Chemistry for Changing Times, 12E By Hill, Kolb, McCreary
Email: Richard@qwconsultancy.com
Chapter 1
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Chapter 1. Chemistry Overview Chapter 1 introduces the student to the concepts of science, and its ability to solve problems as well as its limitations. The student is then introduced to the concept of matter and energy and the physical means by which they are measured. Lecture Outline 1.1
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Science and Technology: The Roots of Knowledge Many students are confused about the differences between science and technology. Technology is the direct application of knowledge to solve problems. Science seeks an understanding of underlying principles. The Baconian Dream and the Carsonian Nightmare Bacon was one of the first to propose the experimental approach to science. Bacon also felt science could solve all problems. Rachel Carson was one of the first “environmentalists.” She pointed out the dangers of DDT and other chemicals damaging to the environment, indicating that the application of science was perhaps not all good. Science: Testable, Reproducible, Explanatory, Predictive, and Tentative The scientific method, from hypothesis to law, theory, or model; is best explained by example. The Limitations of Science The scientific method is dependent on the ability to control all aspects of a problem; this limits the method to “simple” systems. Science and Technology: Risks and Benefits Chemistry: Its Central Role Solving Society’s Problems: Scientific Research Research can be basic or applied. Chemistry: A Study of Matter and Its Changes Matter is anything that has mass. Mass is a measure of the amount of matter present. Weight is mass times gravitational attraction. Matter exhibits chemical and physical properties. Chemical properties tell us how matter will combine to form new and different substances (a chemical change). Physical properties are directly observable: color, state (solid, liquid, or gas), and texture are examples. An excellent demonstration is to burn a candle (chemical change) and boil water (physical change).
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Chapter 1
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Classification of Matter A. States of Matter • Solids, liquids, and gases. B. Substances and Mixtures • Substances have constant composition. • The composition of a mixture is variable. ▪ Homogeneous mixtures: appear the same throughout (milk, paint, and saltwater are examples). ▪ Heterogeneous mixtures: appear different throughout (pizza, raisin bread, and chocolate chip cookies are examples). C. Elements and Compounds • Substances are either elements or compounds. ▪ Elements: fundamental building blocks of all matter. (lead, silver, gold, carbon, and oxygen) ▪ Compounds: two or more elements chemically combined in fixed ratios. (water, ammonia and propane) D. Atoms and Molecules The Measurement of Matter Most students have been introduced to the metric system during their K-12 education; the SI system used in science is based on the metric system. The difference lies in the base units. • Base units: kilograms (kg) for mass, meter (m) for length, and seconds (s) for time. The four other base units are shown in Table 1.4. Density Density is the mass-to-volume ratio of matter. Energy: Heat and Temperature The SI unit of temperature is the kelvin (K); however, the temperature scale used in the chemistry laboratory is normally the Celsius scale. The Celsius scale was developed with the freezing and boiling points of pure water at a pressure of 1 atmosphere as the reference frames. (Students find the history of the development of the Fahrenheit scale much more interesting!) Critical Thinking
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Chapter 1
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Transparencies T-1 T-2 T-3 T-4 T-5 T-6 T-7 T-8 T-9
Inside cover Figure 1.1 Figure 1.4 Figure 1.5 Figure 1.6 Figure 1.7 Figure 1.9
Periodic Table The evaporation of water Risks of death (p. 10) Comparison of copper and sulfur Solids, liquids, and gases A scheme for classifying matter Comparison of metric and customary units Comparison of large and small objects (p. 20) A comparison of the Fahrenheit, Celsius, and Kelvin temperature scales
Demonstrations 1. Place samples of various elements (copper, sulfur, zinc, mercury, aluminum, carbon, etc.) in small stoppered bottles or flasks that can be passed around. 2. Ask students: “Which is heavier, lead or aluminum?” They will say “lead.” Then bring out two flasks, a 50 mL flask with about 10 mL of lead shot and a 500 mL flask about 3/4 full of aluminum pellets. Point out that the aluminum is actually heavier. But we need to compare samples of the same size (volume). Weight per unit volume is density. Lead has the greater density. 3. Compare a yardstick and a meter stick. A meter is slightly bigger than a yard. A sugar cube is approximately 1 mL. An ordinary paper clip weighs about 1 gram. A liter is a little bit (6%) larger than a quart. 4. Use Crispix and Raisin Bran to illustrate the difference between a compound and a mixture. Raisins and bran flakes are easily separated, and the ratio of raisins to flakes can vary. Crispix has a 1:1 ratio of corn and rice flakes, and they cannot be easily separated. 5. Ice floats on water because the density of ice is less than that of water. But alcohol has a much lower density. If you use alcohol instead of water, the ice will sink. Place a glass of water and a glass of alcohol side by side and add an ice cube to each one. 6. The “cartesian diver” is a popular demonstration about density. Completely fill a 2-liter plastic bottle with water, and put just enough water into a medicine dropper so that it can barely float. Put the dropper into the bottle, and cap the bottle. Squeeze and release the bottle to make the “diver” go up and down.
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Chapter 2
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Chapter 2. Atoms Overview Chapter 2 begins with the Greek concept of matter and moves chronologically forward through the development of the atomistic model of matter. Lecture Outline 2.1
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Atoms: The Greek Idea The Greek philosophers had two main ideas regarding matter: The atomistic view (Leucippus and Democritus) The continuous view (Aristotle) Lavoisier: The Law of Conservation of Mass Lavoisier was the first great experimentalist. More than anything, he introduced experimental chemistry to Western civilization. Through his experiments he formulated the law of conservation of mass which states: Matter is neither created nor destroyed during a chemical reaction. Proust: The Law of Definite Proportions The law of definite proportions is strong evidence for the atomistic nature of matter. If matter were continuous, any ratios by mass would be possible. John Dalton and the Atomic Theory of Matter John Dalton summarized the work of Lavoisier and Proust to form the famous atomic theory. Out of Chaos: The Periodic Table John Dalton established relative atomic masses using hydrogen as a base. Mendeleev arranged a table of elements according to increasing atomic weights placing elements with similar properties in the same column. This was the beginning of the modern periodic table! Atoms: Real and Relevant
Transparencies T-10 T-11 T-12
Figure 2.1 Figure 2.4 Figure 2.5 Figure 2.6
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Table 2.1
A sandy beach Mercuric oxide and conservation of mass Law of Definite Proportions and Berzelius’s experiment illustrating the law of definite proportions Electrolysis of water The laws of definite proportions and conservation mass interpreted in terms of Dalton’s atomic theory The law of multiple proportions. A carbon atom can combine with either one or two atoms of oxygen
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