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Chemistry: The Study of Change

Chemistry: The Study of Change. Chapter 1. The Study of Chemistry. Macroscopic. Microscopic. 1.2. Water. Sugar. Gold. Chemistry is the study of matter and the changes it undergoes. Matter is anything that occupies space and has mass. . What about air? Yes it is matter. 1.4.

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Chemistry: The Study of Change

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  1. Chemistry: The Study of Change Chapter 1

  2. The Study of Chemistry Macroscopic Microscopic 1.2

  3. Water Sugar Gold • Chemistry is the study of matter and the changes it undergoes. • Matter is anything that occupies space and has mass. What about air? Yes it is matter 1.4

  4. MAJOR AREAS OF CHEMISTRY • Organic Chemistry - the study of matter which is carbon-based. • Inorganic Chemistry - the study of matter containing all other elements (Inorganic is everything else). • Analytical Chemistry - analyze matter to determine identity and composition (involves qualitative and quantitative) (eg. Determination of Alcohol in blood) • Biochemistry - the study of life at the molecular level (DNA sequencing) • Physical Chemistry - attempts to explain the way matter behaves (eg. Kinetics)

  5. Chemistry Related Other Fields Medical Science Pharmaceutical Industry Oil Industry Paints and Coatings

  6. Chemistry: A Science for the 21st Century • Health and Medicine • Sanitation systems • Surgery with anesthesia • Vaccines and antibiotics • Energy and the Environment • Fossil fuels • Solar energy • Nuclear energy

  7. Chemistry: A Science for the 21st Century • Materials and Technology • Polymers, ceramics • Food and Agriculture • Genetically modified crops • Specialized fertilizers

  8. An element is a substance that cannot be separated into simpler substances by chemicalmeans (Listed in the Periodic Table) • 114 elements have been identified • 82 elements occur naturally on Earth gold, aluminum, lead, oxygen, carbon • 32 elements have been created by scientists technetium, americium, seaborgium • Metals, nonmetals, metalloids Latin names: aurum (Au), ferrum (Fe),natrium (Na)

  9. Classification of Matter • Based on Physical State: • Gas: does not have a fixed shape or a fixed volume. Both the volume & the shape of the gas is that of the container. • Liquid: Has a fixed volume but variable shape = shape of the container. • Solid: Has fixed shape and volume

  10. gas solid liquid The Three States of Matter (water in three forms)

  11. Classification of Pure Substances • Elements: These only have 1 type of atoms. • 2. Compounds: composed of atoms of two or more elements. eg. H2O chemically united in a fixed proportion – 2 H atoms and 1 O atom. It can be separated only chemically. • Ammonia (NH3) – 1N, 3H (combined at high pressure) • Elements exist in monoatomic or polyatomic form. For eg., Hydrogen and oxygen exist as H2 and O2 , that is diatomic molecules where as Ca exist as monoatomic. • Chemical Symbols: Cobalt – Co, Sodium – Na (natium-Latin) • CO is not cobalt, carbon monoxide Pure Substances

  12. soft drink, milk, sugar solution cement, iron filings in sand A mixture is a combination of two or more substances in which the substances retain their distinct identities. • Homogenous mixture – composition of the mixture is the same throughout. • Heterogeneous mixture – composition is not uniform throughout.

  13. magnet distillation Physical means can be used to separate a mixture into its pure components.

  14. Physical and Chemical Properties • Physical Properties: will keep the identity of the substance. Eg. boiling point (B.P), melting point (M.P), density, mass, volume, area. (water changes the physical state by heating and cooling, not the composition, so the M.P and B.P are physical properties) • Chemical Properties: involves chemical change. eg. rusting, burning of H in O H2O Needs a chemical change to bring the original substances, not physical like boiling or melting.

  15. Extensive and Intensive Properties An extensive property of a material depends upon how much matter is being considered - additive • mass (M) • length • volume (V) An intensive property of a material does not depend upon how much matter is is being considered. e.g., density of water 0.99713g/mL at 25 C • density – M/V • color

  16. Hypothesis, Theory, Law • A hypothesis is an educated guess, based on observation. Needs to prove by experiments to be true. • A scientific theory summarizes a hypothesis or group of hypotheses that have been supported with repeated testing. • A law generalizes a body of observations. • Eg., Newton's Law of Gravity • F = Gm1m2/r2

  17. Classifications of Matter

  18. Measurement in Chemistry • In SI units, not in English units. • Meaurement is always necessary to collect data • Mass, volume etc. with proper equipments or glassware • Always use unit. Mass = 5 for a salt is useless. Needs unit like g or kg. • A measurement is useless without its units. • weight – force that gravity exerts on an object

  19. International System of Units (SI) Revised Metric System

  20. Truly systematic • 1 meter = 10 decimeters = 100 centimeters Basic Units of the Metric System Mass gram g Length meter m volume liter L

  21. 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 In Chemistry, the smaller g is more convenient. weight – force that gravity exerts on an object

  22. Volume – SI derived unit for volume is cubic meter (m3) 1 cm3 = (1 x 10-2 m)3 = 1 x 10-6 m3 1 dm3 = (1 x 10-1 m)3 = 1 x 10-3 m3 1 L = 1000 mL = 1000 cm3 = 1 dm3 1 mL = 1 cm3

  23. mass density = volume A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass? m m d = d = V V Density – SI derived unit for density is kg/m3 1 g/cm3 = 1 g/mL = 1000 kg/m3 Chemical Application: g/cm3 or g/mL = 21.5 g/cm3 x 4.49 cm3 = 96.5 g m = d x V

  24. Derived Units in SI • Area: (length x width); unit = m x m = m2 • Volume: length x width x height; unit = m x m x m = m3 • Density: mass/volume; unit = kg/m3 • Speed: distance/time; unit = m/s • Acceleration: speed/time; unit = m/s/s = m/s2 • Force: mass x acceleration; unit = kg m/s2 = N (newton) • Energy: force x distance; unit = kg m/s2 x m = kg m2/s2 = J (joule) • Pressure: force/Area; unit = kg m/s2 ÷ m2 = kg/ms2 = Pa (Pascal)

  25. 0F = x 0C + 32 9 5 K = 0C + 273.15 273 K = 0 0C 373 K = 100 0C 32 0F = 0 0C 212 0F = 100 0C

  26. Conversions between Fahrenheit and Celsius 1. Convert 75oC to oF. 2. Convert -10oF to oC. 1. Ans. 167 oF 2. Ans. -23oC

  27. 0F – 32 = x 0C 0F = x 0C + 32 x (0F – 32) = 0C 5 5 5 0C = x (0F – 32) 0C = x (172.9 – 32) = 78.3 9 9 9 9 9 5 5 Convert 172.9 0F to degrees Celsius.

  28. Chemistry In Action On 9/23/99, $125,000,000 Mars Climate Orbiter entered Mar’s atmosphere 100 km (62 miles) lower than planned and was destroyed by heat (read p.17, ch.1) 1 lb = 1 N 1 lb = 4.45 N F=ma=4.45Newton “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.”

  29. The number of atoms in 12 g of carbon: 602,200,000,000,000,000,000,000 The mass of a single carbon atom in grams: 0.0000000000000000000000199 Scientific Notation N x 10n 6.022 x 1023 1.99 x 10-23 N x 10n N is a number between 1 and 10 n is a positive or negative integer

  30. move decimal left move decimal right Scientific Notation 568.762 0.00000772 n > 0 n < 0 568.762 = 5.68762 x 102 0.00000772 = 7.72 x 10-6 Addition or Subtraction • Write each quantity with the same exponent n • Combine N1 and N2 • The exponent, n, remains the same 4.31 x 104 + 3.9 x 103 = 4.31 x 104 + 0.39 x 104 = 4.70 x 104

  31. Scientific Notation Multiplication (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 • Multiply N1 and N2 • Add exponents n1and n2 Division 8.5 x 104÷ 5.0 x 109 = (8.5 ÷ 5.0) x 104-9 = 1.7 x 10-5 • Divide N1 and N2 • Subtract exponents n1and n2

  32. Scientific Notation • The measuring device determines the number of significant figures a measurement has. • In this section you will learn • to determine the correct number of significant figures (sig figs) to record in a measurement • to count the number of sig figs in a recorded value • to determine the number of sig figs that should be retained in a calculation.

  33. Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit.

  34. Significant Figures • Any digit that is not zero is significant 1.234 kg 4 significant figures • Zeros between nonzero digits are significant 606 m 3 significant figures • Zeros to the left of the first nonzero digit are not significant 0.08 L 1 significant figure • If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg 2 significant figures • If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant 0.00420 g 3 significant figures

  35. How many significant figures are in each of the following measurements? 24 mL 2 significant figures 4 significant figures 3001 g 0.0320 m3 3 significant figures 6.4 x 104 molecules 2 significant figures 2 significant figures 560 kg (5.6 x 102 )

  36. 89.332 + 1.1 one digit after decimal point two digit after decimal point 90.432 round off to 90.4 round off to 0.79 3.70 -2.9133 0.7867 Significant Figures Addition or Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers.

  37. move decimal left move decimal right Scientific Notation 568.762 0.00000772 n > 0 n < 0 568.762 = 5.68762 x 102 0.00000772 = 7.72 x 10-6 Addition or Subtraction • Write each quantity with the same exponent n • Combine N1 and N2 • The exponent, n, remains the same 4.31 x 104 + 3.9 x 103 = 4.31 x 104 + 0.39 x 104 = 4.70 x 104 1.8

  38. 3 sig figs round to 3 sig figs 2 sig figs round to 2 sig figs 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 3.6666 = 16.536366 = 16.5 6.8 ÷ 112.04 = 0.0606926 = 0.061

  39. 6.64 + 6.68 + 6.70 = 6.67333 = 6.67 = 7 3 Significant Figures 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?

  40. 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

  41. UNIT CONVERSION • You need to be able to convert between units • within the metric system • between the English and metric system • The method used for conversion is called the Factor-Label Method or Dimensional Analysis ALWAYS NEEDED

  42. desired unit given unit Dimensional Analysis Method of Solving Problems • Determine which unit conversion factor(s) are needed • Carry units through calculation • 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

  43. We use these two mathematical facts to do the factor label method • a number divided by itself = 1 • Example: 2/2 = 1 • any number times one gives that number back • Example: 2 x 1 = 2 • Example: How many donuts are in 3.5 dozen? • You can probably do this in your head but let’s see how to do it using the Factor-Label Method.

  44. Start with the given information... 3.5 dozen = 42 donuts Then set up your unit factor... See that the units cancel... Then multiply and divide all numbers...

  45. How many mL are in 1.63 L? 1000 mL 1L L2 1.63 L x = 1630 mL mL 1L 1.63 L x = 0.001630 1000 mL Dimensional Analysis Method of Solving Problems Conversion Unit 1 L = 1000 mL

  46. 60 min m x x x 343 60 s 1 mi s 1 hour = 767 1 min 1609 m mi hour The speed of sound in air is about 343 m/s. What is this speed in miles per hour? conversion units meters to miles seconds to hours 1 mi = 1609 m 1 min = 60 s 1 hour = 60 min

  47. Common Relationships Used in the English System A. Weight 1 pound = 16 ounces 1 ton = 2000 pounds B. Length 1 foot = 12 inches 1 yard = 3 feet 1 mile = 5280 feet C. Volume 1 gallon = 4 quarts 1 quart = 2 pints 1 quart = 32 fluid ounces Commonly Used “Bridging” Units for Intersystem conversion Quantity English Metric Mass 1 pound = 454 grams 2.2 pounds = 1 kilogram Length 1 inch = 2.54 centimeters 1 yard = 0.91 meter Volume 1 quart = 0.946 liter 1 gallon = 3.78 liters

  48. Examples of Unit Conversion 1. Convert 5.5 inches to millimeters (139.7 mm) 2. Convert 50.0 milliliters to pints (0.1057 pint) • Convert 1.8 in2 to cm2 (11.613cm2) 1.3 Measurement in Chemistry

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