Unit 1 notes_pp_12_13
- 1. Section 1- Matter and Change Unit 1 – Matter and Energy
- 2. What is chemistry? Chemistry: the study of the composition, structure, and properties of matter, the process that matter undergoes, and the energy changes that accompany these processes. Chemicals: any substance with a definite composition. Examples: H2O and H2O2
- 3. Matter: anything with mass and volume. Weight: Mass: Force of Gravity on an Object (Measured with a scale) Amount of Matter in an Object (Measured with a balance)
- 4. Matter and Its Properties Atom: Smallest unit of an element… Element: A pure substance that CANNOT be broken down into a simpler substance (one type of atom) Example: Cu or He Compound: A substance that can be broken down into a simpler substance (2 or more elements chemically bonded together) Example: H2O or H2SO4
- 5. Physical Property Measured without changing the substance Intensive DOES NOT depend on the amount Extensive DEPENDS on the amount Melting Point, Density, Freezing point Volume, Mass, Energy
- 6. Chemical Property Behavior of a substance interacting with another Coal Burning Iron Rusting
- 7. Matter and Its Properties Physical Change: Chemical Change: Same substance (chemically) present BEFORE and AFTER the change Produces ONE or MORE new substances.
- 8. Matter and Its Properties Indicators of Chemical Change 1. Precipitate (solid formed) 2. Gas is given off 3. Color change 4. Change in energy (Hot OR Cold) 5. Light given off (emission)
- 9. Classification of Matter Homogeneous Materials Elements Is it the same throughout? Is it a combination of things? Can it be separated in to smaller substances?
- 10. Groups = COLUMNS Periods = ROWS
- 11. Unit 1: Matter & Energy Section 2: Measurements and Calculations
- 12. Scientific Method Theory Publish Results
- 13. Units of Measure Qualitative Descriptive (Non-numeric) The fact that the sky is blue Quantitative Measurement (Numeric) A sample of copper ore has a mass of 25.7 grams
- 14. Units of Measure The International System (SI)--1960s Commonly used SI units in CHEMISTRY. 1. Length meters (m) 2. Mass kilograms (kg) 3. Time seconds (s) 4. Quantity moles (mol)
- 15. Units of Measure SI units PREFIXES Prefi x Symbo l How many in a base unit? Nano- n 1,000,000,000 Micro- μ 1,000,000 Milli- m 1,000 Centi- c 100 Deci- d 10 Prefix Symbo l How many base units? Kilo- k 1,000 Mega- M 1,000,000 Giga- G 1,000,000,000 Example: 1 m = 100 cm 1 g = 1,000 mg Example: 1 km = 1000 m 1 mg = 1,000,000 g
- 16. Units of Measure Derived Units: A combination of SI base units. Examples: Velocity, Density Conversion Factors Math used to relate 2 units that measure the same quantity (written as a fraction); Equal to 1 Example 1 m = 1000 mm
- 17. Units of Measure Conversion Factor Practice Example 1 – Convert 22,000 g to kg. Example 2 – Convert 0.0290 m to mm. Example 3 – How many seconds are in 3.11 hours?
- 18. Units of Measure Example 1 – Convert 22 000 g to kg 22000 g x 1 kg = 22000 kg = 1 1000 g 1000 Example 2 – Convert 0.0290 m to millimeters 0.0290 m x 1000 mm = 29 mm = 1 1 m 1 Example 3 – How many seconds in 3.11 hours? 3.11 hours x 60 min x 60 sec = 11196 sec = 1 1 hr 1 min 1 22 kg 29 mm 11196 sec
- 19. Using Scientific Measurement Accuracy Precision Closeness of a measurement to the true or CORRECT value. How close a set of measurements are to one another, regardless of correctness
- 20. Using Scientific Measurement Experiments will always have errors. (human, mechanical, environmental) PERCENT ERROR determines the accuracy of the experiment. (the lower the percent, the better)
- 21. Using Scientific Measurement Significant Figures All digits that occupy places for which ACTUAL measurement was made, including the last estimated digit. See Significant Figure/Scientific Notation Rules Reference Sheet When taking measurements, you ALWAYS read the instrument to one more place than is marked. You estimate that digit.
- 22. Significant Figures Practice 1) 30 504 5 sig. figs. 2) 32.001 20 7 sig. figs. 3) 0.000 123 0 4 sig figs. 4) 560 000 2 sig. figs. 5) 200.003 6 sig. figs.
- 23. Math w/ Significant Figures Put answers in correct Sig. Figs. 1) 2.301 + 6.12 + 4.1158 12.5368 = 12.54 2) 1500 – 301.05 – 251.223 947.727 = 948 3) (9.554)(5.10)(1.5) 73.0881 = 73 4) 44581.2/235.2 189.545918367 = 189.5
- 24. Using Scientific Measurement Scientific Notation How scientists show either BIG or SMALL numbers. 602,200,000,000,000,000,000,000 = 6.022 x 1023 Hint: Positive exponents – Move decimal to the right Negative exponents – Move decimal to the left. Note: Only the numbers before the scientific notation are significant. (ex: 6.022 x 1023 has 4 sig figs)
- 25. Scientific Notation 1) Put 325 000 000 in scientific notation 3.25 x 108 2) Put 0.002405 in scientific notation 2.405 x 10-3 3) Put 7.54 x 104 in regular notation 75 400 4) Put 5.001 x 10-5 in regular notation 0.00005001
- 26. Using Scientific Measurement Direct proportion Inverse Proportion As one variable increase, so does the other. As one variable increases, the other decreases
- 27. Unit 1: Matter & Energy Section 3: Energy Transfer
- 28. What is Energy? Energy Chemical Energy due to chemical bonds Kinetic Energy due to motion Heat (q) is the total KE in a sample Potential Energy due to position
- 29. Energy Units Name Symbol Info Joule J SI Unit for energy calorie c (lowercase) Old energy unit 1 calorie = 4.184 J Calorie C 1000 c = 1 Calorie Food Calorie
- 30. Energy Transfer Energy in the form of HEAT moves back and forth between a system and its surroundings. Temperature: Average KE of the random motion of particles in a substance.
- 31. Energy Transfer Specific Heat Capacity Amount of HEAT needed to raise the temperature of 1 gram of a substance 1 Kelvin (K). Calorimeter A device used to measure energy absorbed or released during chemical or physical change.
- 32. Energy Transfer How to calculate specific heat capacity: q = (m)(ΔT)(Cp) q = Heat (J) m= Mass (g) ΔT = Change in Temperature (K = °C + 273) Cp = Specific Heat Capacity (Given)
- 33. Example Problem #1 40.00 g of water is heated from 10.00°C to 30.00°C. The specific heat of water is 4.184 J/gK. q = (m)(ΔT)(Cp) q=Heat (J) m=Mass (g) ΔT = Change in Temperature (K = °C + 273) Cp = Specific Heat Capacity (Given) q = ? m = 40.00 g ΔT = (30 + 273) - (10 + 273) = 20.00 K Cp = 4.184 J/gK q = (40.00 g) ( 20.00 K) ( 4.184 J/gK) q = 3347.2 J q = 3350 J
- 34. Example Problem #2 135.6 g of water is cooled from 95.80°C to 21.60°C. The specific heat of water is 4.184 J/gK. q = (m)(ΔT)(Cp) q=Heat (J) m=Mass (g) ΔT = Change in Temperature (K = °C + 273) Cp = Specific Heat Capacity (Given) q = ? m = 135.6 g ΔT = (21.6 + 273) - (95.8 + 273) = -74.20 K Cp = 4.184 J/gK q = (135.6 g) ( -74.20 K) ( 4.184 J/gK) q = -42097.39 J q = -4.210 x 104 J
- 35. Energy Transfer Practice Determine the amount of heat absorbed or released in each of the following changes. Use the table of specific heat capacity on page 533 of your text. The specific heat of water is 4.184 J/gK. 1. 40.00 g of water is heated from 10.00°C to 30.00°C. 2. 135.6 g of water is cooled from 95.80°C to 21.60°C. 3. 30.00 g of aluminum is heated from 15.00°C to 35.00°C. 4. 450.0 g of iron is cooled from 125.0°C to 45.00°C. 5. 62.30 g of lead is heated from 21.70°C to 136.4°C.
- 36. Energy Transfer Answers to previous slide 1. +3347.2 J = + 3347 J 2. -42,097.4 J = -4.210 x 104 J 3. +538.2 J 4. -16,164.0 J = -16160 J 5. +921.8 J