1. PACKET #2 Physical Behavior of Matter Textbook: Chapter 2 Reference Table: S & Periodic Table www.regentsprep.org

2. ENERGY The ability to do work. There are different forms; energy can be changed from one form to another. Chemical – released or absorbed in a chemical reaction. Electrical – energy from the flow of electricity (moving electrons) Electromagnetic Radiation – made of waves which include gamma, UV, light, radio, etc… Heat (thermal) – random motion of atoms/molecules (oil burns, chemical - > thermal)

3. ENERGY Mechanical – energy in moving objects, i.e. rock falling, car moving. Nuclear – the energy given off when a nucleus breaks down into smaller nuclei or when two nuclei are united to form a larger nucleus. Kinetic Energy: energy of movement. Potential Energy: stored energy.

4. LAW OF CONSERVATION OF ENERGY Energy may be converted from one form to another, but is never created or destroyed. Example: Lighting a match converts chemical energy to thermal energy (kinetic energy).

5. MATTER Anything that has mass and takes up space. Matter can be classified according to: 1. Physical state of matter (solid, liquid, gas) 2. Composition (element, compound, mixture)

6. Physical States of Matter A. Solid: (s) -Definite Shape; Definite Volume -All the solids have a crystalline structure or regular geometric pattern -Molecules are held tightly together, vibrating in a fixed position. Aqueous: (aq) -When a solid is completed dissolved in water

7. Physical States of Matter B. Liquid: (l) -NO Definite Shape; Definite Volume -Molecules are still close together, but not as close as with a solid -Particles in the liquid phase have no regular arrangement and are in constant motion, moving freely about their container

8. Physical States of Matter C. Gas (Vapor): (g) -NO Definite Shape; NO Definite Volume -Conform to the container they are in i.e. they are compressible -Molecules are far apart and move at very high speeds

10. COMPOSITION Composition of matter can be classified into two categories: SUBSTANCES & MIXTURES

11. Substance Any variety of matter that has the same (constant) properties and composition throughout. A substance is homogeneous which means that it is made up of one thing. Ex: bag of sugar.

12. There are two types of substances (both homogeneous/pure) ELEMENTS: Cannot be decomposed (broken down) into anything simpler by chemical change. Diatomic Molecule: two atoms of the same element bonded together. Ex: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 COMPOUNDS: made up of two or more different elements that are chemically united in a definite proportion. A compound can be decomposed (broken down). Binary Compounds: are composed of only two elements, Ex: NaCl

13. Mixtures Two or more substances that are mixed together, they are not united. Homogeneous mixtures (solutions/aqueous): considered like one thing, even though they are a mixture of two things. Example. Salt or sugar and water; the salt dissolves evenly into the water. Heterogeneous mixtures are mixtures that do not dissolve into each other like oil and water.

14. Techniques for Separating Mixtures 1. Filtration:  A mixture of solid/liquid substances can be separated using this method.  The mixture is poured through a filter.  The liquid components pass through; while the solid components remain on the paper.

15. Techniques for Separating Mixtures 2. Distillation:  Separation based on differing boiling points  A technique used to separate homogenous mixtures (solutions) Ex: Boil a solution of salt & water (water evaporates, salt is left behind) This evaporate is then typically condensed back into a liquid.

16. Techniques for Separating Mixtures 3. Chromatography:  A technique used to separate mixtures on the basis of the differing abilities of substances to adhere to the surfaces of various solids such as paper or starch

17. Techniques for Separating Mixtures 4. Magnet:  Separation on the basis of materials being influenced to different degrees by the presence of a magnetic field  Ex: Can be used to separate a mixture of iron filings & sulfur

18. DIFFERENCE BETWEEN COMPOUNDS AND MIXTURES: Compound – chemically united Mixture – not chemically united Compound – unite in a definite ratio; proportion Mixture – can be different proportions Compound – product has different property than the elements reacting Mixture – each substance maintains its own properties.

19. Composition Matter Pure Substances Mixtures Compounds Elements Heterogeneous Homogeneous

20. Physical Properties Physical properties describe the size, shape, color, texture, flexibility, etc. of an object. Ex.—color, size, shape, density, state of matter, melting point, boiling point, volume PHYSICAL CHANGES occur when physical properties of an object are altered WITHOUT changing the chemical make up. Ex. - melting ice, boiling water

21. Chemical Properties Chemical properties describe a substance’s ability to form new substances; its reactivity. Ex.—Na (sodium) reacts violently with water; Iron (Fe) reacts with oxygen in the air to produce rust. CHEMICAL CHANGES occur when a new substance(s) is made. Ex. - combustion (burning)

22. Identify the following as physical or chemical properties. ColorCombustibilityHardnessDensityMass Melting point Ductility Failure to react with other substances OdorWeightMalleability Tendency to corrode Volume

23. What are these?

24. Particle Diagrams Visual drawings to represent elements, compounds, or mixtures in their phases of matter. Each molecule represents an element; can create various combinations of elements, compounds, or molecules.

25. REMEMBER: Monatomic Molecules... Monatomic Molecules... Diatomic Molecules... Diatomic Molecules... Binary Compounds... Binary Compounds... Compounds vs. Mixtures... Compounds vs. Mixtures...

26. Element X = White Molecule Element Y = Black Molecule Element Z = Red Molecule X Y Z

27. #1: Create five (5) different types of elements. #2: Create three (3) different types of compounds.

28. #3: Create a mixture containing two (2) types of elements and two (2) types of compounds. #4: Create a mixture with one (1) monatomic element, one (1) diatomic molecule, and one (1) binary compound.

29. ENDOTHERMIC (heating) SOLID  LIQUID  GAS Solid  Liquid = melting Liquid  Gas = evaporation Solid  Gas = sublimation

30. EXOTHERMIC (cooling) GAS  LIQUID  SOLID Gas  Liquid = condensation Liquid  Solid = freezing (solidification) Gas  Solid = Deposition

31. Physical States of Matter Energy is absorbed or given off in a chemical reaction. Reactions involving heat energy are classified as: EXOTHERMIC: energy is given off in a chemical reaction. A+B  C + Heat ENDOTHERMIC: energy is absorbed in a chemical reaction. Heat + C  A+B Just remember... BARF (Break-Absorb, Release-Form)

32. ENTROPY A measure of the disorder of a system. The more disorder, the more entropy. Gas molecules are more disorderly than liquid molecules and therefore entropy increases as you go from the liquid to the gas phase. As you go from a gas to a liquid phase, entropy decreases (becomes less disorderly).

33. Entropy: A measure of the disorder of a system. The more disorder, the more entropy. Gas molecules are more disorderly than liquid molecules and therefore entropy increases as you go from the liquid to the gas phase. As you go from a gas to a liquid phase, entropy decreases (becomes less disorderly). Enthalpy: the amount of energy released or absorbed when a compound is formed from the elements in their standard state. Enthalpy can also be described at the heat of reaction, or as ∆H (the difference in heat content of the products and reactants.

34. Potential Energy (PE): Stored energy. Kinetic Energy (KE): Energy of movement. Energy is transferred from one form to another.  Potential Energy  Kinetic Energy

35. Temperature  The temperature of a substance is a measure of the average kinetic energy of its molecules  The higher the temperature the more kinetic energy the molecules have and the faster they move  The lower the temperature the less kinetic energy the molecules have the slower the molecules move  Boiling point and Freezing point are two fixed points on a thermometer

36. Heating Curve – Endothermic (absorbing heat)

37. Heating Curve - Endothermic Heating curves are endothermic because the phase change goes in the direction: solid  liquid  gas Slopes represent phases. Plateaus represent phase changes. There is a temperature change at the slopes. There is no temperature change at the plateaus. Phases (slopes): KE increases; PE remains the same. Phase Changes (plateaus): KE remains the same; PE increases.

38. Cooling Curve – Exothermic (releasing heat)

39. Cooling Curve - Exothermic Cooling curves are exothermic because the phase change goes in the direction: gas  liquid  solid Slopes represent phases. Plateaus represent phase changes. There is a temperature change at the slopes. There is no temperature change at the plateaus. Phases (slopes): KE deacrease; PE remains the same. Phase Changes (plateaus): KE remains the same; PE decreases.

40. Triple Point Phase Diagram Critical Temperature & Pressure: A gas normally liquefies at some point when pressure is applied to it.

41. Triple Point Phase Diagram Under appropriate conditions of temperature and pressure: Solid-liquid equilibrium/solid-vapor equilibrium can also exist Phase Diagram: Graphically portrays the conditions under which different states of equilibrium exist for a particular substance. We use these diagrams to predict the phase of a substance at a given temperature/pressure.

42. Any point on the curves themselves represents the pressure/temperature point at which equilibrium between two phases exists. Any temperature/pressure point that does not fall on line, corresponds to conditions in which only one phase is present. Normal melting/freezing point: Can be found on the graph by identifying it as the temperature point at which the curve AD intersects with 1 atm (760 torr) of pressure Normal boiling point: The temperature point at which the curve AB intersects with 1 atm (760 torr) of pressure. Point B = Critical Point: Represents what the critical temperature and critical pressure are for a particular substance. Point A = Triple Point: Represents the specific temperature/ pressure value at which all three phases exist in equilibrium. (Remember: Every other point on a curve would only represent equilibrium between 2 phases.) B A D C

43. H 2 O & CO 2

44. Practice Questions Referring to the figure below, describe any changes in the phases present when H 2 O is: (a) kept at 0°C while the pressure is increased from that at point 1 to that at point 5 (vertical line) (b) kept at 1.00 atm while the temperature is increased from that at point 6 to that at point 9 (horizontal line). Figure 11.28 Phase diagram of H 2 O.

45. More Using the diagram below, describe what happens when the following changes are made in a CO 2 sample initially at 1 atm and –60°C: (a) Pressure increases at constant temperature to -60°C. (b) Temperature increases from –60ºC to –20ºC at constant 60 atm pressure.

46. Heat Equations: (Table T) q = mC∆T q = mH f q = mH v q = Heat (in joules) m = Mass (in grams) C = Specific Heat Capacity (Table B): the amount of heat required to raise 1 gram of something 1 degree Celsius (ex: 4.18 joules/g°C)  T = Change in temperature (Final - Initial) If q is negative: exothermic reaction (release heat) If q is positive: endothermic reaction (absorb heat) H f = Heat of fusion (energy required to freeze/melt) H v = Heat of vaporization (energy required to boil/condense) If the question is asking for Calories instead of Joules, C for water is 1cal/g°C

47. Table B represents physical constants for water only. If a heat equation question is asking about anything other than water, the constant for the substance will have to be provided (C, H f,H v ). The plateau for vaporization/condensation is longer than melting/freezing because the H v is greater than the H f, and therefore it takes longer to add or release heat.

48. Practice Problems EXAMPLE: Calculate the amount of heat required to raise the temperature of 15g of water from 20˚C to 50˚C. EXAMPLE: When 20 joules of heat are added to 2g of water at 15˚C, the temperature of the water increases to?

49. More... Example: How much heat energy is absorbed when 10g of ice melts to form liquid water at the same temperature? Example: How much heat is absorbed when 70.0g of water is completely vaporized at its boiling point?

50. 1. When 20.0g of a substance is completely melted at its melting point, 3444 J are absorbed. What is the heat of fusion of this substance? 2. The heat of vaporization of a liquid is 1344 J/g. What is the minimum number of joules needed to change 40.0g of the liquid to vapor at the boiling point? 3. How much energy is required to vaporize 10.00g of water at is boiling point? 4. At 1 atmosphere of pressure, 25.0g of a compound at its normal boiling point is converted to a gas by the addition of 34,400 J. What is the heat of vaporization for this compound in J/g? 5. The heat of fusion of a compound is 30.0 joules per gram. What is the number of joules of heat that must be absorbed by a 15.0g sample to change the compound from solid to liquid at its melting point?

51. Hess’ Law Hess’ Law states that if a reaction is carried out in a series of steps, ∆H for the reaction will be equal to the sum of the enthalpy changes for the individual steps.

52. Hess’ Law A + B → X + Y Δ H 1 = 25 kJ X + Y → C + D Δ H 2 = -78 kJ (X + Y cancel one another out) Net Reaction: A + B → C + D Δ H = Δ H 1 + Δ H 2 = 25 + (-78) = -53 kJ

53. Practice Problems 1. N 2 + O 2 → 2NO Δ H 1 = 180 kJ 2NO + O 2 → 2NO 2 Δ H 2 = -112 kJ 2. C graphite → C diamond C graphite + O 2 → CO 2 Δ H 1 = -394 kJ C diamond + O 2 → CO 2 Δ H 2 = -396 kJ 3. N 2 + 2O 2 → 2NO 2 Δ H 1 = 67.7 kJ N 2 + 2O 2 → N 2 O 4 Δ H 2 = 9.7 kJ

54. Gibbs Free Energy (G) The energy of a system available to do work. The change in Gibbs free energy, ∆G, is determined for a given reaction from the equation: ∆G = ∆H - T ∆S ∆G = change in Gibbs free energy ∆H = change in enthalpy T = absolute temperature ∆S = change in entropy

55. Gibbs Free Energy (G) ∆G is used to predict the spontaneity of a reaction A negative ∆G denotes a spontaneous reaction A positive ∆G denotes a nonspontaneous reaction If ∆G is zero, the system is in a state of equilibrium ∆H∆SOutcome -+Spontaneous at all temperatures +-Nonspontaneous at all temperatures ++Spontaneous only at high temperatures --Spontaneous only at low temperatures

56. REGENT’S REVIEW QUESTIONS 1) Which substance represents a compound? A) C(s) B) CO(g) C) Co(s) D) O2(g) 2)

57. 3) Which substance can not be decomposed by a chemical change? A)Ne B) N2O C) HF D) H2O 4)

58. 5) Which statement describes a chemical property of iron? A) Iron can be flattened into sheets. B) Iron combines with oxygen to form rust. C) Iron conducts electricity and heat. D) Iron can be drawn into a wire. 6) Which of the following is an example of a physical change in matter? A) fizzing produced when magnesium metal is added to acid B) magnesium metal burning with a bright white flame C) melting of sodium metal D) sodium metal exploding in water

59. 7) Consider the reaction: H 2 O (l) + energy  H 2 (g) + ½ O 2 (g). Which of the following phrases best describes this reaction? A) endothermic, releasing energy B) exothermic, releasing energy C) exothermic, absorbing energy D) endothermic, absorbing energy 8) Which phase change is endothermic? A) Fe(l)  Fe(s) B) H 2 O (l)  H 2 O (s) C) CO 2 (s)  CO 2 (g) D) NH 3 (g)  NH 3 (l)

60. 9) Which substance can be decomposed by chemical means? A) ammonia B) oxygen C) silicon D) phosphorus 10) Which particle model diagram represents only one compound composed of elements X and Z?

61. 11) Which of the following two substances can not be broken down by chemical change? A) C and CuO B) C and Cu C) CO 2 and Cu D) CO 2 and CuO 12) Which of the following particle diagrams represents a mixture of element X and element Z, only?

62. 13) Which statement describes a chemical property of the element magnesium? A) Magnesium conducts electricity. B) Magnesium is malleable. C) Magnesium reacts with an acid. D) Magnesium has a high boiling point. 14) Solid ZnCl2 and liquid ZnCl2 have different A) empirical formulas B) ion ratios C) physical properties D) formula masses

63. 15) What type of mixture is represented by X in the diagram shown? 16) What type of substance is represented by Z in the diagram shown? 17) Given a mixture of sand and water, state one process that can be used to separate water from the sand. Questions 15 through 17 refer to the following:

64. Review Questions 1) Given the balanced equation representing a reaction: CH 4 (g) + 2O 2 (g)  2H 2 O(g) + CO 2 (g) + heat Which statement is true about energy in this reaction? A) The reaction is exothermic because it releases heat. B) The reaction is endothermic because it releases heat. C) The reaction is endothermic because it absorbs heat. D) The reaction is exothermic because it absorbs heat.

65. 2) Given the balanced equation representing a reaction: Cu + S  CuS + energy Which statement explains why the energy term is written to the right of the arrow? A) The compound CuS is composed of two metals. B) Energy is released as the bonds in CuS form. C) The compound CuS is composed of two nonmetals. D) Energy is absorbed as the bonds in CuS form. 3) At which temperature would atoms of a He(g) sample have the greatest average kinetic energy? A) 25°C B) 273 K C) 298 K D) 37°C 4) Which Kelvin temperature is equal to 56°C? A) -217 K B) 329 K C) 217 K D) -329 K

66. 5) A temperature of 37°C is equivalent to a temperature of A) 310. K B) 98.6 K C) 371 K D) 236 K 6) At which Celsius temperature does lead change from a solid to a liquid? A) 874°C B) 0°C C) 328°C D) 601°C 7) What term refers to the difference between the potential energy of the products and the potential energy of the reactants for any chemical change? A) heat of fusion B) heat of reaction C) heat of deposition D) heat of vaporization

67. 8) For a given reaction, adding a catalyst increases the rate of the reaction by A) providing an alternate reaction pathway that has a lower activation energy B) providing an alternate reaction pathway that has a higher activation energy C) using the same reaction pathway and decreasing the activation energy D) using the same reaction pathway and increasing the activation energy 9) If 4.0 grams of water at 1.0°C absorbs 33 joules of heat, what will be the change in temperature of the water? A) 3.0°C B) 4.0°C C) 2.0°C D) 1.0°C

68. 10) A 5.00 gram sample of water is heated and the temperature rises from 10.0°C to 15.0°C. What is the total amount of heat energy absorbed by the water? A) 84 J B) 42.0 J C) 105 J D) 21.0 J 11) What is the total number of joules of heat energy released when 20.0 grams of water is cooled from 20.0°C to 10.0°C? A) 836 J B) 200. J C) 30.0 J D) 83.6 J 12) What is the total number of calories of heat energy absorbed when the temperature of 200 grams of water is raised from 10°C to 40°C? A) 200 cal B) 8,000 cal C) 30 cal D) 6,000 cal

69. 13) What is the total number of joules released when a 5.00-gram sample of water changes from liquid to solid at 0°C? A) 1,670 J B) 2,260 J C) 334 J D) 11,300 J 14) A hot pack contains chemicals that can be activated to produce heat. A cold pack contains chemicals that feel cold when activated. Based on energy flow, state the type of chemical change that occurs in a hot pack.

70. 15) The graph below represents the uniform heating of a substance, starting below its melting point, when the substance is solid. Which line segments represent an increase in average kinetic energy? A) AB and BC B) AB and CD C) DE and EF D) BC and DE

71. 16) Given the heating curve where substance X starts as a solid below its melting point and is heated uniformly: Identify the process that takes place during line segment DE of the heating curve shown.

72. 17) The graph below represents the uniform heating of a substance, starting with the substance as a solid below its melting point. Which line segment represents an increase in potential energy and no change in average kinetic energy? A) CD B) EF C) BC D) AB