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

5. Classifying Matter According to Its State Liquid Has an indefinite shape and a definite volume. It will take the shape of the container it fills

6. Classifying Matter According to Its State 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 chlorine gas

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

8. 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

9. Pure Substances Matter that has a definite and constant composition is a pure substance Composed of the same substance; no variation There are two classifications of pure substances: Elements Compounds Over 10 million pure substances have been isolated: About 100 are elements (naturally occurring) The rest are compounds

10. Pure Substances Elements Substances which can not be broken down into simpler substances by chemical reactions Fundamental (simplest) substances (building blocks of matter) Each has its own specific properties For example: copper, gold, and lead (consist of only one type of atom)

11. Pure Substances Compounds Composed of two or more elements combined chemically in a definite and constant ratio (e.g., water, table salt, sugar) Compounds can be broken down into elements (simpler substances) by chemical processes

12. Pure Substances Most of matter is in the compound form Properties of the compound not related to the properties of the elements that compose it Water, for example, is composed of hydrogen and oxygen gases combined in a 2:1 ratio

13. Mixtures Mixtures result from the physical combination of two or more substances (elements or compounds) Mixtures can be separated by physical methods Each substance (after separation) retains its identity because the substances are not chemically mixed (united) with each other in fixed proportions Unlike compounds, mixtures do not exist in fixed proportions Mixtures of the same components can vary in composition

14. 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 Many are referred to as solutions Heterogeneous mixture: Composition is not uniform throughout the sample. It contains visibly different parts or phases

15. Examples of Mixtures Homogenous mixtures corn oil tea the air we breath (with no clouds present) seawater Heterogeneous mixtures clouds oil and vinegar trail mix sand triglycerides caffeine flavenoids fluoride

16. Compounds vs. Mixtures Compounds are not mixtures Compounds cannot be separated by a physical process because the two or more elements are bound together Compounds can be subdivided by a chemical process into two or more simpler substances Simpler substances have different properties from the compound itself Mixtures Unlike compounds, mixtures can be separated by a physical process Once separated, each substance in a mixture retains its own individual properties

17. Classification of Matter Chemical Methods Physical Methods Pure Substances Mixture Matter Compounds Elements Homogeneous Mixture Heterogeneous Mixture Pure Substances

18. 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)

19. 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 viewed or measured without changing its identity or composition For example: color, physical state, density, hardness, melting point, boiling point

20. Chemical Properties A chemical property describes the way a substance undergoes a change or resists change to form a new substance Chemical properties cannot be determined by inspection; these properties only become evident during a chemical reaction It describes a substances potential to undergo a chemical reaction due to its composition It also describes a substances failure to undergo a chemical reaction in the presence of another substance

21. Chemical Properties Chemical properties include flammability, toxicity, the ability to react with oxygen (rusting) Hydrogen has the potential to ignite and burn (high flammability). Mercury is a highly toxic element which naturally occurs and is introduced as a contaminant in the environment. Objects made from copper will undergo corrosion (turn green) when exposed to oxygen and moist air for long periods. Objects made from gold objects will resist corrosion when exposed to oxygen and moist air for long periods. Water will react intensely with sodium metal and produces hydrogen gas and heat. In each case (above) the atoms in compounds are rearranged to make new and different compounds.

22. Classifying Properties The boiling point of ethyl alcohol is 78 °C Physical property – describes an inherent characteristic of alcohol: its boiling point Diamond is a very hard substance Physical property – describes an inherent characteristic of diamond: its hardness Propane is a flammable substance Chemical property – describes the ability to burn (react with oxygen) and form new substances: carbon dioxide and water Ammonia is a colorless gas at room temperature Physical property – describes two inherent characteristics of ammonia: its color and physical state

23. 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

24. 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

25. 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, an 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

26. Classifying Changes Melting of snow Physical change – a change of state but not a change in its composition Burning of gasoline Chemical change – reacts (combines with oxygen) to form new compounds: carbon dioxide and water Dissolving salt in water Physical change – only a change in the form of the substance but not a change in its composition

27. Classifying Changes Melting of iron metal Physical change – describes a state change, but the material is still iron Rusting of iron metal Chemical change – combines with oxygen to form a new reddish- colored substance (ferric oxide)

28. Conservation of Mass During a physical change: No new substance is formed During a chemical change: At least one new substance is formed

29. Conservation of Mass 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 nor destroyed

30. 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

31. Energy Energy is 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

32. 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

33. 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

34. 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

35. Energy: Chemical and Physical Change 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 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

36. 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

37. 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

38. 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

39. 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

41. 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

42. 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? K = °C + 273 25 ºC + 273 = 298 K 0 °C = 273 K add 273 to the °C value

43. The Temperature Scales Temperature is the quantity that indicates how warm or cold an object is relative to some standard It is expressed by a number that corresponds to the degree of hotness on a chosen scale: Fahrenheit, Celsius, or Kelvin The Celsius and Kelvin scales are part of the metric system and the Fahrenheit scale belongs to the English measurement system The various temperature scales are a result of the different degrees sizes and different reference points

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

45. 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

46. 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 (PE component only) As heat energy (KE component only) increases the temperature increases

47. Temperature Changes: Heat Capacity Heat energy is the 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

48. 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”

49. 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 of state

50. 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

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

52. Specific Heat Equation Energy (heat) required to change the temperature of a substance depends on: The amount of substance being heated (g) The temperature change (initial T and final T in °C) The identity of the substance

53. 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×

54. 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 cal4.184 J = 1 cal

55. Energy and Heat Capacity Calculations How many grams of water will the 4184 J (from a can of Pepsi One™) heat from 0°C to boiling? TfTf TiTi d water =1.0 g/mL

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

57. Energy and Heat Capacity Calculations If 50.0 J of heat is applied to 10.0 g of iron, by how much will the temperature of the iron increase? Solve for ΔT

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

60. Energy: Chemical and Physical Change 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 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 reactants products reactants products