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CHEM 120: Introduction to Inorganic Chemistry

CHEM 120: Introduction to Inorganic Chemistry. Instructor: Upali Siriwardane (Ph.D., Ohio State University) CTH 311, Tele: 257-4941, e-mail: upali@chem.latech.edu Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00 and 11:00-12:00 M,W,& F. Chapters Covered and Test dates.

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CHEM 120: Introduction to Inorganic Chemistry

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  1. CHEM 120: Introduction to Inorganic Chemistry Instructor: Upali Siriwardane (Ph.D., Ohio State University) CTH 311, Tele: 257-4941, e-mail: upali@chem.latech.edu Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00 and 11:00-12:00 M,W,& F

  2. Chapters Covered and Test dates • Tests will be given in regular class periods  from  9:30-10:45 a.m. on the following days: September 21,     2004 (Test 1): Chapters 1 & 2 • October 6,         2004(Test 2):  Chapters  3, & 4 • October 20,         2004 (Test 3): Chapter  5 & 6 • November 3,        2004 (Test 4): Chapter  7 & 8 • November 15,      2004 (Test 5): Chapter  9 & 10 • November 17,      2004 MAKE-UP: Comprehensive test (Covers all chapters • Grading: • [( Test 1 + Test 2 + Test3 + Test4 + Test5)] x.70 + [ Homework + quiz average] x 0.30 = Final Average •                               5

  3. What is Chemistry Chemistry Study of the composition, structure and properties of matter and the changes (including energy) that matter undergoes. Major Areas of Chemistry Analytical Biological (Biochemistry) Inorganic Organic Physical. Chemistry: the ‘central’ science

  4. The Scientific Method Discovery of Penicillin Alexander Fleming 1. Observation. 2. Formulation of a question 3. Pattern recognition 4. Developing theories. (hypothesis and eventfully theory) 5. Experimentation 6. Summarizing information.( scientific laws)

  5. Matter and properties • Properties are characteristics of matter and are classified as physical or chemical properties. • Matter can exist in 3 physical states.

  6. The three states of matter • _____: atoms or molecules are close together in orderly array; fixed volume and shape. • ______: atoms or molecules are close together but relatively free to move about; has a definite volume but takes on shape of container • ___: atoms or molecules are widely separate and free to move about; have no definite volume or shape--expand to fill entire container

  7. Solid, liquid, gas

  8. Physical Properties • _________________ (and change): measured and observed without changing composition of substance; e.g: color, mass, density, volume, melting and boiling points , odor taste • water boils (physical change) • Chemical identity of substance unchanged.

  9. Chemical Properties • __________________: must carry out chemical change to observe; e.g., burning gasoline, smelling perfume, digesting sugar. Atoms rearrange to form new substances • Chemical identities of substances change.

  10. Classify as chemical or physical change • Water boiling to become steam • Butter becoming rancid • Combustion of wood • Melting of ice in spring • Decay of leaves in winter

  11. Intensive and extensive properties • __tensive property: depends on amount of substance; e.g., mass, volume, energy • __tensive properties are additive • __tensive property: does not depend on amount of substance; e.g., temperature, density, pressure. • __tensive properties are not additive

  12. Classification of matter • All matter is either a pure substance or a mixture.

  13. Pure Substances A. Substance (pure): 1. has only one component 2. definite unvarying composition (as carbon dioxide: 27.3 % carbon, 72.7% oxygen) 3. uniform properties throughout 4. cannot be separated into other components or further purified by physical means.

  14. Substances can be • 1. _________: substances that cannot be separated into simpler substances by chemical means. • 2.__________: substances composed of 2 or more elements chemically united in fixed proportions (as CO2). Compounds can be separated into the elements that make it up by chemical (not physical) means

  15. Mixtures • B. Mixtures: 2 or more substances that can be: 1. combined in any proportion 2. separated into the substances that make it up by physical means

  16. Mixtures can be: • 1. _____ogeneous : composition is not uniform (granite, orange juice) • 2. ____ogeneous: composition is uniform throughout. Solutions are mixtures (salt water, brewed tea and coffee)

  17. Identify as pure substance, homogeneous of heterogeneous mixture • Air • Paint • Perfume • Carbon monoxide

  18. Data, Results and Units • Chemistry is a quantitative science based on experimentation. Take measurements (data). • Each piece of data is the outcome of a measurement. • Results are the outcome of experiments (generally several pieces of data).

  19. Chemistry is a quantitative science based on experimentation. Take measurements. • A measurement is a number and a unit that describes what the measurement is measuring. • In chemistry we use the metric system for taking measurements. We also use scientific notation to make life easier.

  20. Scientific notation • Based on powers of 10 • 1000 = 1 x 10 x 10 x 10 = 1 x 103 • 56,000,000 = 5.6 x 10,000,000 or 5.6 x 107 • 0.000068 = 6.8 x 1/100,000 or 6.8 x 10-5

  21. Express in scientific notation: • 0.000570 • 248,000,000 • Express in decimal notation: • 3.44 x 10-4 • 1.45 x 105 • Move decimal one place to right, , exponent • Move decimal one place to left, exponent

  22. Brief review • Addition or subtraction: both numbers must be expressed as the same power of 10: 145.756 + 5.3 x 10-2 • Multiplication : N1 x N2, add exponents of the nos. • Division: N1/N2 , subtract the exponents of N2 from N1

  23. Exponents on your calculator • 6.56 x 108 (7.054 x 10-6)(4.9 x 1012) • To put exponents in your calculator, use the EE or EXP key!!!!!

  24. Metric system • Base unit as meter (m),(for length), gram (g) (for mass), liter (L) (for volume), etc • In order to express our measuremenmts as #s that are close to whole #s, we use the following prefixes. Prefixes that tell us how far the measurement is from the base unit.

  25. Learn this table!!!! • Prefix symbol meaning Mega- M 106 (1,000,000) • Kilo- k 103 (1,000) • Deka- da 101 (10) • base unit 100 (1) • Deci- d 10-1 (0.1) • Centi- c 10-2 (0.01) • Milli- m 10-3 (0.001) • Micro- m 10-6 (0.000001) • Nano- n 10-9 (0.000000001)

  26. 1 ns (nanosecond) = ? • 1 ns (nanosecond) = • 1mm (millimeter) = ? • 1mm (millimeter) = • 1kg (kilogram) = ? • 1kg (kilogram) = • 1 cL (centiliter) = ? • 1 cL (centiliter) =

  27. Significant figures • How many nos. can I write down, legitimately, when I make a measurement? • You are allowed to write down all nos. that are measured with certainty plus one that is estimated, The no. of the recorded figures (including the estimated no.) are the no. of sig. fig. in the measurement.

  28. How to figure the no. of sig fig in a recorded no. • 1.All nonzero digits are significant 965 12,456 3.41567 • Zeroes may be troublesome. • 2. Zeroes btn 2 nonzero digits are sig. 10.143 20006 230807 3.00001 • 3. Zeroes that come before the first nonzero digit (after the decimal pt) are not sig. 0.000000045 (4.5 x 10-8)

  29. 4. Zeroes at the end of a no. and after a decimal are significant (trailing zeros). 0.076800 10.076800 • 5. Zeroes at the end of a no. without a decimal are ambiguous. 4500

  30. Rounding off of data • 1. If the first "extra" digit is LESS than 5-drop it.  Now the last digit of the number remains the same. • Ex. 4.321 becomes 4.32 • 2. If the first "extra" digit is 5 or MORE than 5,  drop the number and increase the last significant digit by 1. • Ex. 4.336 becomes 4.34 • What is this "even/odd rule" I keep hearing about? • When digit is exactly 5: option will give an even number as the answer (last digit is 0, 2, 4, 6, or 8). Refer to the examples above.

  31. Addition and subtraction rules • Rule for addition and subtraction: the answer can’t have more nos. after the decimal pt than the original no with the fewest nos after the decimal. • What is the result of adding 8.355, 4.687 x 10-3 and 99.1568?

  32. Rule for multiplication and division:the answer can’t have more sig. fig. Than the no. with the least no. of sig. fig. • 4.56 x 106 (7.954 x 10-8)(4.5 x 1011)

  33. What is the answer to the correct no. of significant figures of... • 1.446 x 10-3 - 2.91 x 10-5 • (12.675)(10.03) (9.44 + 6.885)

  34. Experimental quantities that we measure • Mass (amt of matter in an object--location independent) in kilograms(kg) (SI units) but grams (g) more convenient • Weight: force that gravity exerts on an object--location dependent • weight = mass x acceleration due to gravity, • When gravity is constant, mass and weight are directly proportional. (Gravity depends on distance from center of the earth)

  35. Units of mass • 1 pound = 16 ounces; 1 ton = 2000. lbs • 1 pound = 454 g; 2.2 lbs = 1 kg = 103 g • 1 amu (atomic mass unit) = 1.661 x 10-24 g

  36. Length: distance btn 2 points • Length in meters ( SI units); frequently cm: 1 cm = 10-2m • 1 ft = 12 inches; 1 yd = 3 ft; 1 mile = 5280ft • 1 inch = 2.54 cm; 1 yd = 0.91 m (1 m = 39. 56 in)

  37. Volume (amt of space occupied by matter) in m3(SI units) ; • more conveniently liters (1 L = 1 dm3) • or milliliters (1 mL = 1 cm3) • 1 gallon = 4 qts; 1 qt = 2 pts = 32 fluid oz (4 cups = 1 qt) • 1 qt = 0.946 L (1L = 1.06 qt); 1 gal = 3.78 L

  38. Conversion from one unit to another: factor label method • 1. Identify the problem. • 2. Write down all conversion factors that let you go from the given units to the units that you want your answer to be in. (Or write down the relevant equation.) • 3. Set up conversion factors so units cancel and you’re left with the units you want your answer in. (Or substitute in the eqn and solve.)

  39. Convert 32.0 oz to lbs, tons, grams, milligrams 1 lb = 454 g • How many miles in 1.0 km? 1 in = 2.54 cm • Convert 3.0 km to inches, yds, millimeters. • Convert 10.0 pts to qts, gals, and microliters 1 qt = 0.946 L • Convert 68.3 cm3 to cubic kilometers.

  40. 1.64: A newborn is 21 inches in length and weighs 6 lb 9 oz. Convert to metric units. 1 lb = 454 g 1 in = 2.54 cm

  41. Other things we measure • Concentration: amt of a substance in a given unit volume; as the no of red blood cells/L; the number of rose petals in a given volume in a vase of a given size, etc. • Molarity: Moles of a compound in a liter of solution

  42. In chemistry a useful concentration unit is density. • Density = mass volume • d = m V • Units g/mL(for solids and liquids) or g/L (for gases)

  43. Specific gravity • Specific gravity = density of object(g/mL) density of water (g/mL) • Specific gravity is unitless. Normally use 1.00 g/mL for density of water.

  44. Density problems • A lead sphere has a mass of 1.20 x 104 g and a volume of 1.05 x 103 cm3. Calc the density of lead. • 1.76: What volume, in liters, will 8.00 x 102 g of air occupy if the density of air is 1.29 g/L?

  45. 1.78. What is the mass, in grams, of a femur (leg bone) having a volume of 118 cm3? The density of bone is 1.8 g/cm3 ?

  46. We also measure • Time: in seconds (s) • 60 s = 1 minute; 60 minutes = 1 hour, etc

  47. And temperature ( degree of hotness) • Temperature usually in degrees Celsius (°C) • Related to familiar degrees Fahrenheit (°F) by °C = (°F -32) 1.8 • And oF = 1.8oC + 32

  48. As a check water freezes at 0°C (32oF), boils at 100°C (212oF)

  49. Kelvin scale used in many chemistry calculations • Kelvin is “absolute temperature,” related by K = °C + 273.15 • Water boils at 373 K, freezes at 273 K.

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