1. Mark S. Cracolice Edward I. Peters http://academic.cengage.com/chemistry/cracolice Mark S. Cracolice The University of Montana Chapter 2 Matter and Energy

2. Representations of Matter Goal 1 Identify and explain the differences among observations of matter at the macroscopic, microscopic, and particulate levels. Goal 2 Define the term model as it is used in chemistry to represent pieces of matter too small to see.

3. Representations of Matter Anything that has mass (sometimes expressed as weight) and takes up space is called matter. Matter can be observed and/or thought about at different levels: Macroscopic Microscopic Particulate

4. Representations of Matter Macroscopic Samples of Matter: Mountains Rocky cliffs Huge boulders Rocks and stones Gravel Sand Macro- means large

5. Representations of Matter Microscopic Samples of Matter: Tiny animals or plants Cells Crystals on rock surfaces Micro- means small

6. Representations of Matter Particulate Samples of Matter: Too small to see, even with the most powerful optical microscope. Chemists imagine the nature of the behavior of the tiny particles that make up matter, based on experimental measurements and models of matter, and they use that knowledge to carry out changes from one type of matter to another.

7. Representations of Matter Macroscopic, Microscopic, and Particulate Matter:

8. Representations of Matter Model: A representation of something. Geologists model the earth (globe). Biologists model cells. Chemists model atoms and molecules.

9. Representations of Matter Ball-and-Stick Model: Symbolizes atoms as balls and the electrons that connect those atoms as sticks. Space-Filling Model: Shows the outer boundaries of the particle in three-dimensional space.

10. Representations of Matter

11. Models are represented in writing with symbols. Chemical symbols are letters that represent atoms of elements. H represents an atom of hydrogen. O Represents an atom of oxygen. H 2 O represents a molecule of water: Two hydrogen atoms and one oxygen atom.

12. Representations of Matter Chemists make mental transformations between visible macroscopic matter (the liquid water) and models of the particulate-level molecules that make up the matter (red and white space-filling models). Written symbols (H–O–H on the board) serve as simpler representations of the particulate-level models.

13. States of Matter Goal 3 Identify and explain the differences among gases, liquids, and solids in terms of (a) visible properties, (b) distance between particles, and (c) particle movement.

14. States of Matter Familiar examples of the states of matter: The air you breathe is a gas. The water you drink is a liquid. The food you eat is a solid.

15. States of Matter Kinetic Molecular Theory: All matter consists of extremely tiny particles that are in constant motion. Kinetic refers to motion. Molecular comes from molecule, the smallest unit particle that can exist independently and possess the identity of the substance. Theory is a hypothesis that has been tested and confirmed by many experiments.

16. States of Matter The speed at which particles move is faster at higher temperatures and slower at lower temperatures. There is an attraction among particles in all samples of matter. The state of matter of any sample depends on temperature (the speed at which particles move) and the attractions among the particles that make up the sample.

17. States of Matter Particulate-Level Behavior of States of Matter: Gas Particles are independent of one another, moving in random fashion Liquid Particles move freely among themselves, but clump together Solid Particles vibrate in fixed positions relative to one another

18. States of Matter

19. Properties and Changes Goal 4 Distinguish between physical and chemical properties at both the particulate level and the macroscopic level. Goal 5 Distinguish between physical and chemical changes at both the particulate level and the macroscopic level.

20. Properties and Changes Physical Properties Description as detected by senses: Color, shape, odor, etc. Measurable properties: Density, boiling point, etc. Examples: Charcoal is black Glass is hard The normal boiling point of water is 100°C

21. Properties and Changes Physical Changes Alteration of the physical form of matter without changing its chemical identity No new substance formed Examples: Ice melts to liquid water Dry ice changes to gaseous carbon dioxide A rock is ground into sand

22. Properties and Changes In a Physical Change, the Molecules are Unchanged:

23. Properties and Changes Chemical Changes Chemical identity of a substance is destroyed A new substance is formed Examples: Water decomposes to hydrogen and oxygen gases when subjected to an electrical current Iron rusts Food is digested

24. Properties and Changes In a Chemical Change, the Molecules Change:

25. Properties and Changes

26. Chemical Properties The types of chemical change a substance is able to participate in. Examples: A chemical property of water is that it can be decomposed to its elements when subjected to an electrical current. A chemical property of iron is that it will rust under certain conditions. A chemical property of starch is that it reacts to form sugar during digestion.

27. Properties and Changes ChemicalPhysical Changes Chemical identity of a substance is destroyed New substances formed New form of same substance No new substances formed Properties Types of chemical changes possible Description as detected by the senses Measurable properties

28. Substances and Mixtures Goal 6 Distinguish between a pure substance and a mixture at both the macroscopic level and the particulate level. Goal 7 Distinguish between homogeneous and heterogeneous matter.

29. Substances and Mixtures Pure Substance A sample consisting of only one kind of matter, either compound or element; made up entirely of one kind of particle. Unique set of physical and chemical properties. Cannot be separated into parts by a physical change.

30. Substances and Mixtures Mixture A sample of matter that consists of two or more substances. Physical and chemical properties of a mixture vary with different relative amounts of the parts. Can be separated into parts via physical processes.

31. Substances and Mixtures Pure water has a constant boiling point—a physical property. The boiling point of a mixture (solution) changes as the composition of the mixture changes.

32. Substances and Mixtures You cannot distinguish a pure substance from a mixture of uniform appearance by observation alone at the macroscopic level.

33. Substances and Mixtures Solution A homogeneous mixture of two or more components. Homogeneous A sample that has uniform appearance and composition throughout. Examples: Tea, paint, gasoline Heterogeneous A sample with different phases, usually visible. Examples: Carbonated beverages, salad dressings

34. Substances and Mixtures Homogeneous Matter may be Either a Pure Substance or a Mixture:

35. Separation of Mixtures Goal 8 Describe how distillation and filtration rely on physical changes and properties to separate components of mixtures.

36. Separation of Mixtures Most natural substances are mixtures. Separation processes are an important part of chemistry. Nitrogen and oxygen are separated from the mixture called air. Pure water is separated from the mixture called natural water. Gasoline is separated from the mixture called crude oil.

37. Separation of Mixtures A Physical Property, Magnetism, Allows a Mixture of Iron and Sulfur to be Separated:

38. Separation of Mixtures Distillation Separation of the parts of a mixture by heating a liquid solution until one component boils, changing into the gaseous state. The pure substance in the gaseous state is then collected and cooled into the liquid state. Boiling is a physical change. Distillation allows components in a homogeneous mixture to be separated into one or more pure substances.

39. Separation of Mixtures Laboratory Distillation Apparatus:

40. Separation of Mixtures Filtration Separation of the components of a mixture by physical means by using a porous medium, such as filter paper, to separate components based upon relative particles sizes. Filtration is based on the physical properties of a mixture: The particle sizes of a component to be separated must be significantly larger or smaller than the pore size of the filtration medium.

41. Separation of Mixtures Gravity Filtration

42. Elements and Compounds Goal 9 Distinguish between elements and compounds. Goal 10 Distinguish between elemental symbols and the formulas of chemical compounds. Goal 11 Distinguish between atoms and molecules.

43. Elements and Compounds Element Pure substance that cannot be decomposed into other pure substances by ordinary chemical means. Atom Smallest particle of an element that can combine with atoms of other elements to form chemical compounds. Compound Pure substance that can be broken down into two or more other pure substances by a chemical change.

44. Elements and Compounds The Element Silver and a Particulate-Level Model of Silver Atoms:

45. Elements and Compounds Mixtures are separated into pure substances by physical means; compounds are separated into pure substances by chemical changes.

46. Elements and Compounds Elements At least 88 elements occur in nature. Examples: copper, sulfur, gold, silver 11 elements occur in nature as gases; 2 occur as liquids (mercury and bromine); the others occur as solids. Name of an element is always a single word; compound names are usually two words or a polysyllabic compound word.

47. Elements and Compounds Major Elements of the Human Body ElementPercentage Composition by Number of Atoms Hydrogen63.0 Oxygen25.5 Carbon 9.45 Nitrogen 1.35 These four elements make up 99.3% of the atoms in your body.

48. Elements and Compounds Familiar Objects that are Composed of Nearly Pure Elements:

49. Elements and Compounds Familiar Objects that are Compounds:

50. Elements and Compounds Elemental Symbols Letters that symbolize elements The first letter of the name of the element, written in uppercase, is often its symbol. Examples: Hydrogen, H Oxygen, O Carbon, C If more than one element begins with the same letter, a second lowercase letter is added. Examples: Helium, He Osmium, Os Chlorine, Cl

51. Elements and Compounds Chemical Formulas Symbolic representations of the particles of a pure substance. A combination of the symbols of all the elements in a substance. The formula of most elements is the same as the symbol of the element, e.g., helium He; sodium, Na. Other elements exists in nature as molecules and their formulas indicate the number of atoms of the element in the molecule, e.g., hydrogen, H 2 ; oxygen, O 2.

52. Elements and Compounds Formula Unit Molecule or simplest ratio of particles for non-molecular species. Ammonia molecules have the formula NH 3 : 1 atom of nitrogen and 3 atoms of hydrogen. Magnesium chloride exists as an orderly, repeating pattern of magnesium and chlorine in a 1:2 ratio; Its formula unit is MgCl 2.

53. Elements and Compounds Law of Definite Composition or Law of Constant Composition Any compound is always made up of elements in the same proportion by mass (weight). No matter its source, water (H 2 O) is 11.1 parts hydrogen per 88.9 parts oxygen.

54. Elements and Compounds The Properties of a Compound are Different from the Properties of the Elements that Make Up the Compound: Water, H 2 O Liquid at 25°C, melts at 0°C, boils at 100°C Hydrogen, H 2 Gas at 25°C, melts at –259°C, boils at –253°C Oxygen, O 2 Gas at 25°C, melts at –219°C, boils at –183°C

55. Elements and Compounds Particulate and Macroscopic Views of Elements and Compounds:

56. Elements and Compounds Particulate and Macroscopic Views of Elements and Compounds:

57. Elements and Compounds Particulate and Macroscopic Views of Elements and Compounds:

58. Elements and Compounds Summary of the Classification System for Matter:

59. Electrical Character of Matter Goal 12 Match electrostatic forces of attraction and repulsion with combinations of positive and negative charge.

60. Electrical Character of Matter Two of the fundamental forces that govern the operation of the universe are: Force of gravity Electromagnetic force The electromagnetic force plays an important role in understanding chemistry. It includes electricity and magnetism.

61. Electrical Character of Matter Force Field: Region in space where the force is effective. Electrostatic Force: The force of an electrical charge that does not move. A charged object exerts an invisible electrostatic force.

62. Electrical Character of Matter Experimental evidence shows that: There are only two types of electrical charge, positive and negative. Two objects having the same charge, both positive or both negative, repel each other. Two objects having unlike charges, one positive and one negative, attract each other.

63. Electrical Character of Matter Electrostatic forces show that matter has electrical properties. These forces are responsible for the energy absorbed or released in chemical changes.

64. Chemical Change Goal 13 Distinguish between reactants and products in a chemical equation. Goal 14 Distinguish between exothermic and endothermic changes. Goal 15 Distinguish between potential energy and kinetic energy.

65. Chemical Change Chemical Equation A symbolic representation of chemical change, with the formulas of the beginning substances to the left of an arrow that points to the formulas of the substances formed. Reactant Original substance Product Substance formed as a result of chemical change 2 H 2 O 2 H 2 + O 2 Reactant Products

66. Chemical Change Exothermic Reaction A chemical change that releases energy to its surroundings. Example: Burning charcoal C + O 2 CO 2 + energy

67. Chemical Change Endothermic Reaction A chemical change that absorbs energy from its surroundings. Example: Decomposition of water to its elements 2 H 2 O + energy 2 H 2 + O 2

68. Chemical Change Energy The ability to do work. Potential Energy Energy due to position in a field where forces of attraction and/or repulsion are present. Gravitational Potential Energy Position in the earth’s gravitational field. Electrical Potential Energy Position in an electrical field.

69. Chemical Change Minimization of energy: One of the driving forces that cause chemical reactions to occur. Chemical energy comes largely from the rearrangement of charged particles in an electrostatic field.

70. Chemical Change Kinetic Energy Energy of motion The temperature of an object is proportional to the average kinetic energy of its particles.

71. Conservation Laws Goal 16 State the meaning of, or draw conclusions based on, the Law of Conservation of Mass. Goal 17 State the meaning of, or draw conclusions based on, the Law of Conservation of Energy.

72. Conservation Laws The Conservation Law In any change, the sum of mass plus energy is conserved; they are neither created nor destroyed. ∆E = ∆m  c 2 Matter is an extremely concentrated form of energy.

73. Conservation Laws Law of Conservation of Mass In a nonnuclear change, mass is conserved; it is neither created nor destroyed. Mass may change form, however. In any ordinary chemical change, Total mass of reactants = Total mass of products

74. Conservation Laws Law of Conservation of Energy In a nonnuclear change, energy is conserved; it is neither created nor destroyed. Energy may change form, however. The energy lost in one form is always exactly equal to the energy gained in another form.

75. Conservation Laws Common Events in which Energy Changes from One Form to Another: