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A Molecular Approach, 4e - Notes for Chapter (1).doc

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Category: Chemistry
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Chapter 1. Matter, Measurement, and Problem Solving Chapter 1. Matter, Measurement, and Problem Solving Chapter 1. Matter, Measurement, and Problem Solving Student Objectives 1.1 Atoms and Molecules Define atoms, molecules, and the science of chemistry. Represent simple molecules (carbon monoxide, carbon dioxide, water, hydrogen peroxide) using spheres as atoms. 1.2 The Scientific Approach to Knowledge Define and distinguish between a hypothesis, a scientific law, and a theory. Understand the role of experiments in testing hypotheses. State and understand the law of mass conservation as an example of scientific law. Understand that scientific theories are built from strong experimental evidence and that the term “theory” in science is used much differently than in pop culture. 1.3 The Classification of Matter Define matter and distinguish between the three main states of matter: solid, liquid, gas. Define and understand the difference between crystalline and amorphous solids. Define mixture, pure substance, element, compound, heterogeneous, and homogeneous. Differentiate between mixtures and pure substances; elements and compounds; and heterogeneous and homogeneous mixtures. Use the scheme on page 7 to classify matter. Define and understand the methods of separating mixtures: decantation, distillation, and filtration. 1.4 Physical and Chemical Changes and Physical and Chemical Properties Define, recognize, and understand the difference between physical and chemical changes. 1.5 Energy: A Fundamental Part of Physical and Chemical Change Define energy, work, kinetic energy, potential energy, and thermal energy. State and understand the law of conservation of energy. 1.6 The Units of Measurement Understand the importance of reporting correct units with measurements. Know the differences between the three most common sets of units: English system, metric system, and International System (SI). Know the SI base units for length, mass, time, and temperature. Know the three most common temperature scales (Fahrenheit, Celsius, and Kelvin), the freezing and boiling points of water on each scale, and the relationships between the scales. Calculate temperature conversions between each scale. Know and use the SI prefix multipliers for powers of ten. Know and calculate using the derived units of volume and density. 1.7 The Reliability of a Measurement Understand that all measurements have some degree of uncertainty and that the last digit in a measurement is estimated. Know how to determine the number of significant figures in a measurement using a set of rules. Know how to determine the number of significant figures after calculations. Distinguish between accuracy and precision. 1.8 Solving Chemical Problems Understand dimensional analysis and know how to use conversion factors. Understand the problem-solving strategy: sort, strategize, solve, and check. Convert from one unit to another. Make order-of-magnitude estimations without using a calculator. Rearrange algebraic equations to solve for unknown variables. Section Summaries Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples Teaching Tips Suggestions and Examples Misconceptions and Pitfalls Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 1.1 Atoms and Molecules Definitions of atoms, molecules Composition of water and hydrogen peroxide Definition of chemistry Intro figure: Disneyland ride showing water molecules unnumbered figures: models of H2O and H2O2 and structure of graphite and diamond 1.2 The Scientific Approach to Knowledge Definitions of hypothesis, falsifiable, experiments, scientific law, theory Scientific method: Observations and experiments lead to hypotheses. More experiments may lead to a law and a theory. A theory explains observations and laws. unnumbered figure: painting of Antoine Lavoisier Figure 1.1 The Scientific Method The Nature of Science: Thomas S. Kuhn and Scientific Revolutions Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 1.1 Atoms and Molecules Chemistry involves a great deal of what can't be seen directly, requiring representations and models. The intro figure shows hemoglobin, but the actual molecule is not a green and blue ribbon. Chemists look at microscopic, macroscopic, and symbolic representations of atoms and molecules interchangeably. If you say “water”, you might mean the formula H2O or a molecular model or a large collection of molecules (e.g., a glass of water). Students need help recognizing which representation to think about when a chemical name is used. 1.2 The Scientific Approach to Knowledge Experiments test ideas. They are designed to support a hypothesis or to disprove it. Good scientific hypotheses must be testable or falsifiable. Theories are developed only through considerable evidence and understanding, even though theories often are cited in popular culture as unproven or untested. Figure 1.1 shows how the scientific method is cyclic and allows for the refining of ideas. Conceptual Connection 1.1 Laws and Theories The box about Thomas Kuhn can help to clear misconceptions of science being completely objective and immutable. Theories are not as easily dismissible as pop culture suggests. Scientific knowledge constantly evolves as new information and evidence are gathered. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 1.3 The Classification of Matter States of matter: their definitions and some of their characteristics gas liquid solid crystalline amorphous Classification of Matter pure substance element compound mixture heterogeneous homogeneous Separating mixtures decantation distillation filtration Figure 1.2 Crystalline Solid unnumbered figure: illustrations of solid, liquid, and gas phases Figure 1.3 The Compressibility of Gases unnumbered figure: classification of matter Figure 1.4 Separating Substances by Distillation Figure 1.5 Separating Substances by Filtration 1.4 Physical and Chemical Changes and Physical and Chemical Properties Differences between physical and chemical changes Examples and classifying changes Figure 1.6 Boiling, a Physical Change Figure 1.7 Rusting, a Chemical Change Figure 1.8 Physical and Chemical Changes Example 1.1 Physical and Chemical Changes and Properties 1.5 Energy: A Fundamental Part of Physical and Chemical Change Definitions of work and energy Classification and types of energy kinetic thermal potential Definition and examples of the law of conservation of energy unnumbered figure: illustration of work (physical definition) Figure 1.9 Energy Conversions Figure 1.10 Using Chemical Energy to Do Work Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 1.3 The Classification of Matter Properties of matter define its state: gas, liquid, or solid. Temperature is one example, and everyone recognizes steam, water, and ice. Ask for additional examples such as dry ice or liquid nitrogen. Compressibility is a property that differentiates especially gases from liquids and solids. The thickened glass at the bottoms of old windows helps students appreciate the amorphous nature of glass. Classifying additional examples of matter, e.g. mayonnaise, Jell-O, and milk, according to the scheme demonstrates some of the challenges. Students are likely to have varying personal experience with distillation and filtration. Kitchen analogies may be useful: steam condenses on the inside of a pot lid; macaroni and water are poured into a colander; wine is often decanted. The differences between the space-filling models from Section 1.1 and the ball-and-stick model of diamond may be missed by some students. Students may not have experience with elemental forms other than diamond and charcoal. 1.4 Physical and Chemical Changes and Physical and Chemical Properties Conceptual Connection 1.3 Chemical and Physical Changes Boiling (especially) does not change a substance’s chemical identity. Confront the confusion that can occur when a physical change accompanies a chemical one: burning liquid gasoline produces gases. (physical or chemical or both?) 1.5 Energy: A Fundamental Part of Physical and Chemical Change The examples of work being done by a person moving a box and chemical energy ultimately moving the car are consistent and simple. Additional examples using gravitation (very familiar) are straightforward. Several examples are cited for the law of conservation of energy; ask students to name and describe other forms of energy (solar, mechanical, chemical, electrical) and devices that convert between these forms. Work is a form of energy and thus has the same units as energy. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 1.6 The Units of Measurement Loss of Mars Climate Orbiter because of inconsistent units Systems of measurement and units English system metric system International System (SI) SI base units length: meter mass: kilogram time: second temperature: Kelvin Temperature scales and conversions Fahrenheit to Celsius and vice versa Celsius to Kelvin and vice versa Derived units volume (cubic meter, cubic centimeter, liter, milliliter) density, mass per unit volume (g/mL, g/cm3) unnumbered figure: Mars Climate Orbiter unnumbered figures: heights in meters of Empire State Building and basketball player Table 1.1 SI Base Units unnumbered figure: electronic balance Figure 1.11 Comparison of the Fahrenheit, Celsius, and Kelvin Temperature Scales unnumbered figure: The Celsius Temperature Scale Example 1.2 Converting between Temperature Scales Table 1.2 SI Prefix Multipliers Figure 1.12 The Relationship between Length and Volume Table 1.3 Some Common Units and Their Equivalents Table 1.4 The Density of Some Common Substances at 20°C Example 1.3 Calculating Density Chemistry and Medicine: Bone Density 1.7 The Reliability of a Measurement Significance and reporting of numerical values estimating measurements Counting significant figures or digits nonzero digits interior zeroes leading zeroes trailing zeroes exact numbers Significant figures in calculations multiplication and division (fewest significant figures) addition and subtraction (fewest decimal places) rounding (best only after the final step) Precision vs. accuracy Scientific integrity and data reporting unnumbered figures: CO concentration in L.A. county; two tables with different significant figures for the data Figure 1.13 Estimation in Weighing Example 1.4 Reporting the Correct Number of Digits Example 1.5 Determining the Number of Significant Figures in a Number Example 1.6 Significant Figures in Calculations unnumbered figure: accuracy and precision Chemistry in Your Day: Integrity in Data Gathering Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 1.6 The Units of Measurement Students are amazed and horrified that NASA could lose an expensive spacecraft because of inconsistent units. Metric and SI units are unfamiliar to most Americans. That a nickel has a mass of 5 g and that a yard is nearly as long as a meter gives a good frame of reference. Conceptual Connection 1.4 The Mass of a Gas The practical examples of different temperatures on the Celsius scale (unnumbered figure) provide practical reference points. Several of the large SI unit prefixes (mega, giga, tera) are already familiar from memory capacity in computers. Conceptual Connection 1.6 Density The Chemistry and Medicine box on bone density provides an open-ended conceptual question about designing an experiment to measure bone density; this may be good for a brief in-class discussion. A common misconception is that 100 cm3 is equal to 1 m3. Some students initially are confused that density can be used as a conversion factor even when the units are inverted. 1.7 The Reliability of a Measurement Use a 400-mL beaker and a 100-mL graduated cylinder to measure quantities of water. Make the point about the importance of estimating measurements. Add the quantities of water together and ask the students to calculate the final volume...to the correct precision. Two tables present air quality data (with different precision) that might appear in a newspaper or other publication. Initiate a discussion of the certainty of digits in reported data. Water-quality standards have evolved substantially since the advent of instrumental methods for quantitative analysis. Ask the question: Does zero mean that a particular analyte is not present? The number on a calculator display requires interpretation; only the user knows the certainty of the values entered. A discussion about why integrity in data reporting is particularly important in science is appropriate. It should point out that scientists report how they did the experiments so others can try to repeat and verify the work. Use recent examples from the media. Students presume that calculators are flawless but forget that calculators do only what the user dictates. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 1.8 Solving Chemical Problems Converting from one unit to another dimensional analysis multiple approaches to any problem General problem-solving strategy sort strategize solve check Calculations using units raised to a power Order-of-magnitude estimations Using equations Example 1.7 Unit Conversion Example 1.8 Unit Conversion Example 1.9 Unit Conversions Involving Units Raised to a Power Example 1.10 Density as a Conversion Factor Example 1.11 Problems with Equations Example 1.12 Problems with Equations Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 1.8 Solving Chemical Problems General chemistry classes at most schools have students with a wide range of math skills. A quick review of algebra may be useful. Emphasize that watching an instructor work problems is not nearly as effective as working those same problems on one’s own. Give students time to work a problem or two in class; allow them to work in small groups. Emphasize the good practice of writing units and keeping track of units in every calculation. Simple dimensional analysis prevents many headaches throughout the year of general chemistry. Promote estimation as part of the problem-solving model. Tell the students to ask themselves, “Does this answer make sense?” Reduce the reliance on blindly entering numbers into a calculator and transcribing whatever answer comes up. Cognitive load theory says that a person can remember 7–9 items in short-term memory. A problem loaded with unit conversions, spurious facts, and many steps does not test a person’s understanding of an underlying idea or concept. It becomes a measure of cognitive ability outside the realm of chemistry. Students often want to follow one particular “recipe” to solve one particular kind of problem. Procedure for Solving Unit Conversion Problems Additional Problem (Example 1.7 Unit Conversion) Convert 1.76 miles to meters. Sort Begin by sorting the information in the problem into Given and Find. Given 1.76 mi Find m Strategize Devise a conceptual plan for the problem. Begin with the given quantity and symbolize each conversion step with an arrow. Below each arrow, write the appropriate conversion factor for that step. Focus on the units. The conceptual plan should end at the find quantity and its units. In these examples, the other information needed consists of relationships between the various units as shown. Conceptual Plan mi km m Relationships Used 1 km = 0.6214 mi 1 km = 1000 m (These conversion factors are from Tables 1.2 and 1.3.) Solve Follow the conceptual plan. Begin with the given quantity and its units. Multiply by the appropriate conversion factor(s), cancelling units, to arrive at the find quantity. Round the answer to the correct number of significant figures by following the rules in Section 1.7. Remember that exact conversion factors do not limit significant figures. Solution 2832.31 m = 2830 m Check Check your answer. Are the units correct? Does the answer make physical sense? The units (m) are correct. The magnitude of the answer (2830) makes physical sense since a meter is a much smaller unit than a mile. Additional Problem for Unit Conversion Involving Units Raised to a Power (Example 1.9) Calculate the number of cubic meters of concrete necessary to support a deck if each of 14 concrete piers require 4750 cubic inches. Sort Begin by sorting the information in the problem into Given and Find. Given 14 piers, 4750 in3 Find m3 Strategize Write a conceptual plan for the problem. Begin with the given information and devise a path to the information that you are asked to find. Notice that for cubic units, the conversion factors must be cubed. Conceptual Plan piers in3 m3 14 piers Relationships Used 1 m = 39.37 in (Conversion factor from Table 1.3) 1 pier = 4750 in3 (Given) Solve Follow the conceptual plan to solve the problem. Round the answer to three significant figures to reflect the three significant figures in the least precisely known quantity (4750). These conversion factors are all exact and therefore do not limit the number of significant figures. Solution 1.0897 m3 = 1.09 m3 Check The units of the answer are correct and the magnitude makes sense. The unit meters is larger than inches, so cubic meters are much larger than cubic inches. Additional Problem for Density as a Conversion Factor (Example 1.10) An experimental automobile has a 100.0 liter fuel tank filled with ethanol. How many pounds does the fuel add to the mass of the car? Sort Begin by sorting the information in the problem into Given and Find. Given 100.0 L Find lb Strategize Devise a conceptual plan by beginning with the given quantity, in this case the volume in liters (L). The overall goal of this problem is to find the mass. You can convert between volume and mass using density (g/cm3). However, you must first convert the volume to cm3. Once you have converted the volume to cm3, use the density to convert to g. Finally, convert g to lb. Conceptual Plan L mL cm3 g lb Relationships Used 1000 mL = 1 L 1 mL = 1 cm3 d (ethanol) = 0.789 g/cm3 1 lb = 453.59 g (These conversion factors are from Tables 1.2, 1.3 & 1.4.) Solve Follow the conceptual plan to solve the problem. Round the answer to three significant figures to reflect the three significant figures in the density. Solution 173.94 lb = 174 lb Check The units of the answer (lb) are correct. The magnitude of the answer (174) makes physical sense since a liter of water has a mass of 1 kilogram or about 2.2 pounds; 100 liters of water is about 220 lbs. Ethanol has a lower density than water (about 80% or 8/10). Additional Problem for Solving Problems Involving Equations (Example 1.12) What is the mass in grams of an ice cube that is 1.1 inches per side? Sort Begin by sorting the information in the problem into Given and Find. Given l = 1.1 in Find g Strategize Write a conceptual plan for the problem. Focus on the equation(s). The conceptual plan shows how the equation takes you from the given quantity (or quantities) to the find quantity. The conceptual plan may have several parts, involving other equations or required conversions. In these problems, you must use the geometrical relationships given in the problem as well as the definition of density. Conceptual Plan l V V = l 3 in3 cm3 g Relationships Used V = l 3 [volume of a cube with a length of l] 2.54 cm = 1 in d (ice) = 0.917 g/cm3 (These conversion factors are from Tables 1.3 and 1.4.) Solve Follow the conceptual plan. Solve the equation(s) for the find quantity. Gather each of the quantities that must go into the equation in the correct units. (Convert to the correct units if necessary.) Substitute the numerical values and their units into the equation(s) and compute the answer. Round the answer to the correct number of significant figures. Solution 20.0008 g = 20. g Check The units (g) are correct. The magnitude of the answer (20.) seems to make physical sense. 2 Copyright © 2017 by Education, Inc. 3 Copyright © 2017 by Education, Inc.

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