Why Study Chemistry?

1. Why Study Chemistry? • To be better informed • To be a knowledgeable consumer • To make better decisions for yourself and society • To learn problem-solving skills • To enhance analytical thinking

2. Chemistry as the “Central Science” • Chemistry = the study of matter and the transformation it undergoes • EVERYTHING is a CHEMICAL • Table salt = sodium chloride, NaCl • Table sugar = sucrose, C12H22O11 • Clothes: Wool? Cotton? Polyester? • Body: lipids, Proteins, Carbohydrates, DNA/RNA • You name it– it’s a chemical!

3. Chemistry as the “Central Science” • Chemistry is the driving force behind many “liberal arts” • Composition of paints? Colors? • Economies of industrial nations • #1 commercial chemical is sulfuric acid– LOTS of uses! • All idustry involves chemical processes • Economies of Developing Nations • Agriculture depends on chemicals as fertilizers, pesticides • Politics and Natural Resources

4. The Study of Chemistry • Chemistry is everywhere! • Matter is everywhere! • Thus, chemistry matters! • Chemistry involves the study of matter – its properties and behavior. • Macroscopic observations are rooted in microscopic structure.

5. Assignment: Chemistry in your major • Find a current news story or historical example that demonstrates the importance of chemistry to your major • For example: chemical resource as a key issue in a political / economic rift; wars fought over chemical resources; etc • Write a 2 paragraph summary on issue and its relevance to your studies

6. Chemistry as the “Central Science”

7. Classification of Matter 4 Physical States: solid, liquid, gas, plasma Solid: Fixed shape and fixed volume; Atoms tightly packed together

8. Classification of Matter Liquid: No fixed shape but maintains a fixed volume Atoms loosely packed together, slide around each other

9. Classification of Matter Gas: No fixed shape or volume Atoms not really associated with neighbors at all

10. Classification of Matter Plasma: mix of subatomic particles with not organization (sun)

11. States of Matter

12. States of Matter Gases, liquids and solids are all made up of microscopic particles, but the behaviors of these particles differ in the three phases. The following figure illustrates the microscopic differences.

13. States of Matter

14. Physical Properties = characteristics of a material Color Mass Temperature Odor Density Hardness Solubility Conductivity (heat or electrical) Freezing/boiling point Chemical Properties = describe how a material reacts with another type of matter Ability to burn Ability to rust / corrode Ability to make a solution acidic or basic Lack of ability to react with something Properties of Matter

15. Properties of Matter • physical – measured without changing substance, e.g. physical state, color, odor, density, boiling point • chemical – describes a substance’s reactivity, e.g. flammability, corrosiveness • extensive – depends on the amount of matter present, e.g. mass, volume • intensive – does not depend on the amount of matter present, e.g. density, color, temperature

16. Properties: “ The characteristics that give each substance its unique identity “ • Physical Properties: “ Properties that can be observed without changing the identity of a substance “ • Color • Melting Temperature - a physical change of state • Electrical conductivity • Density • Boiling Temperature - a physical change of state • Solubility • Hardness

17. Chemical Properties: “ Properties that result in changes in the identity of one or more reactants “ • The rusting of iron • Hydrogen and oxygen burning to form water • The baking of bread • The absorption of oxygen by hemoglobin

18. Physical Changes = a change in a physical property; does NOT change the chemical composition or atomic arrangement of the material Increase in temperature Phase changes Cutting into smaller pieces Chemical Changes = changes that alter the identity of a material, a change in the chemical composition or atomic arrangement of the material Wood burns in air to produce CO2 and H2O Cooking an egg (change molecular structure of the proteins, loss of water) Formation of rust (iron to iron oxide) Changes in Matter

19. Changes in Matter: Is it Physical or Chemical? • Properties: “ The characteristics that give each substance its unique identity “ • Physical Properties: “ Properties that can be observed without changing the identity of a substance • Color • Melting Temperature - a physical change of state • Electrical conductivity • Density • Boiling Temperature - a physical change of state • Solubility • Hardness continue…..

20. Changes in Matter (cont) • Chemical Properties: “ Properties that result in changes in the identity of one or more reactants “ • The rusting of iron • Hydrogen and oxygen burning to form water • The baking of bread • The absorption of oxygen by hemoglobin continue…..

21. Changes in Matter (cont) Reactants Products Hydrogen + Oxygen Water Chemical Reactions: “ Process in which one or more pure substances are converted to one or more different pure substances “ Reactants: “ Substances that undergo change in a chemical reaction “ • Reactants are on the left side of the chemical equation Products: “ Substances formed as the result of a chemical reaction “ • Products are on the right side of the chemical equation

22. Changes in Matter - Physical & Chemical • Physical Change: “ A change that alters the physical form ofmatter without changing its chemical identity “ • Chemical Change: “ A change which changes the chemical identity of the substance and creates one or more new substances “ continue…..

23. Changes in Matter - Physical Change • Example of a Physical Change: A Melting Ice Sickle Solid Water Liquid Water continue…..

24. 2H2O 2H2 + O2 Changes in Matter - Chemical Change • Example of a Chemical Change: The Electrolysis of Water(H2O) Particulate Viewpoint Oxygen Gas Hydrogen Gas Negative Electrode Positive Electrode continue….. The Chemical Identity of Water ( H2O ) is changed into the elements Hydrogen ( H2 ) and Oxygen ( O2 )

25. 2 C (s) + O2 (g) 2 CO 2.7 Using Chemical Symbols (cont) Chemical Equations: “ Representations of chemical reactions by the formulas of reactants and products “ At the Macroscopic Level: “ Carbon, a solid plus oxygen gas yields carbon monoxide “ At the Particulate Level: “ Two atoms of carbon plus one diatomic molecule of oxygen yields two molecules of carbon monoxide “ Equation Coefficients: “ Gives the relative amount of each compound involved in the chemical equation “ Balanced Chemical Equations: “ The number of each kind of atom on the reactant side must equal the number of each kind of atom on the product side “

26. Classification of Matter Matter - Anything that occupies space and has mass (solid, liquid or gas) Heterogeneous Mixture: Non-uniform composition Homogeneous Matter: Uniform composition Physically Separable Into Pure Substances: Fixed composition; cannot be further purified Solution: Homogeneous mixture Physically Separable Into Chemically decomposable Into Compounds: Elements united in fixed ratios Elements: Cannot be subdivided by chemical or physical changes Combine Chemically to

27. The Chemical View of Matter • What are elements and chemical compounds made of? • What is the difference between a mixture and a pure substance? • What is the difference between a chemical and a physical process? • What is the basic theme of chemistry? • How are symbols for the elements used in formulas and equations to communicate chemical information? continue….

28. Macroscopic, Microscopic & Particulate Matter • Matter: - “ Anything that has mass and takes up space • (occupies volume) “ • Matter can be studied on three levels: • Macroscopic Level: “ Matter that can be seen with the human eye “ • Beach Sand, Trees, Cars, Pen, CD, Mountains, • Planets, Galaxies, etc • Length: 101 to 109 meters continue…..

29. Macroscopic, Microscopic & Particulate Matter (cont) • Microscopic Level: “ Matter that is too small to be seen by the naked eye, but can be seen under a • microscope • Very small plants, individual bacteria, cellular • structures, DNA Molecule, Semiconductors, etc • Length: 10- 6 meters continue…..

30. Macroscopic, Microscopic & Particulate Matter (cont) • Particulate Level: “ Matter too small to be seen with even the most powerful optical microscope “ • Particulate matter consists of the tiny particles • that make up all matter • Molecules, atoms, protons & electron • Length: 10 - 10 meters (1 Angstrom = 10 - 10 meters ) continue…..

31. Elements - The Most Simple Kind of Matter Pure Substance: “Something that with a uniform, fixed composition at the submicroscopic level” • Recognized by the unchanging nature of their properties Element: “A pure substance composed of only one kind of atom” Atom: “The smallest particle of an element” • Atoms of different elements are different and are shown on the periodic table • Each element has a one or two letter abbreviation • Hydrogen - H • Helium - He • Sodium - Na • Lithium - Li

33. Elements

34. Main Group Elements Transition Metals Main Group Elements The Periodic Table and the Elements (cont) continue…. Inner Transition Elements

35. Chemical Compounds - Atoms in Combination Chemical Compounds: “ Pure substances made of atoms of different elements combined in definite ways” Examples: • H2O Water • NaCl Sodium Chloride • C2H6O Ethanol • C6H12O6 Sugar

36. Chemical Compounds (cont) • Compound: “ Any pure substance that can be decomposed by a • chemical change into two or more pure substances • is a compound “ - (another definition) • Compounds are made up of elements • Examples of Compounds: • Water - H2O Ethanol - C2H6O • Salt - NaCl Sugar - C6H12O6 • Examples of Mixtures of Compounds: • Pepper • Beer, Wine & Soda Pop • Milk • Cheese continue…..

37. Using Chemical Symbols Chemical Formulas: “ Combinations of the symbols for the elements that represent the stable combinations of atoms in molecules “ Examples: Water H2O Carbon dioxide CO2 Ammonia NH3 Methane CH4 Carbon Tetrachloride CCl4 Subscripts: “ Indicate the relative numbers of atoms of each kind “

38. H Ammonia H N H Water H O H H Methane H C H H Using Chemical Symbols (cont) Structural Formulas: “ Formulas that show the connections between atoms in molecules “

39. Microscopic view of the molecules of the compound water (gas phase). Oxygen atoms are red and hydrogen atoms are white.

40. Mixtures and Pure Substances • Homogeneous Sample: “ Matter that has a uniform appearance and • composition throughout “ • A mixture of water and alcohol • Sugar dissolved in water • Gold blended with silver (18 karat gold) • The air we breathe - a mixture of oxygen and nitrogen Solutions: “Homogeneous mixtures, either liquid, solid or gaseous” continue…..

41. Mixtures and Pure Substances (cont) Table salt is stirred into water (left), forming a homogeneous mixture called a solution (right) continue…..

42. Mixtures and Pure Substances (cont) • Heterogeneous Sample: “ Matter that does not have a uniform • appearance and composition • throughout “ • A mixture of cooking oil and water (two phases develop) • Concrete (sand, rock, cement, etc) • A mixture of sand, sawdust, iron fillings and water continue…..

43. Mixtures and Pure Substances (cont) Sand and water do not mix to form a uniform mixture continue…..

44. Homogeneous Same composition throughout sample Ex- milk, tea, others? Heterogeneous Different samples of the same mixture have different compositions Ex- air in the room others? Mixtures:

46. Microscopic view of a gaseous mixture containing two elements (argon and nitrogen) and a compound (water).

47. Classification of Matter

48. Substance has a definite or fixed composition Composition does not vary from sample to sample Mixture Has a varied composition Each individual component can be separated by physical means Ex: salt and pepper, sugar in water, sea water Substances vs Mixtures

49. Energy • The “fuel” of the universe • The capacity of something to do work • chemical, mechanical, thermal, electrical, radiant, sound, nuclear • The SI unit of energy is the Joule (J) • Other common units are • Calories (cal) • Kilowatt-hour (kW.hr) • Types of energy: • Potential • Kinetic • Heat • Energy cannot be created nor destroyed (but it does change from one type to another!)

50. Changes in Matter - Energy Energy: “ The ability to cause change or, in formal terms of physics, the ability to do work “ Potential Energy: “ Energy in storage “ • There is potential energy in gasoline called chemical energy • Chemical energy is release as heat and light when it burns • Chemical energy can also be released as electrical energy Kinetic Energy: “ Energy in motion “ • Examples are - Muscle in movement, a rocket in flight, • inflation of a car air bag during collision