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I. The Study of Chemistry

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  1. I. The Study of Chemistry A. The Molecular Perspective of Chemistry • Studying the properties and behavior of matter • the physical material of the universe • It is anything that has mass and occupies space • It is comprised of combinations of only about 100very basic substance called elements. • Provides a background to understanding the properties of matter in terms of atoms • the almost infinitesimally small building blocks of matter • Atoms can combine to form molecules (chemical combination of atoms) • First Assignment: Learn the names and symbol for the first 48 elements by Friday.

  2. Oxygen Ethanol Water Aspirin Carbon Dioxide Ethylene glycol

  3. Why Study Chemistry? • It provides an important understanding of our world and how it works • Improvement of health care • Conservation of natural resources • Protection of the environment • Increased food production • Development of new materials • It is, by its very nature the central science • Astronomy, atmospheric science, biology, geology, environmental science, medicine, physics, material science, and polymers • The language of chemistry is a universal scientific language

  4. How Do I Study Chemistry? • It takes lots of practice--homework, reviewing notes, reading the text. • It is different than some other disciplines • MICHELANGELO Buonarroti, Italian painter, sculptor and architect (1475-1564). If a block of marble were at the front of this room I suspect we would select Michelangelo to teach us this art form. • Antoine Lavoisier (1743-1794:guillotined) is called the Father of Modern Chemistry but he thinks water’s formula is HO. He knows of about a dozen elements, nothing about polymers, nuclear chemistry, ceramics, etc. • You must study differently for chemistry--study nearly every day. The best predictor of your final grade is your grade on the first exam.

  5. B. Classifications of Matter • States of Matter • a gas, a liquid, or a solid • states of matter differ in some of their simple observable properties • gases (vapors) have no fixed volume or shape. They can be compressed to occupy a smaller volume or allowed to expand to occupy a larger volume • liquids have a distinct volume independent of the container that they occupy. They assume the shape of the portion of the container they occupy • solids have a definite shape and a definite volume

  6. Pure Substances • Most forms of matter are not chemically pure • air • gasoline • Sidewalks • Pure substances have distinct properties and a composition that does not vary from sample to sample • All substances are elements or compounds • elements cannot be decomposed into simpler substances • carbon, helium, iron, oxygen, chlorine, etc. • compounds are composed of two or more elements • Water (H2O) is composed of two elements, hydrogen and oxygen • Mixtures are combinations of two or more substances in which each substance retains its own chemical identity

  7. Atoms of an element Molecules of an element Mixture of elementsand a compound Molecules of a compound

  8. Elements • At the present time 114 elements are known • Elements vary widely in their abundance

  9. The symbol for each element consist of one or two letters, with thefirst letter capitalized. You will need to know the symbols and names for the first 100 elements

  10. Compounds • Most elements interact with other elements to form compounds • Hydrogen burns in oxygen to form water (one O and two H atoms)

  11. The observation that the elemental composition of a pure compound is always the same is known as the law of constant composition (law of definite proportions) • Mixtures • Most of the matter we encounter consists of mixtures of different substances • Each substance in a mixture retains its own chemical identity and properties • Mixture composition can vary • a cup of sweetened coffee. chocolate chip cookies. water found in nature, rocks, wood, cement, steel, etc. • Some mixtures are uniform throughout • Homogeneous mixtures or solutions • Some mixtures are not uniform throughout • Heterogeneous mixtures

  12. C. Properties of Matter • Every substance has a unique set of properties • characteristics that allow us to recognize and distinguish one substance from another • Properties of matter can categorized as either physical or chemical • Physical properties can be measured without changing the identity and composition of the substance • Color, odor, density, melting point, boiling point, hardness, etc. • Chemical properties describe the way a substance may change or react to form other substances • A common chemical property is flammability

  13. Some properties, such as temperature, melting point, and density, do not depend on the amount of sample being examined — intensive properties. • used to identify substances • Some properties, such as mass and volume, do depend on the amount of sample being examined — extensive properties.

  14. Physical and Chemical Changes • During physical changes a substances changes its physical appearance, but not its composition • ice melting to become water • water evaporating to become steam • All state changes are physical changes • During chemical changes (chemical reactions) a substances is changed into a chemically different substance • propane burning to form carbon dioxide and water • scrambling an egg • a change in state will not revert the substance back to its original form

  15. Separation of Mixtures • Mixtures can be separated into their constituent components • mixture components retain their own properties • Take advantage of the differences in the properties Heterogeneous Mixtures • visual differences • magnetic differences • state differences Homogeneous Mixtures • Boiling point difference • Polarity differences

  16. D. Units of Measurement • Many properties of matter are quantitative; that it, they are associated with numbers To say that the length of a pencil is 17.5 is meaningless • The units used for scientific measurements are those of the Metric System

  17. 1. SI Units • 1960 - international agreement specifying a particular choice of seven metric units for scientific measurements • SI = Systéme International d’Unités

  18. Prefixes are used to indicate decimal factions ormultiples of various units • Exponential notation is used to avoid ambiguity with regard to value certainty, see section 1.8. Learn these prefixes in Table 1.3.

  19. 2. Length and Mass • SI base unit of length is the meter (m) • 1 m = 100 cm = 39.37 inches (slightly longer than a yard) • Mass is a measure of the amount of material in an object. • Mass is different from weight, which depends upon gravity • SI base unit of mass is the kilogram (kg) • This base unit is unusual because it uses a prefix, kilo-, instead of the word gram alone. • Other units of mass are obtained by adding prefixes to the word gram

  20. 3. Temperature • We sense temperature as a measure of hotness and coldness • Temperature determines the direction of heat flow • Heat always flows spontaneously from a substance at higher temperature to one at lower temperature. • The temperature scales commonly employed in scientific studies are the Celsius and Kelvin scales • The Celsius scale, the everyday scale of temperature in most countries throughout the world, was originally based on the assignment of 0˚C to the freezing poing of water and 100˚C to its boiling point at sea level.

  21. The Kelvin Scale is the SI temperature scale and the SI unit of temperature is the Kelvin (K). • Zero on this scale is - 273.15˚C, once believed to be the lowest attainable temperature. Because of this, 0K is known as Absolute Zero. • Both the Celsius and Kelvin scales have equal-sized units — that is, a kelvin is the same size as a degree Celsius. K = ˚C + 273.15 • The freezing point of water, 0˚C, is 273.15 K • Notice that the degree symbol (˚) is not used with temperatures on the Kelvin scale.

  22. The common temperature scale in the US is the Fahrenheit scale • Freezing point of water = 32˚F • Boiling point of water = 212˚F ˚C = 5/9 (˚F - 32) or ˚F = 9/5 (˚C) + 32

  23. 4. Volume • The volume of a cube is given by its length cubed, (length)3 • The basic SI unit of volume is the cubic meter • Another common unitof volume is the liter (L),which equals a cubicdecimeter, dm3. • It is slightly larger than aquart • There are 1000 milliliters (mL)in a liter • Each milliliter is the same volume as acubic centimeter

  24. Devices Used Most Frequently in Chemistry to Measure Volume

  25. 5. Density • Widely used to characterize substances • Amount of mass in a unit volume of the substance • Densities of solids and liquids are commonly expressed in grams per cubic centimeter (g/cm3) or grams per milliliter (g/mL) • The density of water is 1.00 g/ml • Most substances change volume when heated or cooled — Densities are temperature dependent • Assume 25˚C unless otherwise noted

  26. E. Uncertainty in Measurement There are two kinds of numbers in scientific work: Exact numbers (those whose values are known exactly) 12 eggs in a dozen, exactly 1000 g in a kilogram, and exactly 2.54 cm in an inch and inexact numbers (those whose values have some uncertainty) numbers obtained by measurement Uncertainties always exist for in measured quantities

  27. 1. Precision and Accuracy • Precision is a measure of how closely individual measurements agree with one another • Accuracy refers to how closely individual measurements agree with the correct, or “true” value.

  28. 2. Significant Figures • Precision of a measured number is indicated using the concept of significant figures. • Those digits in a measured number (or result of a calculation with measured numbers) that include all certain digits plus a final one have some uncertainty. • Three measurements (9.12, 9.11, and 9.13 cm) • Avg = 9.12 First two digits (9.1) are certain, the next digit is estimated, so it has some uncertainty.

  29. The greater the number of significant figures, the greater the certainty implied for the measurement • In any measurement that is properly reported, all nonzero digits are significant. Zeros can be used either as part of the measured value or merely to locate the decimal point. • Zeros between nonzero digits are always significant • Zeros starting a number are never significant • Zeros ending a number to the right of the decimal point are always significant • Zeros ending a number to the left of the decimal point may or may not be significant • Exponential notation is the solution

  30. Significant Figures in Calculations Suppose that 0.0634 g of a compound will dissolve in 25.31 g of water. Howmany grams will dissolve in 100 g ofwater?

  31. Significant Figure Calculation Rules Multiplication and Division. When multiplying or dividing measuredquantities, give as many significant figures in the answer as there arein the measurement with the least number of significant figures. If the leftmost digit to be dropped is ≥ 5, round the last significantfigure up, otherwise simply drop the nonsignificant digits.

  32. Significant Figure Calculation Rules Addition and Subtraction. When adding or subtracting measuredquantities, give the same number of decimal places in the answer asthere are in the measurement with the least number of significant figures. If the leftmost digit to be dropped is ≥ 5, round the last significantfigure up, otherwise simply drop the nonsignificant digits.

  33. Dimensional Analysis A.This process is called1. Dimensional analysis (we will use this name)2. Factor-label method, Unit conversion method, or the Unit Factor Method • Dimensional analysis is a problem solving aid used to help ensure that the solutions to problems yield the proper units. • It provides a systematic way of solving any numerical problems and of checking solutions for possible errors. • The key to using dimensional analysis is the correct use of conversion factors to change one unit into another. • This method will work for many chemical calculations. When it is not convenient you must still cancel the units to be certain your answer has the proper dimensions.

  34. 1. A conversion factor is a fraction whose numerator and denominator are the same quantity expressed in different units. 2. Two important mathematical realities • Doing the same thing to both sides of an equation does not change the relationship • Multiplying one does not change the anything

  35. 3. Since conversion factors are the number one in a different form, multiplying a measurement and its units by any number of conversion factors changes the value and the units but not the reality of the measurement itself. How many seconds in a century?

  36. B. Conversion Factors: • Constructed from any two terms that are equal to each other. If two quantities are equal and if one is then divided by the other, the quotient equals one and it is called a Unit Factor or a Conversion Factor. Remember: any quantity can be multiplied by one without changing the result. • Examples of Conversion Factors, Memorize these. • a. 1.00 inch = 2.54 cm Length • b. 454 g = 1.00 lb Mass • c. 1.00 L = 1.057 qt Volume • d. 1.00 mL = 1.00 cm3 = 1.00 cc • (learn these first four conversions): • e. 1.00 mol Mg = 24.3 g of Mg • f. Many others that you will learn.

  37. Volume, like other concepts that will be encountered throughout the semester, requires the use of relationships that necessitate the raising of numbers to a power, cubing in this specific case • It is imperative to remember to raise both the number and the units to the appropriate power and not just the units 1 in = 2.54 cm 1 in3 ≠ 2.54 cm3 1 in3 = (1 in)3 = (2.54 cm)3 = 16.39 cm3

  38. C. Recipe for Dimensional Analysis 1. On the far right, write down the units of the answer. 2. Analyze the information given or known (the units) to select the proper starting quantity. This information may be from the problem, the periodic chart, some physical law, or something you learned. 3. Analyze the dimensions (units) of the answer and the dimensions (units) of the starting quantity to determine the proper unit factors (conversion factors) to convert the units given into the units of the answer. 4. This may require several unit factors. 5. Cancel the units and do the numerical calculation.

  39. Examples: 1. How many inches in 2.57 ft? DO ANS: 2. How many nm in 3.72 yds? DO ANS:

  40. 3. A major contributor to global environmental pollution is from coal burning power plants. Carbon dioxide, a greenhouse gas, is the major product and all coal contains various amounts of sulfur that eventually contributes to acid rain. A typical coal burning power plant burns 2500 tons of coal per day, with a sulfur content of 3%, and the density of solid bituminous coal is 1346 kg/m3 (solid anthracite coal has a density of 1506 kg/m3). a. What volume, in cubit feet, of bituminous coal is burned by a typical coal burning power plant each day? ANS:

  41. How many grams of sulfur are burned each day? ANS: If a railroad car holds 100 tons of coal how many car loads are burned each day? ANS:

  42. 4. Determine the volume in ft3 of a piece of Pb that has a mass of 13.4 kg, if the density is11.2 g/cm3. DO ANS: 5. A gallon of milk weighs 8.00 lbs. a. What is the mass of one pint in grams? DO ANS: b. What is the density of milk in g/mL? DO ANS:

  43. Douglas Isbell/Don Savage Headquarters, Washington, DC Nov. 10, 1999 (Phone: 202/358-1547) Embargoed until 2 p.m. EST RELEASE: 99-134 MARS CLIMATE ORBITER FAILURE BOARD RELEASES REPORT, NUMEROUS NASA ACTIONS UNDERWAY IN RESPONSE Wide-ranging managerial and technical actions are underway at NASA's Jet Propulsion Laboratory, Pasadena, CA, in response to the loss of the $125 million Mars Climate Orbiter and the initial findings of the mission failure investigation board, whose first report was released today. "The 'root cause' of the loss of the spacecraft was the failed translation of English units into metric units in a segment of ground-based, navigation-related mission software, as NASA has previously announced," said Arthur Stephenson, chairman of the Mars Climate Orbiter Mission Failure Investigation Board. "The failure review board has identified other significant factors that allowed this error to be born, and then let it linger and propagate to the point where it resulted in a major error in our understanding of the spacecraft's path as it approached Mars. WEB: http://mars.jpl.nasa.gov/msp98/news/mco991110.html