºC + 273 = 298 K 25 ºC + 273 = 298 K

38. Example • A cake is baked at 350 °F. What is this in Centigrade/Celsius? In Kelvin? 318 °F

39. Temperature Changes:Heat Capacity • Heat is the total amount of energy in a system • It is function of the amount of motion (kinetic energy) contained in molecules • It is also a function of the potential energy of the molecules • It involves the exchange of thermal energy caused by a temperature difference

40. Heat vs. Temperature • Within a quantity of matter: • Heat has units of Joules and temperature has units in degrees • Temperature relates only to kinetic energy within a molecule • Heat is the total amount of energy in a molecule: It contains a kinetic and potential energy component • Heat energy can be added or removed without a change in temperature • As heat energy increases the temperature increases

41. Temperature Changes: Heat Capacity • Heat energy isthe form of energy most often released or required for chemical and physical changes • Every substance must absorb a different amount of heat to reach a certain temperature • Different substances respond differently when heat is applied • The amount of heat required to raise the temperature of a given quantity of a substance by 1 ºC is called its heat capacity

42. Temperature Changes: Specific Heat • If 4.184 J of heat is applied to: • 1 g of water, its temperature is raised by 1 °C • 1 g of gold, its temperature is raised by 32 °C • Some substances requires large amounts of heat to change their temperatures, and others require a small amount • The precise amount of heat that is required to cause a given amount of substance (in grams) to have a rise in temperature is called a substance’s “specific heat”

43. Specific Heat • The amount of heat energy (q) needed to raise 1 gram of a substance by 1 °C • Specific to the substance • The higher the specific heat value, the less its temperature will change when it absorbs heat • SH values given in table 3.4, page 71 • Only for heating/cooling not for changes in state

44. Specific Heat Expression with Calories and Joules • 1 cal is the energy needed to heat 1 g of water 1 °C • 1 cal is 4.184 J • Make a conversion factor from the statements

45. Specific Heat Equation • The rearrangement of the SH equation gives the expression called the “heat equation” • q = heat • C = specific heat (different for each substance) • m = mass (g) • ∆T = change in temperature (°C) C

46. Specific Heat Equation • Energy (heat) required to change the temperature of a substance depends on: • The amountof substance being heated (g) • Thetemperature change (initial T and final T in °C) • Theidentity of the substance

47. Energy and T • The amount the temperature of an object increases depends on the amount of heat added (q) • If you double the added heat energy (q), the temperature will increase twice as much. • When a substance absorbs energy, q is positive, temperature increases • When a substance loses energy, q is negative, temperature decreases 2× 2×

48. Energy and Heat Capacity Calculations • Use same problem solving steps as before (Chapter 2) • State the given and needed units • Write the unit plan to convert the given unit to the final unit • State the equalities and the conversion factors • Set up the problem to cancel the units • Pepsi One™ contains 1 Calorie per can. How many joules is this? 1 Cal = 1000 cal 4.184 J = 1 cal

49. Energy and Heat Capacity Calculations • The 4184 J from the Pepsi One™ will heat how many grams of water from 0°C to boiling?

50. Energy and Heat Capacity Calculations • How many grams of water would reach boiling if the water started out at room temperature (25°C)?