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INTRODUCTION

INTRODUCTION. What is Chemistry? Definition: Chemistry is the study of matter and its changes from one substance to another. Chemistry is central to all sciences and overlaps with physics, biology, geology, and astronomy.

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INTRODUCTION

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  1. INTRODUCTION

  2. What is Chemistry? Definition: Chemistry is the study of matter and its changes from one substance to another. Chemistry is central to all sciences and overlaps with physics, biology, geology, and astronomy. Learning about chemistry teaches you about the benefits and risks associated with chemicals and will help you to be an informed citizen and make intelligent choices concerning the world around you. Chemistry teaches you to solve problems and communicate with others in an organized and logical manner.

  3. Paleontology Biology Medicine Biophysics Biochemistry Geochemistry Geology Chemistry Physics Physical Chemistry Chemical Engineering Nuclear Chemistry Cosmo chemistry Astrophysics Astronomy

  4. Major Chemistry subdivisions: • Analytical (qualitative and quantitative) Chemistry: Qualitative Analytical Chemistry: What is a sample of matter composed of? Quantitative Analytical Chemistry: How much “stuff” is in a sample of matter? • Biochemistry: The study of living systems (Biology + Chemistry) • Organic Chemistry: The study of properties and reactions of compounds that contain Carbon • Inorganic Chemistry: The study of properties and reactions of compounds that are not Carbon based • Physical Chemistry: The study of the physics involved with chemical changes.

  5. Major Hints in being a successful Chemistry Student Study Everyday!! A 5 unit class requires at least 10 hours/week of study time to earn a “C” Don’t fall behind – you will never catch up!!! Always ask for help when you need it – don’t wait. Working in groups is very helpful IF you have the right group!!! Read the chapter and/or online lecture notes before class to familiarize yourself with the material.

  6. ELEMENTS

  7. NAMING ELEMENTS The chemical symbol and the name of the elements must be known before starting any discussion in chemistry. There is a correlation between the name and the chemical symbol. The chemical symbol is usually the first or the first and second letter of the name. H stands for the hydrogen atom He is helium Li is lithium Or it could be the first and third letter Mg is magnesium Mn is manganese Or it could be the letters for the “Old” (Latin, Greek, etc.) name Fe is iron (ferrum: latin Cu is copper (cuprum)

  8. NAMES OF ELEMENTS TO KNOW H = hydrogen atom Si = silicon Ag = silver He = helium P = phosphorus Sn = tin Li = lithium S = sulfur Zn = zinc Be = beryllium Cl = chlorine Ba = barium B = boron Ar = Argon Ni = nickel C = carbon K = potassium (kalium) Pb = lead N = nitrogen atom Ca = calcium Xe = xenon O = oxygen atom Cr = chromium Au = gold F = fluorine atom Fe = iron (ferrum) Ne = neon Hg = mercury Na = sodium (natrium) Cu = copper (cuprum) Mg = magnesium I = iodine atom Al = aluminum Br = bromine atom

  9. NAMING ELEMENTS There are 7 homonuclear diatomic elements. These elements do not exist “alone” readily under the conditions experienced on the surface of the Earth. H - hydrogen atom H2 - hydrogen N - nitrogen atom N2 - nitrogen O - oxygen atom O2 - oxygen F - fluorine atom F2 - fluorine Cl - chlorine atom Cl2 - chlorine Br - bromine atom Br2 - bromine I - iodine atom I2 - iodine You must be extra careful (& specific) when naming these elements

  10. Reading Chemical Formulas of Compounds BaCO3 (NH4)3PO4 CuSO4•5H2O CH3COOH CuCO3•Cu(OH)2

  11. SCIENTIFIC METHOD

  12. SCIENTIFIC METHOD 1. FACT : An observable event; indisputable evidence which does not explain but simply is. 2. HYPOTHESIS: A guess to try to explain an observation. 3. EXPERIMENT: A systematic exploration of an observation or concept. 4. THEORY:An explanation of the facts; it can be proven by experiment and it confirms an hypothesis. 5. LAW:A theory which has undergone rigorous experimentation and no contradiction can be found. Note: MODEL: A visual or mathematical device or method used to demonstrate a theory or concept.

  13. EXPERIMENTS In ancient Greece philosophers, like Aristotle, did very little physical experimentation. Mental exercises were the preferred method for solving problems. SCIENTIFIC METHOD depends on experimentation therefore the ability to make measurements is vital when using scientific method.

  14. ACCURACY vs. PRECISION • Accurate & preciseinaccurate but precise • inaccurate & imprecise

  15. MEASUREMENTSRounding Off Numbers Rule 1. When the first digit after those you want to retain is 4 or less, that digit and all other to its right are dropped. The last digit retained is not changed. The following examples are rounded off to four digits: 74.693 ≡ 74.69 1.00629 ≡ 1.006 Rule 2. When the first digit after those you want to retain is 5 or greater, that digit and all others to the right are dropped and the last digit retained is increased by one. These examples are rounded off to four digits: 1.026868 ≡ 1.027 18.02500 ≡ 18.03 12.899 ≡ 12.90

  16. MEASUREMENTSSignificant Figures 1. All nonzero numbers are significant figures. 2. Zero’s follow the rules below. Zero’s between numbers are significant. 30.09 has 4SF Zero’s that precede are NOT significant. 0.000034 has 2SF Zero’s at the end of decimals are significant. 0.00900 has 3 SF Zero’s at the end without decimals are either. 4050 has either 4SFor 3SF

  17. MEASUREMENTSSignificant Figures & Calculations Significant figures are based on the tools used to make the measurement. An imprecise tool will negate the precision of the other tools used. The following rules are used when measurements are used in calculations. Adding/subtracting The result should be rounded to the same number of decimal places as the measurement with the least number of decimal places. Multiplying/dividing The result should contain the same number of significant figures as the measurement with the least number of significant figures.

  18. Adding & Subtracting MEASUREMENTSSignificant Figures & Calculations 345.678 + 12.67 0.07283 - 0.0162789 1587 - 120 358.348 0.0565511 1467 All Answers are Incorrect!!! 358.35 0.05655 1470 or 1.47 x 103 Multiplication & Division (12.034)(3.98) = 47.89532 47.9 is correct 98.657 ÷ 43 = 2.294348837 2.3 is correct (13.59)(6.3) = 12 7.13475 7.1 is correct

  19. MEASUREMENTSScientific Notation Many measurements in science involve either very large numbers or very small numbers (#). Scientific notation is one method for communicating these types of numbers with minimal writing. GENERIC FORMAT: # . # #… x 10# A negative exponent represents a number less than 1 and a positive exponent represents a number greater than 1. 6.75 x 10-3 is the same as 0.00675 6.75 x 103 is the same as 6750

  20. MEASUREMENTSScientific Notation 5.289003 x 1011 3.400 x 10-12 Give the following in scientific notation (or write it out) with the appropriate significant figures. 1. 528900300000 = 2. 0.000000000003400 = 3. 0.23 = 4. 5.678 x 10-7 = 5. 9.8 x 104 = 2.3 x 10-1 0.0000005678 98000

  21. PRACTICE PROBLEMS 2.37 x 10-4 6.55 x 109 1.24 x 10-6 9.46 x 104 Show your work for the following questions on the back. Always give the correct significant figures. 1. Express each of the following numbers in scientific notation & 3 significant figures. A) 6545490087 _______ C) 0.0002368 _______ B) 0.000001243 _______ D) 94560 _______ 2. 0.00496 - 0.00298 = ________________ 3. (3.36-5.6) / (82.98 + 2.4) = ______________________ 4. 4.45 x 10- 23 / 8.345 x 10-53 = ________________ 5. [(26.7 x 10-8) (47 x 1013)]4 / (8.54 x 1017)1/2 = __________ 1.98 x 10 -3

  22. Dimensional Analysis Dimensional Analysis (also call unit analysis) is one method for solving math problems that involve measurements. The basic concept is to use the units associated with the measurement when determining the next step necessary to solve the problem. Always start with the given measurement then immediately follow the measurement with a set of parentheses. Keep in mind, try to ask yourself the following questions in order to help yourself determine what to do next. 1. Do I want that unit? If not, get rid of it by dividing by it if the unit is in the numerator, (if the unit is in the denominator, then multiply). 2. What do I want? Place the unit of interest in the opposite position in the parentheses. Numerator Denominator

  23. Dimensional Analysis 1. Calculate the number of weeks in 672 hours. 2. How many miles will a car travel in 3.00 hours at an average speed of 62.0 miles per hour?

  24. STUDY PROBLEMS A study of gemstones and dimensional analysis: The basic unit for gemstones is the carat. One carat is equal to 200 milligrams. • The Star of India sapphire (Al2O3, corundum) weighs 1.126 x 105 mg. What is the weight of the gemstone in carats? • The Cullinan Diamond was cut into nine major stones and 96 smaller brilliants. The total weight of the cut stones was 1063 carats, only 35.0% of the original weight! What weight (in milligrams) of the Cullinan Diamond was not turned into gemstones? • D.J. promised to bake 200 dozen cookies and deliver them to a bake sale. If each cookie weighs 3.5 ounces, how many grams will 200 dozen cookies weigh? 1 oz. is the same as 28.3495 g. 4. One box of envelopes contains 500 envelopes. A case of envelopes contains 28 boxes of envelopes and cost $123.49. What is the cost, in cents, of an envelope?

  25. MEASUREMENTS

  26. MEASUREMENTS • There are different types of measurements that can be made in the laboratory like mass, time, volume, and length. • These measurements can be made using either the metric system or the English system. The metric system is based on increments of 10. 1 base = 100 centibases “c” = centi 1 base = 1000 millibases “m” = milli 1 kbase = 1000 bases 1 base = 106 microbases “m” = micro k = kilo 1 base = 109 nanobases “n” = nano • The first step to understanding measurements is to learn the types, symbols, & units associated with these measurements.

  27. MEASUREMENTS • There are different types of measurements that can be made in the lab for length, mass, volume, temperature, area, time, heat and pressure.

  28. MEASUREMENTS • Putting it all together: Length (variable in a math equation = L ) Þsymbol for units:cm stands for centimeter, mm is millimeters, mm is micrometer, & nm is nanometer. Mass (variable “m”) Þsymbol for units:cg stands for centigram, mg is milligram, mg is microgram, & ng is nanogram. Volume (variable “V”) Þsymbol for units:cL stands for centiliter, mL is milliliter, mL is microliter, & nL is nanoliter. Note: One Liter is defined to be exactly 1000 cm3 1 mL = 0.001 L = 1 cm3

  29. SIMPLE MEASUREMENT CONVERSIONS How many meters are in 2608 centimeters? How many milliliters are in 2.96 liters? How many kilograms are equal to 0.3648 grams?

  30. MEASUREMENTS Since two different measuring systems exist, a scientist must be able to convert from one system to the other. CONVERSIONS Length: 1 in = 2.54 cm 1 mi = 1.61 km Mass: 1 lb = 454 g 1 kg = 2.2 lb Volume: 1 qt = 946 mL 1 L = 1.057 qt 4 qt = 1 gal 1 mL = 1 cm3 Temperature: °F = (1.8 °C) + 32 °C = (°F – 32) K = °C + 273.15 1.8

  31. STUDY PROBLEMS _____1. Water boils at 212 oF, what is the boiling point of water in oC? What is this temperature in Kelvin? _____2. Convert 25.0 ng to cg _____3. Convert 25.0 ks to min _____4. Convert 25.0 lb to mg _____5. Convert 25.0 ft3 to L 6. How many liters of gasoline will be used to drive 725 miles in a car that averages 27.8 miles per gallon? 7. What is the percent of salt in a saltwater solution if 246.99 cm3 of sodium chloride was added to 4.00 gal of pure DI water?

  32. STUDY PROBLEMS 8. If 2830.5 cg of a sample of haematite iron ore [iron (III) oxide, Fe203] were dissolved in concentrated hydrochloric acid and then the solution was diluted to 250 dm3. What is the mass of ore in pounds and what volume (in qt) of solution was made? 9. An instructor gives a sample of powered metal to each of four students (W, X, Y, & Z), and they weigh the samples on different balances. Their results for three trials are as follows. The true value is 1.921 x 10-2 lbs. A) Calculate the average mass for each data set with the correct significant figures. B) Which student was the most accurate in weighing? C) Which student was the most precise? D) Which student had the best combination for accuracy and precision?

  33. DENSITY

  34. Introduction to Density • Density is the measurement of the mass of an object per unit volume of that object. d = m / V • Density is usually measured in g/mL or g/cm3 for solids or liquids. • Volume may be measured in the lab using a graduated cylinder or calculated using: • Volume = length x width x height if a box or • V = pr2h if a cylinder. • Remember 1 mL = 1 cm3

  35. How to measure the density of a solid in the laboratory. • Obtain a clean graduated cylinder. • Fill the graduated cylinder with enough water to cover the object. Record the volume • Carefully place the object into the water filled graduated cylinder. • Record the new water level. • The volume of the object is the Vfinal – Vinitial.

  36. Study Problem on Density Record your final answer on this sheet and show all of your work on the back. Use scientific notation when necessary and always given your answers with the appropriate significant figures. • 1. Calculate the volume, in liters, of a box that is 346 mm long by 6.75 cm wide by 17.88 in high. • 2. Calculate the density of a liquid if 63.76 mL of the liquid has a mass of 3456. cg. • 3. What volume of mercury will have a mass of 4.0 kg if it has a density of 13.6 g/mL? •  4. What is the density (in g/mL) of a substance that weighs 9.272 lb. and occupies a volume of 12.48 qt. • 5. The mass of an irregular shaped object was measured to be 35.58 g. A student fills a 50.0 mL graduated cylinder with 23.0 mL of water, then places the object into the graduated cylinder. The new volume is 47.5 mL. Calculate the density of the object. • 6.  A piece of indium weighing 9.942 g is placed in 39.7 cm3 of ethyl alcohol (density = 0.798 g/mL) in a graduated cylinder. The alcohol level increases to 48.8 cm3. The density of indium based on this data is? Is the measurement accurate?

  37. STATES OF MATTER

  38. Three States of Matter

  39. Definitions

  40. The Atomic-Molecular Theory of Matter A “microscopic” view

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