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Chemistry: The Study of Matter Chapter 1 http://xkcd.com/435/ Worldviews The overall perspective with which one sees and interprets the world. Naturalistic Worldview Matter is everything and science is the only path to “truth”. Christian Worldview Science is the discovery of God's Handiwork in creating matter and all the universe “They exchanged the truth about God for a lie, and worshiped and served created things rather than the Creator-who is forever praised.” Romans 1:25 Old Testament Chemistry Genesis 4:22 Metallurgy – Extracted pure metals from ores (raw earth material) via smelting (heat decomposition). Created useful alloys (purposeful mixing of metals for desirable properties) Bronze alloy – Copper and Tin Steel alloy – Iron and Carbon Exodus 30:25 Apothecary - “The original pharmacists” Used chemicals and herbs for medicinal purposes Greek Chemistry ~ 430 BC Democritus' theory – Philosophical atomism (no evidence) All matter is made up of tiny identical atoms and the difference in materials is based on the shape, position, and arrangement of these atoms. "Atomos" - indivisible Alchemists The "original chemists” Attempted to make gold from other substances. Impossible challenge (without nuclear reactions) Nevertheless, resulted in organized approach to science Laboratory techniques Equipment Terminology Why Study Chemistry? Creation Mandate Gen. 1:26, 28 God blessed them, and God said to them, "Be fruitful, multiply, fill the earth, and subdue it. Rule the fish of the sea, the birds of the sky, and every creature that crawls on the earth." (Genesis 1:28 HCSB) Career Foundation Pharmacy Medical Engineering Dietician Agriculture Environmental Material science • Critical thinking Skills • Problem solving • Deductive logic • Scientific inquiry Chemistry: A Science for the st 21 Century • Materials and Technology • Plastics, ceramics, liquid crystals • Room-temperature superconductors? • Molecular computing? Binary data stored in DNA • Food and Agriculture • Genetically modified crops • “Natural” pesticides • Specialized fertilizers GFP: Green fluorescent protein Fields of Chemistry • Organic – carbon-containing compounds (>70% of substances, plastics, drugs) • Inorganic - all elements minus carbon (metals and coordinating elements) Fields of Chemistry • Biochemistry – organic chemical processes in living things (Biomolecules: proteins, DNA, lipids, carbohydrates) • Analytical – Create/improve chemical techniques used in all branches for precise quantitative measurements. (Purification, sample analysis, water/soil testing) • Physical - foundational theories, detailed study of interaction and energy changes (e- probability, thermodynamics, quantum) The Study of Chemistry: We observe the Macroscopic Macroscopic Microscopic Chemistry explains what’s happening on the Microscopic scale 1886 2-6 year Process Today 2 Cu + H2O + CO2 + O2 → Cu(OH)2 (s) + CuCO3 (s) Oxidized mixture called “Patina” Making Observations Qualitative observations Describes the quality of an object Color, taste, texture, appearance, smell, etc. Think Adjectives Quantitative observations Describes an object using numbers Count, length, weight, volume Think Units List 2 observations of each type you could make The scientific method is a systematic approach to researching phenomena, acquiring new knowledge, or correcting and integrating previous knowledge. Macroscopic Microscopic/Symbolic Explain Observations A hypothesis is a tentative explanation for a set of observations. tested modified The scientific method is a systematic approach to research. Hypothetical Method Actual Method * http://www.wired.com/wiredscience/2013/04/whats-wrong-with-the-scientific-method/ A theory is a unifying principle that attempts to explain a body of experimental observations. Theories offer explanations for what we observe. • Atomic Theory • Cell Theory Theories tell us why we should expect it. • Big bang theory Do not confuse scientific theories as improbable explanations filled with inconsistency. They are often incapable of absolute proof, but all available data are still in support of them. A law is a concise statement that is always the same under the same conditions. Laws describe observations Often mathematical equations Laws tell us what we should expect (∝ = directly proportional) Newton's 2nd Law: Force = mass x acceleration 2nd law of thermodynamics: Entropy > 0 Charles’s Law: V ∝ T Chemistry is the study of matter and the changes it undergoes. Matter is anything that occupies space and has mass. Substances are pure forms of matter that have definite composition and distinct properties. liquid nitrogen gold ingots Talc (mineral) A mixture is a combination of two or more substances in which the substances retain their distinct identities. 1. Homogenous mixture – composition of the mixture is the same throughout Solutions (soft drink), gas mixtures (air), solder (Sb/Pb alloy) 2. Heterogeneous mixture – composition is not uniform throughout cement; oil and water; iron filings in sand; insoluble compounds Substance or Heterogeneous/Homogenous mixture? • Kool-Aid • Distilled water • Skittle/M&M’s • Bronze statue • Copper pipe • Vinaigrette dressing Mixtures can be separated into their pure components by some physical means. Distillation – Separating liquid mixtures by their boiling points Filtration– Separating mixtures by their phase Separating Sand/Iron via a magnet Pure Mixture Physical Properties: can be measured or observed without changing the composition or identity of a substance. • Density: amount of mass per volume of space • Malleability: Hammered into a thin sheet • Ductility: Drawn into long thin strings • Conductivity: Ability to transfer either heat and/or electricity • Phase transition temperatures: temp. melting/boiling occurs • Appearance: color, luster, texture • Solubility: amount dissolvable in solvent (water) • Hardness: measured by Mohs scale (1: Talc - 10: diamond) Extensive and Intensive Properties of matter An extensive property depends upon how much matter is being considered. • mass • length • volume An intensive property of a material does not depend upon how much matter is being considered. • Density • Temperature • Color •Viscosity Types of Changes A physical change does not alter the composition or identity of a substance. ice melting sugar dissolving in water A chemical change alters the composition or identity of the substance(s) involved. Metal rusting Hydrogen burns in air to form water A chemical change may result in one or more of the following: • Solid production from solution (precipitation) • Color change • Odor originates • Gas production (effervescence) • Temperature change (potential flame) Physical or Chemical Change? • Grinding coffee beans • Food rotting • Lighting a match • Cutting paper in half • Water boiling to steam • Jewelry tarnishing • Dissolving orange Kool-Aid in water An element is a substance that cannot be separated into simpler substances by chemical means. • 118 elements have been identified • 98 elements occur naturally (some only in trace amounts) Mercury Aluminum Sulfur Carbon • 20 elements have been synthetically created by scientists Plutonium Americium Atoms: the basic particles that make up the different elements • Ex. Li, Be, B, , F, Ne, Au • Either 1 or 2 letter symbol; first letter capitalized Atoms possess subatomic particles: Neutrons (N0) - no charge, but have mass Protons (P+) - positively charged and have mass Electrons (e-) - negatively charged, but little mass When an atom has equal Protons and Electrons it is Neutral ex. a neutral Helium atom contains 2 P+ and 2 e- Ion: When P+ and e- are unbalanced in an atom, it is Charged. ex. an ionized Sodium ion (Na+1) has 11 P+ and 10 eTedEd: Just how small is an atom? https://www.youtube.com/watch?v=yQP4UJhNn0I Elemental symbols Si ≠ SI Aurum Kalium Ferrum Plumbum Argentum Natrium * * *Many are derived from their Latin names *Hydrargyros "water silver" *Wolframite: W containing ore A compound is a substance composed of multiple elements chemically united (bonded). Compounds can only be separated (broken down) into their pure components (elements) by chemical means. Lithium fluoride: LiF Quartz: SiO4 dry ice – CO2 (carbon dioxide) *Non-compounds are not necessarily always monoatomic (C, He): Can have many element atoms in a substance: P4, S8, Cl2 Classifications of Matter ex. Carbonated Water ex. Iron Ore (compound mixture) Fe3O4 Fe2O3 FeCO3 H2O + CO2 C + O2+ H2 Minerals Fe3O4 Fe2O3 FeCO3 Fe + O2+ C Review of the Nucleus: The Nucleus: Crash Course Chemistry #1 https://www.youtube.com/watch?v=FSyAehMdpyI Lab Glassware Non-Quantitative Borosilicate Glass (SiO2 + B2O3) • Withstands higher temperatures Erlenmeyer Flask Beaker • Lower thermal expansion (hot to cold) • Less likely to shatter Used Quantitatively • Used to contain chemicals/reactions • Used to heat liquids • Not used to heat solids Crucible used for prolonged heating of solids Buret Graduated Cylinder Volumetric Flask Lab Warnings Toxic: poisonous to living organisms (dose dependent) • Acute (rapid onset) • Chronic (slow-development) Oxidizer: Electron thief, very reactive (chemical burn) Corrosive: Damage/ destroys on contact (chemical burn) Carcinogenic: Causing cancer via altering genome (DNA) Flammable: Easily burns/ignites in contact with heat/spark "The dose makes the poison" The Three States of Matter: Effect of a Hot Poker on a Block of Ice gas liquid solid A Comparison: The Three States of Matter Defined shape, incompressible Undefined shape, incompressible Undefined shape, Compressible A Comparison: The Three States of Matter Kinetic Molecular Theory: describes motion of particles in various states of matter Solids: Minimal particle motion locked in place; only vibrations Liquids: Greater freedom of motion particles shift/slide Gases: Random, fast movement of particles (non-interacting) SciShow: How to supercool water; https://www.youtube.com/watch?v=NMSxuORKynI Temperature is a measurement of the movement of particles. °F = 9 x °C + 32 Absolute Scale Relative Scales 5 K = 0C + 273.15 0 K = -273.15 0C 0 K = -460 ° F “Water based” ✔ “Weather/human based” Absolute Zero: Theoretical temp where all atomic movements stops Example 1.3 Temperature Conversions a) A certain solder has a melting point of 224°C. What is its melting point in Fahrenheit? (b) Helium has the lowest boiling point of all the elements at -452°F. Convert this temperature to degrees Celsius. (c) Mercury melts at only -38.9°C. Convert the melting point to Kelvin. Inversely proportional: as one increases, the other decreases Directly proportional: as one increases, so does the other Matter - anything that occupies space and has mass mass – measure of the quantity of matter SI unit of mass is the kilogram (kg) 1 kg = 1000 g = 1 x 103 g weight – force that gravity exerts on an object weight = mass g (F = ma) on earth, g = 1.0 on moon, g ~ 0.1 La Grande K 1 Kg Pt/Ir alloy World’s Roundest Object https://www.youtube.com/watch?v=ZMByI4s-D-Y A 1 kg bar will weigh 1 kg on earth 0.1 kg on moon International System of Units (SI) Base Units Utilized in this class Used as Relative Standards for comparison All other units are derived from these units and are known as Derived Units Velocity: m/s Force: 1 Newton = 1 kg•m/s2 Volume: m3 Prefixes can be used to simplify for extremely large or small quantities of base units “mu” Used most often in this class, be sure to memorize. Prefix examples Driving 321,000 meters to LR = 321 kilometers Radio station 90.9 MHz = 90,900,000 Hz A mosquito weighs 2.5 milligrams (mg) = 0.0025 grams (g) A dust mite weighs 0.0002 meters = 200 micrometers (mm) Conversion factors can be written/used 2 ways -6 6 1 Mm = 10 m 10 Mm = 1 m -3 10 3 10 1km = m meter (base) -2 1 cm = 10 m -3 1 mm = 10 m -6 1 mm = 10 m -9 1 nm = 10 m Or km = 1 m meter (base) 2 10 cm = 1 m 3 10 mm = 1 m 6 10 mm = 1 m 9 10 nm = 1 m I favor the forms using (+) exponents Volume – SI derived unit for volume is cubic meter (m3) (cm3 is more commonly used) 1 cm3 = (0.01 m)3 = 1 x 10-6 m3 1 L = 1000 mL = 1000 cm3 = 1 dm3 1 mL = 1 cm3 Density – SI derived unit for density is kg/m3 1 g/cm3 = 1 g/mL = 1000 kg/m3 *more commonly used density = mass volume m d= V 2 L of Os = 100 lbs Example 1.1 Gold is a precious metal that is chemically unreactive. It is used mainly in jewelry, dentistry, and electronic devices. A piece of gold ingot with a mass of 301 g has a volume of 15.6 cm3. Calculate the density of gold. gold ingots Example 1.2 The density of mercury, the only metal that is a liquid at room temperature, is 13.6 g/mL. Calculate the mass of 5.50 mL of the liquid. m d= V Chemistry In Action On 9/23/99, $125M Mars Climate Orbiter entered Mars’ atmosphere 100 km (62 miles) lower than planned and was destroyed by heat. 1 lb = 1 N 1 lb = 4.45 N Failed to convert English to metric units “This is going to be the cautionary tale that will be embedded into introduction to the metric system in elementary school, high school, and college science courses till the end of time.” Scientific Notation The number of atoms in 12 g of carbon: 602,200,000,000,000,000,000,000 6.022 x 1023 The mass of a single carbon atom in grams: 0.0000000000000000000000199 1.99 x 10-23 We can factor out powers of 10 to simplify very large or small numbers N is the base number between 1 and 10 N x 10n Exponent (n) is a positive or negative integer Scientific Notation • Base x 10exponent • Base number ≥ 1 and < 10 5 320,000 = 3.2 x 10 . 0.000074 . = 7.4 x 10-5 Decimal moved left so (+) Decimal moved right so (–) Scientific Notation Practice Write these in scientific notation • 0.00578 • 579 • 96,000 • 0.0140 Write these in long notation • 2.0 x 103 • 3.58 x 10-4 • 4.651 x 107 • 9.87 x 10-2 Mathematics in Scientific Notation Addition or Subtraction: Must have same exponent 1. Write each quantity with the same exponent n 4.31 x 104 + 3.9 x 103 = 4.31 x 104 + 0.39 x 104 = 2. Combine N1 and N2 3. The exponent, n, remains the same 4.70 x 104 Tip: change smaller number to match larger exponent 1.36 x 10-1 – 4 x 10-3 = 1.36 x 10-1 – 0.04 x 10-1 = = 1.32 x 10-1 Mathematics in Scientific Notation Multiplication: Add exponents 1. Multiply N1 and N2 2. Add exponents n1 and n2 (4.0 x 10-5) x (7.0 x 103) = (4.0 x 7.0) x (10-5+3) = 28 x 10-2 = 2.8 x 10-1 Division: Subtract exponents 1. Divide N1 and N2 2. Subtract exponents n1 and n2 8.5 x 104 ÷ 5.0 x 109 = (8.5 ÷ 5.0) x 104-9 = 1.7 x 10-5 Bell Ringer a) Write in scientific notation • 8,705,000 m • 0.0000045 L • 0.00237 sec • 9,300 g b) Rewrite above numbers using the nearest SI prefix c) Perform the below mathematics in Sci. Notation • (9.01 x 103 g) + (3.8 x 102 g) • (2.61 x 107 m) x (9.87 x 10-2 m) • (3.98 x 10-2 m) – (8.2 x 10-3 m) • (8.4 x 109 g) ÷ (2.0 x 104 mL) Precision indicates to what degree we know our measurement. (Arithmetic precision) A measurement of 8.0 grams could be made on an average countertop food scale (balance). (~$20) A high-precision milligram scale could weigh the same sample with a much higher precision (8.0235 grams) (~$1,500) Significant Figures: Used to prevent uncertainty from rounding of various measured quantities with various levels of precision. 1) Any digit that is not zero is significant 1.234 kg 34,000 mm 4 significant figures 2 significant figures 2) Zeros between nonzero digits are significant 606 cm 50,050 s 3 significant figures 4 significant figures 3) Zeros to the left of the first nonzero digit are not significant 0.08 mL 0.00054 ML 1 significant figure 2 significant figures 4) If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg 20.000 g 2 significant figures 5 significant figures 5) If a number is less than 1, then only the zeros at the end are significant 0.00420 g 3 significant figures 0.1000 g 4 significant figures Significant Figures Every significant figure is shown when using Scientific notation. ____ m 0.001400 4 significant figures 1.400 x -3 10 Not 1.4 x -3 10 __ mL 500 2 significant figures 5.0 x 2 10 Not 5 x 2 10 Exampl 1.4 Unit Conversions e Determine the number of significant figures in the following measurements: (a) 478 cm (d) 0.0430 kg 22 10 (b) 600,001 g (e) 1.310 × (c) 0.85 m (f) 7000 mL atoms Example 1.4 Solution (a) 478 cm -- Three, because each digit is a nonzero digit. (b) 600,001- Six, because zeros between nonzero digits are significant. (c) 0.825 m -- Three, because zeros to the left of the first nonzero digit do not count as significant figures. (d) 0.0430 kg -- Three. The zero after the nonzero is significant because the number is less than 1. (e) 1.310 × 1022 atoms -- Four, because the number is greater than one so all the zeros written to the right of the decimal point count as significant figures. Example 1.4 solution (f)7000 mL -- This is an ambiguous case. The number of significant figures may be four (7.000 × 103), three (7.00 × 103), two (7.0 × 103), or one (7 × 103). This example illustrates why scientific notation must be used to show the proper number of significant figures. If no decimal is present it is usually assumed only non-zeros are significant. If a decimal is present, than all zero’s are significant. 7,000 mL ≠ 7,000. mL They display differing degrees of precision. Significant Figures Addition or Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers. Use the least precise number. 89.392 L + 1.1XX 90.492 ± 50 mL one significant figure after decimal point round off to 90.5 ± 1.0 mL 3.70XX -2.9133 0.7867 two significant figures after decimal point round off to 0.79 Significant Figures Multiplication or Division The number of significant figures in the result is set by the original number that has the smallest number of significant figures. 4.51 x 3.0006 = 13.532706 = 13.5 round to 3 sig figs 3 sig figs 6.8 ÷ 112.04 = 0.0606926 = 0.061 2 sig figs round to 2 sig figs Example 1.5 Carry out the following arithmetic operations to the correct number of significant figures: (a) 11,254.1 g + 0.1983 g (b) 66.59 L − 3.113 L (c) 8.16 m × 5.1355 kg (d) 0.0154 kg ÷ 88.3 mL (e) (2.64 × 103 cm) + (3.27 × 102 cm) Example 1.5 Solution Solution In addition and subtraction, the number of decimal places in the answer is determined by the number having the lowest number of decimal places. (a) (b) Example 1.5 Solution In multiplication and division, the significant number of the answer is determined by the number having the smallest number of significant figures. (c) (d) (e) First we change 3.27 × 102 cm to 0.327 × 103 cm and then carry out the addition (2.64 cm + 0.327 cm) × 103. Following the procedure in (a), we find the answer is 2.97 × 103 cm. Bell Ringer a) Perform the below mathematics in Sci. Notation. using Significant Figures in your answer. 1. (9.8 x 105 g) + (6.75 x 104 g) 2. (5.98 x 10-6 m) – (7 x 10-8 m) b) Rewrite the first 2 solutions using the nearest SI prefix 3. (2.612 x 1010 m) x (9.87 x 10-3 m) 4. (7 x 102 mg) ÷ (1.875 x 104 mL) Significant Figures Exact Numbers Numbers from definitions or numbers of objects are considered to have an infinite number of significant figures. •The average of three measured lengths: 6.64, 6.68 and 6.70? 6.64 + 6.68 + 6.70 3 = 6.67333 = 7 = 6.673 Because 3 is an exact number, not a measured number; It is not used for sigfigs. • How many feet are in 6.82 yards? 6.82 yards x 3 ft/yard = 20.5 ft = 20 ft 1 yard = exactly 3 ft by definition Accuracy – how close a measurement is to the true value Precision – how close a set of measurements are to each other accurate & precise precise but not accurate not accurate & not precise Percent Error A way to determine how accurate your measurements are to a known value. |Obtained value – Actual value| x 100% Actual Value Ranges between 0 and 100% Ex. I weigh a 3 kg block on three different scales: 3.2 kg, 3.0 kg, 3.1 kg = 3.1 kg average 3.1 – 3.0 x 100% = 3.3% error 3.0 Dimensional Analysis of Solving Problems (Train-Tracks) 1. Determine which unit conversion factors are needed 2. Carry units through calculation 3. If all units cancel except for the desired unit(s), then the problem was solved correctly. given quantity x conversion factor = desired quantity given unit x desired unit given unit = desired unit Train Track Example How many inches are in 3.0 miles? Identify beginning information Draw a train track 3 miles Write measurement as a fraction Train Track Example How many inches are in 3.0 miles? • We are going from a larger measurement to a smaller one. • Find a conversion factor you know that changes miles into something smaller. Conversion Factor: 1 mile = 5,280 feet • Write your conversion factor on the track so that miles cancels out and you are left with the unit feet. 3 miles 5280 feet 1 mile Always need same units on opposite sides to cancel out Train Track Example How many inches are in 3.0 miles? We now need another conversion factor between Feet and Inches: 1 foot = 12 inches Again, place conversion factor so that the previous unit cancels out. 3.0 miles 5280 feet 1 mile 12 inches 1 foot Train Track Unit Conversions How many inches are in 3.0 miles? 3.0 miles 5,280 feet 1 mile 12 inches 1 foot Inches are the only remaining unit ✔ Multiply all numbers on the top Divide all numbers on the bottom 3.0 x 5,280 x 12 1x1 = 190,080 inches = 1.9 x 105 inches (2 sig figs) More practice: Convert 1.40 x 10-6 Mg to cg Example Metric to Metric Conversion Problems Convert 2.79 x 105 mm to km Don’t try to convert directly from mm to km. Go to the base unit (m) first Conversion factors: 1,000 mm = 1 m; 1,000 m = 1 km 2.79 x 105 mm 1m 103 mm 1 km 103 m Kilometers are the only remaining units ✔ 2.79 x 105 = 2.79 x 10(5-3-3) 103 x 103 = 2.79 x 10-1 km (3 sig figs) More practice: Convert 3.4 x 109 cg to Mg Example A person’s average daily intake of glucose (a form of sugar) is 0.0833 pound (lb). What is this mass in milligrams (mg)? (1 lb = 453.6 g.) A metric conversion is needed to convert grams to milligrams (1 mg = 1 × 10−3 g) (Or we could write: 1,000 mg = 1 g) Either Conversion factor will work 0.0833 lb 453.6 g 1 lb 103 mg 1g = 37,784.88 mg = 3.78 x 104 mg (3 sig figs) Example 2-D Conversion Problems (Unit1/Unit2) Convert 70.0 miles/hour to m/s We convert one unit at a time, followed by the other Conversion factors: 1 mile = 1,609 meters; 1 hour = 60 min; 1 min = 60 sec *Note: to cancel out hours (on bottom) it must appear again on the top 70.0 miles 1 hour 1609 meter 1 mile 1 hour 60 min 1 min 60 sec Meter/sec are the only remaining units ✔ 70.0 x 1,609 x 1 x 1 1 x 1 x 60 x 60 = 31.286 m/s = 31.3 m/s More practice: Convert 3.4 kg/L to g/mL 2-D Conversion Problems (Unit#) Example Convert 2.5 x 10-5 m3 to mm3 Conversion factors: 1 m = 1,000 mm 1 m3 ≠ 1,000 mm3 1 m3 = (1,000)3 mm3 ✔ 1. Write the 1-D units first (1m = 103m) 2. Add exponent to entire conversion factor 3 2.5 x 10-5 m3 1000 mm 1m = 2.5 x 10-5 m3 109 mm3 1 m3 2.5 x 10-5 x 109 = 2.5 x 104 mm3 More practice: Convert 34 yd2 to ft2 Example Convert 2-D Conversion Problems # (Unit ) 6.70 x 103 ft2 to inches2 Conversion factors: 1 ft = 12 in 1 ft2 ≠ 12 in2 1. Write the 1-D units first (1ft = 12 in) 1 ft2 = 122 in2✔ 6.70 x 103 ft2 12 in. 1 ft Alternate 2nd Method 2. Write the same conversion factor again until they cancel 12 in. 1 ft = 9.65 x 105 in2 More practice: Convert 34 yd3 to ft3 Example An average adult has 5.2 L of blood. What is the 3 volume of blood in m ? 1 mL = 1 cm3 5.2 L 103 mL 1L 1 cm3 1 mL 1 m3 1003 cm3 = 5.2 x -3 10 3 m Example 2-D Conversion Problem The circumference of the earth is approximately 2.49 x 104 miles long. If the speed of sound travels at 760 mph, how many days would it take a sound wave to circulate Earth’s circumference. Starting with length and velocity; needing time Conversion factor: 760 miles = 1 hour 2.49 x 104 miles 1 hour 760 miles 1 day 24 hours = 1.4 days More Practice Conversion problems • Convert 3.0 mL to ounces (33.8 oz = 1 L) • 42.0 km/h to ft/ms • 1.67 Mm to mm • 0.55 Acres to m2 (247 acre = 1 km2) • 2.35 x 1012 inches to cm (1 ft = 0.305 m) • 10.6 g/mm3 to kg/m3 • 3.50 x 104 mL to cL Review Unit Conversion & Significant Figures: Crash Course Chemistry #2 www.youtube.com/watch?v=hQpQ0hxVNTg Sample of Topics to Study • Democritus • Alchemy • 5 Fields of Chemistry • Macro- / Microscopic • Quantitative /Qualitative • Scientific: Theory / Law • Matter / Substance • Mixture: Hetero- / Homogenous • Physical Properties • Extensive / Intensive Properties • Chemical change properties • Elements/Atoms/Symbols • Subatomic particles • Compound • Temperature scale conversions • Phases of Matter • Metric Units • Using Prefixes • Significant Figures (+ math) • Scientific Notation (+ math) • Dimensional Analysis • Accuracy • Precision • % Error