Introduction to Chemistry

1. Introduction to Chemistry Unit 1

2. Lab Safety • Heat and Flames: • Flammable material should never be poured near a flame. • If a fire does erupt – I will deal with it. Evacuate. • When heating a test tube – always point it away from others • Hot and cold glass looks the same – use tongs! • Never leave a lit burner unattended • Eyewear: • Must be worn at all times chemicals, flames, or glassware is used. • Full coverage of eyes is required • I need to know if you wear contacts. There may be times that you will not be allowed to wear them.

3. Lab Safety Continued • Lab Procedure • Follow all directions – get help if you don’t understand something • Make sure glassware is clean and not broken • No horseplay, throwing things, etc. – ever! • Do not touch lab equipment until the lab time begins. • Treat all chemicals as dangerous even if they’re not – be oversafe (never taste, dispose of as directed, etc.) • Add acids to water and not water to acids. • Apparel/Food/Drinks/Etc. • Long hair needs to be pulled back No dangling jewelry • Closed-toe shoes required • No gum, food or drinks (even water) near the lab areas. • Accidents • Emergencies I will handle, evacuate • Small cut: tell me and first aid will be applied

4. Chemistry • Chemistry: The study of the composition, structure, and properties of matter and the changes it undergoes. • Chemical: Any substance that has definite composition.

5. Matter • Review (Discuss with your partner – define mass, volume and matter) • Mass= a measure of the amount of matter. • Volume = the amount of 3-d space an object occupies • Matter= Anything that takes up space and has mass • Matter cannot be created nor destroyed – only changed • Examples that ARE Matter: desk, air, you • Examples that ARE NOT Matter: energy, light, heat • Two types of matter = pure substances and mixtures • Pogil

6. Pure Substances • Pure Substancescannot be broken down physically, only chemically • Element = a pure substance that cannot be broken down into simpler, • stable substances and is made of one type of atom. • Atom = smallest particles of an element that is still that element • Elements can be metals, nonmetals or metalloids • Compound = a substance that can be broken down into simpler substances. Each is made from atoms of two or more elements that are chemically bonded. • Molecule– Particle made of 2 or more atoms • Smallest particles of a molecular compound that this still the compound • Some elements also exist at molecules = diatomic elements (H2,O2, etc.) • Many compounds are made of molecules, but others are made of ions, charged particles that form crystal lattice structures instead • Water = molecular, salt = ionic

7. Pure Substances • Identifying and Decomposing Pure Substances • Apure substance has a fixed composition and differs from a mixture in the following ways: • Every sample has the same characteristics, exactly • Compare density with known, accepted densities • Boiling points for pure substances are fixed • Freezing and melting points are also fixed – pure substances have higher melting points than impure • How to decompose or break down pure substances?? • Electrolysis – uses electricity to break down substances • Separates differently charged particles such as Na and Cl in salt • Chemical Change

8. Mixtures • Mixtures (Homogeneous and Heterogeneous) can be broken down or separated by physical means • Homogeneous = same/uniform throughout; all samples of the substance are the same. • Lemonade (solid/liquid), soda pop (gas/liquid), fog (liquid/gas) • Alloys (Steel, Bronze, solid/solid), air (gas/gas), vinegar (liquid/liquid) • Heterogeneous = non-uniform throughout; all samples of the substance are different • Chocolate chip ice cream, chex mix • How can you separate mixtures? • Boil, Distillation (boiling, but recovers water) • Settle, Filtration • Others??? (spoon, magnet, etc.)

9. Classification of Properties • Properties of Matter can be classified as either extensive or intensive • Extensive = depends on the amount of matter that is present (volume, mass, etc.) • Intensive = do not depend on the amount of matter present (melting point, boiling point, density, ability to conduct electricity) • Matter can also be described using both chemical and physical properties. • Physical properties = characteristics that can be observed or measured without changing the identity of the substance. (color, odor, melting and boiling point,etc.) • Chemical properties = characteristics that relate to a substance’s ability to undergo changes into a different substance

10. Chemical and Physical Changes • Baggie Lab • Physical Changes – Changes in state, same chemical properties (does not involve a change in the identity of the substance) • Change in physical properties • Examples: Melting/Freezing/Grinding/Cutting/Boiling • Change of state = a physical change of a substance from one state to another. • Same smells, tastes, etc.

11. Physical and Chemical Changes • Chemical Changes – Changes in the properties of substance • Book definition = A change in which one or more substances are converted into different substances. • Chemical equations are used to show the substances that change (the reactants, left side) produce new type of matter (the products, right side) • Arrows are used in reaction to show yields, plus signs mean the substances are added together • Melting sugar – changes the type of matter • Energy (often heat) causes both physical and chemical changes • The energy may be heat or light • Typically more energy = chemical change at some point, but definitely not always. • 4 Things always identify a chemical change – unexpected color change, temperature change, formation of a gas/bubbles, formation of a solid/precipitate

12. States of Matter • Partner Activity: Draw a beaker with a solid, a liquid and a gas; describe each beaker. • Solids • Packed tightly together • Particles vibrate • Definite shape and volume • Liquid • Fairly close together • Moving slowly, can still move/slide past each other • Defined volume, but not shape • Gases • Far apart • Particles move quickly • No defined shape or volume

13. Kinetic Molecular Theory of Gases • Read page 329 in book and complete activity in packet • Only applies to ideal gases, non-polar gases at high temperate and low pressure. • As temperature drops or pressure increases, properties change. • Complete Chart on Page 10 – look up definitions you do not know.

14. Phase Changes • Solid  Liquid = melting • Liquid  Gas = evaporation and boiling • Gas  Liquid = condensation • Liquid  Solid = freezing • Solid  Gas = sublimation • Gas  Solid = deposition • Vaporization • Stronger forces between liquid particles = lower equilibrium vapor pressure (less particles can evaporate) • Volatile liquids = evaporate readily, weak attractive forces (rubbing alcohol – compare to water)

15. Boiling, Freezing and Melting • Boiling • Boiling = conversion of liquid to a vapor within liquid and at the surface • At the boiling point, the energy absorbed is used to evaporate and thus temp change DOES NOT occur as long as pressure remains constant • Energy must be constantly added to maintain boiling • Energy overcomes attractive forces between molecules of the liquid • Freezing and Melting • Freezing = change from liquid to solid, involves loss of energy in the form of heat • Particle order increases as it becomes solid • Melting and freezing occur at constant temp • Only after all of the ice is done melting/freezing does the temperature change

16. Phase Diagrams • A diagram illustrating the conditions under which the phases of a substance exist • Each curve represents areas (temp and pressure conditions) when the two phases on either side of the curve can coexist at equilibrium • 3 important points: • triple point = the temp and pressure at which all three phases can coexist at equilibrium • critical temp = the temp above which the substance cannot exist as a liquid • critical pressure = lowest pressure at which the substance can exist as a liquid at the critical temp. • Diagram also indicates normal freezing and boiling points

17. Phase Diagrams • triple points = A and X • critical temp = 374 ○ C and 31.1 ○ C • critical pressure = 218 atm and 73 atm