The top of a small table is 0.500 m 2 .

1. The top of a small table is 0.500 m2. • Calculate how many square inches it is (1.000 inch =2.540 cm) • Is the table top a pure substance? • If the table is sawed in half, is that a physical or chemical change? • If the table is burned, is that a physical or chemical change? • If the table floats in water (before being burned), what does that tell you about the spacing of the molecules compared to water?

2. Is ice water homogeneous or heterogeneous? • Is freezing a chemical or physical change? • Is baking cookies a chemical or physical change? • How do molecules differ in a liquid than a solid? • A soda can has a radius of 3.25 cm and a height of 12.00 cm. • Calculate the volume of the can • The can has a mass of 557 g. Calculate the density. • Convert the density to kg/m3.

3. Pure Substances • Pure Substances – either all one element or all one compound. Have definite unchanging properties • Types a) Elements –Au, O2 b) Compounds – pure water

5. Homogeneous Mixtures • Uniform mixture of two or more elements and compounds (only one phase) • Types a) Solutions – liquid homogeneous mixtures Examples Kool-Aid Salt Water

6. Homogeneous Mixtures b) Alloys – Solid uniform mixtures. Usually metals. Stainless Steel – Iron and Cr “Gold” Ring – Au, Cu & Zn

7. Heterogeneous Mixtures • Nonuniform mixture, composed of two or more phases • Phase – One physical state with definite boundaries • Examples: Ice Water – Two Phases (water & ice) Italian dressing – More than two phases Granite – Multiple phases Zinc and Sulfur

9. Physical Properties • Can be observed without changing the substance into another substance • Examples: Melting Point/Freezing Point Boiling Point Hardness Solubility Malleability Ductility

12. Solids Can “wiggle” in place (these are the wiggle lines)

14. Liquids They wander in random patterns quite close to one another.

15. Gases

16. Plasma • Plasma - 4th state of matter • Ionized gases • Electrons removed from the atoms • Positive ions remain • Present in: • Stars • Lightning • Arc welding • Most common state of matter in the universe

17. Plasma hydrogen and helium plasma (sun) e He+ e He+ e H+ H+ e He+ e He+ e e H+ H+ e

19. Can an elephant swim?

20. Can an elephant swim?

21. Physical Properties : Density • History – Archimedes story • Density = mass per unit volume • D = m/V Mass  Grams V  mL or (cm3) D  g/mL

22. Physical Properties : Density • Intensive property – does not depend on amount present • Volume formulas Rectangular Cylinder Irregular shape (Archimedes)

23. Physical Properties : Density 6. Solving for variables a. Algebra 3 = x/8 2 = 6/x b. Density Solve for m Solve for V

24. Physical Properties : Density • A ring has a mass of 8.99 g and a volume of 0.796 mL. Is the ring gold (19.3 g/mL)? (Ans: 11.3 g/mL) • A substance masses 47.5 g. It is put into a grad cylinder containing 12.5 mL of water. After immersing the substance, the total volume is 31.8 mL. What is the density? (Ans: 2.46 g/mL)

25. Physical Properties : Density • Salt has a density of 2.16 g/mL. What is the volume of 485 g of salt? (Ans: 225 mL) • What is the mass of 1.520 L of kerosene (r = 0.8200 g/mL)? (Ans: 1246 g) • What volume would 0.450 kg of lead occupy (r = 11.3 g/cm3) (Ans: 39.8 cm3) • A cylinder has a radius of 2.00 cm and a height of 10.0 cm. If the cylinder is filled with 0.289 grams of a gas, density? (Ans: 0.00230 g/cm3)

26. A rectangular sample of aluminum has the measurements 1.34 cm by 2.58 cm by 10.00 cm. • Calculate the volume of the sample. (34.6 cm3) • Calculate the mass of the sample. The density of aluminum is 2.70 g/cm3. (93.3 g) • Would the block of aluminum float in mercury? (13.6 g/mL) • An irregular sample of aluminum is found to have a mass of 50.00 grams. The sample is placed in 10.00 mL of water. Calculate the final volume reading of the graduated cylinder. (28.5 mL)

27. Chemical Properties • Chemical Change – one substance changed into another substance • Chemical Properties – tendency of a substance to react with other substances • Examples Flammability Will it rust (oxidize)? Acid or base?

28. Chemical Properties • Atoms rearrange and bonds are made and broken: 2H2 + O2 2H2 O

29. Chemical Properties Recognizing a chemical change: • Properties change (rust is different than iron) • Gets hotter or colder • Exothermic – Gives off heat (burning gas, exercise) • Endothermic – Absorbs heat (cooking) • Color change • Gas given off • Light produced (glow stick)

30. Law of Conservation of Mass • Antoine Lavosier (1789) • Law of cons. of mass – mass is neither created nor destroyed in a chemical reaction

32. Law of Conservation of Mass • Two Examples 1. 2H2 + O2 2H2O 10g + 80g  90g 2. Burning wood ? Wood + O2 Gas + Ash.

33. Law of Conservation of Energy Energy can never be created or destroyed. It can only change form. • Battery: turns chemical to electrical to mechanical energy. • ALWAYS lose some energy in transformations (usually as waste heat)

34. Einstein:1905 Law of cons. of mass/energy – mass and energy cannot be created or destroyed. They only change form E = mc2 E= Energy m = mass c = speed of light

35. Einstein:1905 2. A small amount of matter can be destroyed to release a large amount of energy (nuclear processes) (20 g U  18 g U)

36. Energy – The capacity to do work • 1 Joule = 1 Newton-meter • 1 calorie = amount of heat to raise one gram of water by 1o C 1 calorie = 4.18 Joule (1 nutritional Calorie = 1000 calories, 1 kilocalorie)

37. Two types • Potential Energy - Stored Chemical Energy – energy stored in chemical bonds • Plants absorb energy from the sun • That energy is released through digestion/burning • Kinetic Energy – energy in motion Energy of moving atoms and molecules

38. Specific Heat Specific Heat – Amount of heat needed to raise the temperature of one gram of a substance by one degree Celsius or Kelvin • Unit – cal/goC or J/goC • Symbol = Cp • Higher the specific heat, the more energy needed to raise the temperature

39. Predict the specific heat of the following (high or low): Car hood Pot of water Beach sand Plants & Trees

41. Calculating Heat q = mCpDT q = heat in Joules m = mass (grams) Cp = specific heat (J/goC) DT = Tfinal – Tinitial

42. How much heat must be supplied to a 500.0 gram iron pan (Cp = 0.444 J/g oC) to raise its temperature from 20.0oC to 100.0oC? q = mCpDT q = (500.0g)(0.444J/g oC)(100.0oC -20.0oC) q = (500.0g)(0.444J/g oC)(80.0oC) q = 17,800J or 17.8 kJ

43. Suppose we use a similar pan, except it is now made of Aluminum (Cp = 0.895 J/goC)? q = mCpDT q = (500.0g)(0.895 J/goC)(100oC -20oC) q = (500.0g)(0.895 J/goC)(80oC) q = 35,800 J or 35.8 kJ

44. What temperature change would 50.0 g of rock undergo if they absorbed 452 Joules of heat? (Assume the specific heat is 0.836 J/goC).) ANS: 10.8 oC (which is also 10.8 K)

45. A sample of copper (Cp = 0.092 cal/goC) undergoes a temperature change from 24.0oC to 76.0 oC upon the addition of 1300.0 Joules of heat. • Convert the heat to calories • Calculate the mass of the copper (65 g)

46. A 500.0 g sample of zinc absorbed 4850 J of heat. The temperature increased from 20.00 oC to 45.00 oC. • Calculate the specific heat of the sample in J/goC. (0.388 J/goC) • Calculate how many calories of heat the sample absorbed. (1160 cal) • A separate sample of zinc absorbs the same amount of heat. However, the temperature only rises from 20.00oC to 28.00oC. Calculate the mass of the sample in kilograms. (1.56 kg)

47. Heating Curves • Changes of state do not have a temperature change. • Melting/Freezing • Boiling/Condensing • A glass of soda with ice will stay at 0oC until all of the ice melts. • Graph “flattens out” during changes of state

48. Steam heats up Temperature (oC) Boiling Water warms up Melting Ice warms up Heat (Joules)