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Chapter 3

Chapter 3. Matter and Energy. Homework. Assigned Problems (odd numbers only) “Questions and Problems” 3.1 to 3.41 (begins on page 61) “Additional Questions and Problems” 3.49 to 3.69 (page 87-88) “Challenge Questions” 3.71 and 3.75, (page 88). Matter.

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Chapter 3

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  1. Chapter 3 Matter and Energy

  2. Homework • Assigned Problems (odd numbers only) • “Questions and Problems” 3.1 to 3.41 (begins on page 61) • “Additional Questions and Problems” 3.49 to 3.69 (page 87-88) • “Challenge Questions” 3.71 and 3.75, (page 88)

  3. 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 • Any material we use such as water, wood, clothing, etc.

  4. Matter and 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 • Nearly all changes that matter undergoes involves the release or absorption of energy • Energyis the part of the universe that has the ability to do work

  5. Classification of Matter

  6. Pure Substance • Matter that has a definite and constant composition • Always contains the same substance, never varies • Either elements or compounds, all of one type • A pure sample of water only contains water molecules • Pure table salt contains only salt

  7. 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

  8. Compounds • 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)

  9. 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 • Not mixed in a fixed ratio, no chemical combination between the two substances

  10. 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 • Mixtures • Unlike compounds, mixtures can be separated by a physical process • Retain the properties of their individual components

  11. Types of Mixtures • Two types of mixtures: • 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. Types of 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 • i.e. copper (all elements are pure substances)

  13. Classification of Matter Physical Methods Chemical Methods

  14. Properties of Matter • Many properties used to identify chemical substances • Two types • Physical Properties • Chemical Properties • Properties can be: • Directly observable • The interaction of the matter with other substances

  15. States of Matter • Solid • Has a rigid, definite shape and definite volume • 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

  16. Physical Properties • Physical Properties • Characteristics of matter that can be observed or measured without changing its identity or composition • Characteristics that are directly observable • Color, odor, physical state, density, melting point, boiling point • Physical Changes • Cutting a piece of metal, melting ice

  17. Physical Change • A process that alters the appearance of a substance but does not change its identity or composition • No new substance is formed • Most common are changes of state

  18. Chemical Properties • Chemical Properties • Describes the ability of a substance to react and change into a new substance • Properties that matter exhibits as it undergoes changes in chemical composition • During a chemical change, the original substance is converted into one or more new substances with different chemical and physical properties

  19. Chemical Change • A change in the fundamental components of the substance: • A substance undergoes a change in chemical composition • Also called a chemicalreaction • Conversion of material(s) into one or more new substances • Wood burning, iron rusting, alka seltzer tablet into water

  20. 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)

  21. 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)

  22. 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)

  23. Temperature • A measure of how hot or cold a substance is compared to another substance • 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

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

  25. 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

  26. 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

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

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

  29. Energy • Capacity to do work or supply heat • Electrical, radiant, mechanical, thermal, chemical, nuclear • Two forms of Energy • Potential: Stored energy • Kinetic: Motion energy • All physical changes and chemical changes involve energy changes

  30. Forms of Energy • Potential energy: • Determined by an objects position • Chemical energy is potential energy 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 physical changes and chemical changes involve energy changes • These changes convert energy from one form to another

  31. 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

  32. Specific Heat • 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

  33. 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 substance to have a rise in temperature is called a substance’s “specific heat”

  34. 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 3.7, page 76 • Only for heating/cooling not for changes in state

  35. 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

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

  37. 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

  38. 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×

  39. Converting Energy Units • 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

  40. Calculating Mass Using Specific Heat • The 4184 J from the Pepsi One™ will heat how many grams of water from 0°C to boiling?

  41. Calculating Mass Using Specific Heat • How many grams of water would reach boiling if the water started out at room temperature (25°C)?

  42. Calculating The Temperature Change Using Specific Heat Values • If 50.0 J of heat is applied to 10.0 g of iron, by how much will the temperature of the iron increase?

  43. end

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