1. Unit 2: Nature of Matter and Kinetic Theory

2. Part 1: The Nature of Matter

3. properties = characteristics and behavior of matter (includes changes that matter undergoes). What color is it? Is it solid, liquid or gas Is it reactive? structure = composition what matter is made of how matter is organized. How do we classify matter?

4. Examples of physical properties : solubility, - dissolves in water? melting point, boiling point color, density, electrical conductivity, physical state (solid, liquid, or gas).

5. physical change - change in matter that does not involve a change in the chemical identity Change of state is a physical change:

6. Classify by purity Is it a pure substance or mixture? Pure substance = sample of matter that has definite chemical and physical properties, can be either an element or a compound Classifying Matter

8. compound = pure substance that can be broken down into simpler substances. element = substance that cannot be broken down into simpler substances. Element or Compound? salt gold

9. Compounds Are More Than One Element formula = combination of the chemical symbols that show what elements make up a compound and the number of atoms of each element CompoundFormula caffeine salt water C 8 H 10 N 4 O 2 NaCl H 2 O

10. Compounds Are More Than One Element ****The properties of the compound are different from the properties of the elements that compose the compound. silver + bromine = silver bromide

11. substance is not changed = no fixed composition the basic identity of each Mixture = made up of different kinds of matter Pure Substance or Mixture?

12. Homogeneous mixtures are the same throughout. Also known as a solution. Pure substance or a mixture?

13. When you dissolve sugar in water, sugar is the solute—the substance being dissolved. The substance that dissolves the solute is the solvent. in this case it is water solute + solvent = solution When the solvent is water, the solution is called an aqueous solution.

14. heterogeneous mixture is one with different compositions, depending upon where you look Pure Substance or Mixture?

16. Pure Substance Mixture element compound homogenous heterogeneou s Matter

17. a substance must be separated chemically a mixture can be separated physically

18. An example of a pure substance in everyday life is _____. a. pond water b. a cola drink c. sugar d. concrete c. sugar

19. A soft drink is an example of a(n) _____. a. compound b. heterogeneous mixture c. element d. homogeneous mixture d. homogenous mixture

20. Identify each of the following as either a compound or a mixture. A. sand B. water C. juice mixture compound mixture

21. In ocean water, salt is a(n) _____. a. alloy b. solution c. solute d. solvent c. solute

22. pure substance? element or compound? a mixture? Heterogeneous or homogenous Aluminum foil Pure substance, element

23. pure substance? element or compound? a mixture? Heterogeneous or homogenous bowl of cereal mixture, heterogeneous

24. pure substance? element or compound? a mixture? Heterogeneous or homogenous whipped cream Mixture, homogenous

25. pure substance? element or compound? a mixture? Heterogeneous or homogenous oil and vinegar dressing Mixture, heterogeneous

26. pure substance? element or compound? a mixture? Heterogeneous or homogenous aspirin - acetylsalicylic acid Pure substance, compound

27. pure substance? element or compound? a mixture? Heterogeneous or homogenous orange juice with pulp Mixture, heterogeneous

28. pure substance? element or compound? a mixture? Heterogeneous or homogenous gold Pure substance, element

29. pure substance? element or compound? a mixture? Heterogeneous or homogenous salt Pure substance, compound

30. pure substance? element or compound? a mixture? Heterogeneous or homogenous peanut butter Mixture, homogenous

31. Chemical Properties Chemical properties are those that can be observed only when there is a change in the composition of the substance. Rusting is a chemical reaction in which iron combines with oxygen to form a new substance, iron oxide. Examples of chemical property: flammability reactivity

32. Chemical Changes chemical change - the change of one or more substances into other substances. A chemical property always relates to a chemical change = chemical reaction.

33. - production of bubbles - release or absorption of energy - color change ***only way to be sure is to check the composition of the sample before and after the change. Clues that a chemical change has occurred:

34. Below are listed changes that can be observed in everyday life. Tell whether it is a physical change or a chemical change. 1.an icicle melting 2.charcoal burning 3.magnetizing a piece of steel 4.iron rusting 5.rubbing alcohol evaporating from the skin physical change chemical change physical change chemical change physical change

35. chemical change involves only a rearrangement of the atoms. Atoms DO NOT just appear or disappear. ******Law of Conservation of Mass****** In a chemical change, matter is neither created nor destroyed. Chemical Reactions

36. Chemical Reactions and Energy All chemical changes also involve some sort of energy change. Energy is either taken in or given off as the chemical change takes place. Energy is the capacity to do work. Work is done whenever something is moved.

37. Chemical Reactions and Energy Energy is also produced and released in the form of heat and light. Chemical reactions that GIVE OFF heat energy are called exothermic reactions.

38. Chemical reactions that ABSORB heat energy are called endothermic reactions. Chemical Reactions and Energy

39. Classify each of the following as a chemical or physical property. A. density B. reactivity C. color D. melting point physical property chemical property physical property

40. Part 2: The Kinetic Theory States of Matter – solid – liquid – gas – plasma

41. Intermolecular Forces (IMF) Attractive forces between molecules.  Much weaker than chemical bonds within molecules.

42. The Kinetic Theory of Matter 1. Matter is composed of PARTICLES. 2. Particle movement is rapid, constant, and random (Brownian motion)

43. The Kinetic Theory of Matter 3.All collisions are perfectly ELASTIC (NO energy lost).

44. Kinetic theory of matter Kinetic energy (K.E.) = energy of motion gases have the least restriction on motion – have the most K.E. solids have the most restriction on motion – have the least K.E.

45. Kinetic model of gases Gases: matter with variable shape and variable volume Gas particles move in a straight line until they collide with container or each other

46. Kinetic model of liquids Liquids: matter with variable shape and definite volume Particles slide past each other but are so close together they do not move in a straight line

47. Kinetic model of solids solids: matter with definite shape and definite volume Particles cannot move past each other, they are in constant motion bouncing off neighbors

48. Other forms of matter Plasmas - gaseous mixture of ions - exists at high temperatures most common form of matter in the universe but least common on Earth itself

49. Plasmas continued an ionized gas that conducts electricity -forms at very high temps when matter absorbs energy and breaks apart The sun is made of plasma - also found in fluorescent lights

50. Temperature and kinetic energy temperature—the measure of the average K.E. of particles in a sample Kelvin (K) – SI base unit of temperature; measures average K.E.

51. Temperature and kinetic energy When temp increases, particle motion increases. When temp decreases, particle motion decreases. A temp of 300 K has twice the kinetic energy as 150 K.

52. Temperature and kinetic energy 0 Kelvin = absolute zero = no molecular motion No degrees sign ( ° ) is used with Kelvin numbers There will never be negative numbers for Kelvin temperatures!.

53. densitycompressibilityintermolecular forces solidmost densedifficult to compress strong liquid gasleast denseeasily compressed weak Comparing solids, liquids, and gases

54. Kinetic energyspace between particles organization solidleast amount of kinetic energy very little space between particles most organized liquid gasmost amount of kinetic energy a lot of space between particles least organized Comparing solids, liquids, and gases

55. Changing states and energy changes Going from a more energetic state (gas) to a less energetic state (solid) requires a release of energy – exothermic

56. Going from less energetic (solid) to more energetic (gas) requires absorption of energy -- endothermic

57. Vapor Pressure and boiling Vapor Pressure Vapor Pressure - pressure of vapor above a liquid at equilibrium high vapor pressure = volatile volatile = easily evaporates The greater the fraction of molecules which can escape the liquid, the greater the vapor pressure

58. What happens to the vapor pressure if you increase the temperature of a liquid in a closed container? – causes the vapor pressure above the liquid to increase.

59. equilibrium vapor pressure - when the number of vapor molecules rejoining the water equals the number leaving to go into the vapor phase

60. If there is equilibrium between the liquid state and the gas state, what is true about the rate of evaporation and the rate of condensation? They are equal

61. Vapor pressure and boiling point Boiling Point - temp at which v.p. of liquid equals external pressure -depends on atmospheric pressure & IMF Normal B.P. - b.p. at 1 atm

62. Effects of Intermolecular Forces (IMF) When IMF’s are weak – vapor pressure is high vapor pressure is high – volatility is high – boiling point is low

63. Heat of Fusion Melting point – temp of a solid when it becomes a liquid = freezing point (temp when liquid becomes a solid)

64. B. Heating Curves Freezing/Melting point Solid Liquid Boiling point Gas

65. Heating Curves IMPORTANT: temp does not change during the actual phase change. Increasing the temp will only make the change happen faster.

66. Phase Diagrams Shows the phases of a substance at different temps and pressures.

67. triple point -the point (temperature and pressure) on a phase diagram at which three phases of a substance can coexist. All six phase changes can occur at the triple point: freezing and melting, evaporation and condensation, sublimation and deposition.

68. Phase Diagrams critical point -at extremely high temperatures and pressures, the liquid and gaseous phases become indistinguishable, in what is known as a supercritical fluid

69. The End!