1. Energy & Matter 2.1, 1.1, 1.2, 1.3 Element Song

2. 1. Energy (2.1) A.Energy: The capacity to do work or produce heat. 1.7 types of energy: mechanical thermal (heat) radiant (light) sound electrical chemical nuclear

3. 2. Kinetic Energy: Energy of motion. Ex. thermal, mechanical

4. 3. Potential Energy: stored energy; determined by position.

5. Ex. Electrical PE, chemical PE

6. Potential and Kinetic Energy

7. 4. Energy can be transferred from a system to its surroundings. Ex. Photosynthesis is light → chemical

8. Energy Transformations:

9. 5. Energy absorbing changes are called endothermic. If energy is released the change is called exothermic.

10. B. Measuring Energy: 1. Common Unit: calorie The amount of heat needed to raise 1 g of water 1 o C. (One calorie = 1g°C) 2. SI Unit for energy: Joule (J) C. Law of Conservation of Energy: Energy is neither created nor destroyed, it just changes form. Energy is neither created nor destroyed, it just changes form.

11. Temperature (2.1): 1. Energy can be transferred in the form of heat. heat. 2. Temperature is a measurement of heat or kinetic energy. (how fast the average kinetic energy. (how fast the average particle is moving!) particle is moving!) Heat vs. Temperature Animation Heat vs. Temperature Animation Kinetic Energy (Temperature) and Melting Kinetic Energy (Temperature) and Melting

12. 3. Common Temperatures Fahrenheit (  F) Celsius (°C)Kelvin (K) Background Information Popular (1686-1736) scientists (1701-1744) SI Unit Absolute scale (1824-1907) Boiling Point of Water 212100373 Body Temperature 98.637310 Room Temperature 7020293 Freezing Point of Water 320273 Absolute Zero -459.67-2730

15. Room Temp 20°C → Room Temp 70°F → ← Room Temp 293 K 98.6°F

16. 4. Kelvin:°C = K 5. The zero point on the Kelvin scale is called absolute zero (-273°C) 6. All motion of particles stops! Therefore the kinetic energy is zero. K = °C + 273 °C = K - 273

17. Q: Did you hear about the man who got cooled to absolute zero? A: He's 0K now.

18. 2. Matter is anything that has mass and takes up space. 1. Volume: Amount of space an object takes up. 2. Mass: Quantity of matter in a substance. Constant everywhere. Ex) the moon 3. Weight: Force produced by gravity acting on a mass. This is different in different locations.

19. Mass Vs. Weight Mass does not depend on gravity. The mass of an object remains the same in all locations. Weight depends on gravity. Weight equals Mass x gravity. The weight of an object changes with location. Weight and Mass Demo

20. B. Properties of Matter (1.2): 1. Physical: density, color, melting point, viscosity, surface tension, specific heat 2. Chemical: flammability, reactivity with other chemicals or air (O 2 ) C. States of Matter (1.1): StateShapeVolumeMovementStructure Solid definite vibrational - slowhighly organized - crystal Liquid indefinitedefinitetranslational - mediummedium - fluid Gas indefinite translational - fastlow - random Plasma is the 4 th state of matter “ionized gas” like the sun/fluorscent lights

21. D. Kinetic Theory of Matter (2.1) D. Kinetic Theory of Matter (2.1) Gases 1. Gases possess the greatest amount of kinetic energy. speed of the particles distance between them. 2. Two factors that determine the state of matter of a substance: speed of the particles and the distance between them. attraction 3. These two factors contribute to the attraction between the particles. change phase 4. Substances change phase when they overcome these attractions. kinetic energy 5. The overall kinetic energy (temperature) will remain constant until the entire substance has completely changed phase.

22. 6. Heating Curve for Water melting freezing 0 (◦C) 100 Vaporization ( boiling/evaporation) condensation Solid Liquid Vapor (gas) Cooling Curve Heating Curve

23. E. Changes in Matter (2.1): E. Changes in Matter (2.1): 1. Physical Changes: identity a. Do NOT change the identity of the substance. substance b. Often change what the substance looks like. cutting c. Examples:cuttingdyeing changes of state

24. States of Matter & Phase Changes melting freezing condensation vaporization evaporation –at the surface boiling - throughout deposition sublimation Solid Liquid Gas (Vapor) Gases are in the gaseous state at room temp. Vapors are in the solid or liquid state at room temp.

25. 3 States of Matter

26. States of Matter Comparison of the three states of matter

27. Ice Density =.92g/mL Density = 1.00g/mL @ 4 ◦ C Density =.998 g/mL @ 20ºC

29. 2. Chemical Changes: identity a. Alter the identity of the substance. different properties b. The new substance has different properties than the original substance. burning, rusting c. Examples of Chemical Changes: burning, rusting d. Signs that a chemical change has occurred: gas released (bubbles/odor/fizz/smoke) 1. gas released (bubbles/odor/fizz/smoke) color change (can be physical too) 2. color change (can be physical too) formation of a precipitate (insoluble solid that 3. formation of a precipitate (insoluble solid that falls out of solution.) falls out of solution.) temperature change  (can be physical also) 4. temperature change  (can be physical also)

31. F.Law of Conservation of Matter (2.2): Matter is neither created or destroyed it just changes form. G. Classification of Matter (1.3) unique 1. Pure substances: Substances that have a unique set of physicalchemical physical and chemical properties. a. Elements: The smallest part of an element is atom. an atom. simpler substances. 1. Cannot be separated into simpler substances. symbols 1 2 2. Represented by symbols that have 1 or 2 letters. Ex) K, Na, Au, Ag, Hg, Fe, Co (three lettered symbols are temporary) (three lettered symbols are temporary)

32. 3. Examples: 1 H Hydrogen 1.008 Atomic Number: # of protons Element Symbol: 1 or 2 letters (1 st is a capital) Element Name Atomic Mass: (weighted average of all an elements’s isotopes) b. Compounds: 2 1. Made up of 2 or more kinds of atoms chemically chemically combined in a fixed proportion. formulas. 2. Represented by formulas. CO, CO 2, H 2 O, NH 3 3. Examples: CO, CO 2, H 2 O, NH 3

33. 2. Mixtures: different a. Heterogeneous Mixture: Visibly different throughout. Will separate upon standing. salad dressing (emulsion), chocolate Ex) salad dressing (emulsion), chocolate chip cookies, sand & water (suspension) chip cookies, sand & water (suspension) same b. Homogeneous Mixture: The same throughout. May be clear, will not separate. Kool-aid (solution) Ex) Kool-aid (solution) milk (colloid) milk (colloid) gold jewelry (alloy)

34. Examples of Alloys Brass is an alloy of copper and zinc. Steel is an alloy of carbon and iron. Bronze is an alloy of copper and tin.

35. Gold – Element & Alloys

36. Microscopic look at mixtures

37. SuspensionsColloidsSolutionsAlloys ex) sand & water ex) milk ex) Kool-Aid ex) gold jewelry

39. H. Separating Mixtures (1.3) 1. Heterogeneous Mixtures can be separated by: residue filtrate Ex) sand & water a. Filtration- Material remaining on the filter paper is called the residue. The filtrate goes through the filter paper. Ex) sand & water

41. Separation of Homogeneous Mixtures: liquids boiling point residue distillate a. Distillation- separates liquids (and 1 solid) by differences in boiling point. The remaining material is called the residue. The material that goes through is called the distillate. alcohol & H 2 O Ex) alcohol & H 2 O

42. Another Look at Distillation A Closer Look at Distillation

43. Separation of Homogeneous Mixtures Evaporate crystallizesalt and water b. Crystallization- Evaporate liquid and the solid will crystallize. Ex) salt and water

44. separate solubility (density) black ink - rainbow c. Chromatography – used to separate pigments and ink by differences in solubility (density) on a strip of paper. Ex) black ink - rainbow

45. Another look at Paper Chromatography

46. 3. Separating Compounds: elementswater into H 2 and O 2 a. Electrolysis – decomposes a compound into its elements. Ex) water into H 2 and O 2