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CHAPTER ONE

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  1. CHAPTER ONE The Foundations of Chemistry

  2. Why is Chemistry Important? Materials for our homes Components for computers and other electronic devices Fuel Body functions Cooking

  3. Some definitions / Vocabulary • Chemistry • Science that describes matter – its properties, the changes it undergoes, and the energy changes that accompany those processes • Matter • Anything that has mass and occupies space. (In other words: anything that has mass and volume) • Energy • The capacity to do work or transfer heat. • Types of energy • Kinetic and potential energy • Heat energy, light energy, chemical energy, mechanical energy

  4. Natural Laws • The Law of Conservation of Mass • During a chemical or physical change the mass of the system remains constant • The Law of Conservation of Energy • Energy cannot be created or destroyed in a chemical reaction or in a physical change. It can only be converted from one form to another. • The Law of Conservation of Matter and Energy • Read at home

  5. States of Matter • Liquid • Solid • Gas

  6. States of Matter Change States • heating • cooling Steam Water Ice

  7. Substances • Substance • matter all samples of which have identical composition and properties • Examples • water • sulfuric acid • Properties • physical properties – physical changes • chemical properties – chemical changes

  8. Physical Properties • Physical properties • changes of state • density, color, solubility • always involve only one substance • A substance cannot be broken down or purified by physical means!

  9. Mixtures • Mixture • a combination of two or more substances • can be separated by physical means • Homogeneous mixtures • have uniform properties throughout • examples: salt water; air • Heterogeneous mixtures • do not exhibit uniform properties throughout • examples: iron+sulfur; water+sand

  10. Chemical Properties • Chemical properties • chemical reactions • always involve changes in composition • always involve more than one substance • Examples • burning of methane • rusting of iron • oxidation of sugar

  11. Decomposition of Water hydrogen Element Element oxygen water Compound

  12. Compounds and Elements • Compounds • If a substance can be decomposed into simpler substances through chemical changes, it is called a compound • Elements • If a substance cannot be decomposed into simpler substances by chemical means, it is called an element

  13. Compounds and Elements • Important to remember • both compounds and elements are substances • a compound consists of 2 or more elements • Law of Definite Proportions • different samples of any pure compound contain the same elements in the same proportion by mass • Symbols of elements • found on the periodic chart (learn Table 1-2) • www.webelements.com

  14. Scientific Notation • Use it when dealing with very large or very small numbers: • 42,800,000. = • 0.00000005117 =

  15. Measurements in Chemistry QuantityUnitSymbol • length meter m • mass kilogram kg • time second s • current ampere A • temperature Kelvin K • amt. substance mole mol

  16. Metric Prefixes NameSymbolMultiplier • mega- M 106 • kilo- k 103 • deci- d 10-1 • centi- c 10-2 • milli- m 10-3 • micro-  10-6 • nano- n 10-9 • pico- p 10-12

  17. Metric Prefixes: Examples 1000 m = 0.008 s = 30,000,000 g = 0.07 L =

  18. Use of Numbers • Exact numbers • obtained from counting or by definition • 1 dozen = 12 things for example • Measured numbers • numbers obtained from measurements are not exact • every measurement involves an estimate

  19. Significant Figures • Significant figures • digits believed to be correct by the person making the measurement

  20. in my judgement! certain figures estimated figure Significant Figures • Side B: 13.6 mm >13.5 mm but <13.7 mm 13.6 mm

  21. estimated figure Significant Figures 13.6 mm • we always report only 1 estimated figure • the estimated figure is always the last one of the significant figures certain figures + significant figures

  22. Significant Figures - Rules • Exact numbers (defined quantities) have an unlimited number of significant figures. We do not apply the rules of significant figures to them. • Leading zeroes are never significant: • 0.000357 has three significant figures • Zeros between nonzero digits are always significant: • 20.034 1509 1.0000005

  23. Significant Figures - Rules Trailing zeros • Zeroes at the end of a number that contains a decimal point are always significant: • 35.7000 0.07200 40.0 41.0 • Zeroes at the end of a number that does not contain a decimal point may or may not be significant (use scientific notation to remove doubt): • 173,700 may have 4, 5, or 6 significant figures

  24. Significant Figures - Rules Addition/Subtraction Rule • The position of the first doubtful digit dictates the last digit retained in the sum or difference. Multiplication/Division Rule • In multiplication or division, an answer contains no more significant figures than the least number of significant figures used in the operation. • Study examples 1-1 & 1-2 in the book

  25. The Unit Factor Method • The basic idea of the method: • multiplication by unity (by 1) does not change the value of the expression • Principles: • construct unit factors from any two terms that describe identical quantity • the reciprocal of a unit factor is also a unit factor • Study examples 1-3 through 1-9 in the book

  26. The Unit Factor Method • 1 ft = 12 in Unit factors: • Example: Express 77.5 inches in feet 77.5 in = 77.5 in x = 6.46 ft • See Table 1-7 for various conversion factors

  27. 3 ft 1 yrd 12 in 1 ft 2.54 cm 1 in 10 mm 1 cm More examples • 9.32 yrd = ? mm 1. We use the following knowledge to build unit factors: 1 yrd = 3 ft 1 in = 2.54 cm 1 ft = 12 in 1 cm = 10 mm 2. Multiply 9.32 yrd by unit factors to get the value expressed in mm: 9.32 yrd x x x x = 8.52·103 mm

  28. Density • density = mass volume • tells us how heavy a unit volume of matter is • usually expressed as “g/ml” for liquids and solids and as “g/L” for gases • Table 1-8 lists densities of some common substances

  29. Density: Example • Example: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3. • Learn examples 1-11 through 1-13 in the book

  30. Specific Gravity • Sp. Gr. = d (substance) d (water) • tells us how much heavier or lighter a substance is compared to water: Sp. Gr. < 1 – lighter than water Sp. Gr. > 1 – heavier than water • specific gravity has no units – it is a dimensionless quantity • See example 1-14 in the book

  31. Specific Gravity: Example • Example 1-15: Battery acid is 40% sulfuric acid, H2SO4, and 60% water by mass. Its specific gravity is 1.31. Calculate the mass of pure H2SO4 in 100.0 mL of battery acid. • What do we know? 1. The mass percentage of H2SO4 and H2O in the sample of battery acid. 2. Specific gravity of battery acid. 3. Density of water (1.00 g/mL). • To find the mass of H2SO4, we need to know the mass of 100.0 mL of battery acid.

  32. Specific Gravity: Example Therefore,

  33. Heat and Temperature • Heat and Temperature are not the same thing: • Heat is a form of energy • T is a measure of the intensity of heat in a body • Heat always flows spontaneously from a hotter body to a colder body – never in the reverse direction Body 1 T1 Body 2 T2 Heat hotter T1 > T2 colder

  34. Temperature Scales • 3 common temperature scales  Fahrenheit  Celcius  Kelvin 0ºF – freezing (salt+H2O) 30ºF – freezing H2O 90ºF – human body 0ºC – freezing H2O 100ºC – boiling H2O 0 K – absolute zero 273.15 K – freezing H2O http://home.comcast.net/~igpl/Temperature.html

  35. MP BP • Fahrenheit 32 oF 212 oF • Celsius 0.0 oC 100 cC • Kelvin 273 K 373 K Temperature Scales & Water Melting (MP) and boiling (MP) points of water on different temperature scales

  36. Temperature Conversion degrees Kelvin degrees Celcius ? K = ?ºC + 273 ?ºC = ? K - 273 degrees Fahrenheit degrees Celcius ?ºF = (?ºC)·1.8 + 32 ?ºC = (?ºF – 32)/1.8 • Examples 1-16 & 1-17 in the book • http://www.lenntech.com/unit-conversion-calculator/temperature.htm

  37. Heat • Chemical and Physical changes: • evolution of heat (exothermic processes) • absorption of heat (endothermic processes) • Units of measurement: • joule (J) – SI units • calorie (cal) – conventional units • 1 cal = 4.184 J • A “large calorie” (1 large cal = 1000 cal = 1 kcal) is used to express the energy content of foods

  38. Specific Heat • The specific heat (Cp) of a substance: • the amount of heat (Q) required to raise the temperature of 1 g of the substance 1ºC (or 1 K) • Units of measurement:

  39. Specific Heat: Example 1 • Knowing specific heat, we can determine how much energy we need in order to raise the temperature of a substance by T = T2 – T1: • Calculate the amount of heat necessary to raise the temperature of 250 mL of water from 25 to 95ºC given the specific heat of water is 4.18 J·g-1 ·ºC-1. • What do we know? • the temperature change • the specific heat of water • the volume of water • the density of water

  40. Specific Heat: Example 1 • Examples 1-18 through 1-20 in the book

  41. Specific Heat: Example 2 • Given specific heats of two different substances, we can also calculate the heat transfer between them: • 0.350 L of water at 74.0ºC is poured into an aluminum pot at room temperature (25.0ºC). The mass of the pot is 200 g. What will be the equilibrium temperature of water after it transfers part of its heat energy to the pot? The specific heats of aluminum and water are 0.900 and 4.18 J·g-1 ·ºC-1, respectively. You might encounter this kind of problem at your first exam

  42. Specific Heat: Example 2 • What do we know? • the pot and water come to equilibrium, that is eventually they have the same temperature • the specific heat of aluminum and water • the mass of aluminum • the volume of water • the density of water • finally, the Law of conservation of energy which tells us that the amount of heat lost by water is the same as the amount of heat gained by the aluminum pot

  43. Specific Heat: Example 2 Let’s denote the final temperature as Tf. Then the changes in temperature for water and aluminum are: Note that we used the unit factor method to convert L to mL

  44. Specific Heat: Example 2 Solving this equation with respect to Tf, we obtain Tf = 68.6ºC • Try to solve the equation yourself and analyze why the answer is given with 3 significant figures

  45. Reading Assignment • Read Chapter 1 • Learn Key Terms (pp. 40-41) • Go through Chapter 2 notes available on the class web site • If you have time, read Chapter 2

  46. Homework Assignment Textbook problems (optional, Chp. 1): • 11, 13, 15, 18, 27, 29, 30, 32, 36, 41, 43, 47, 49, 57, 62, 68, 80 OWL: • Chapter 1 Exercises and Tutors – Optional • Introductory math problems and Chapter 1 Homework problems – Required (homework #1; due by 9/13)