Type:
Solution Manual
General Organic and Biological Resource: Chemistry An Integrated Approach Edition:
4th Edition
Author(s):
Kenneth W. Raymond
Chapter 1 Science and Measurements Outline Notes I. Scientific Method (Section 1.1) A. Science and Chemistry Overview 1. Science greatly impacts our lives (i.e. DNA discovery) 2. Science covers a wide range of subjects including chemistry, biology, biochemistry, geology, astronomy, physics, etc. 3. Chemistry is the study of matter and its changes 4. Chemistry lays the groundwork for many other fields of science 5. Examine Figure 1.1 B. Scientific Method 1. Observation is the first step where information is gathered (i.e. Newton’s apple) 2. Laws are statements that describe things consistently and reproducibly observed (i.e. Newton’s Law of Gravity) 3. A hypothesis is a tentative and testable explanation based on observations or known facts (i.e. a clinician diagnosis a patient’s problem) 4. Experiments test the hypothesis (i.e. medical tests) a. Designed to be directly related to the problem or question at hand b. If the experiment invalidates the hypothesis then the hypothesis must be revised c. If a hypothesis survives repeated testing in can become part of a theory 5. A theory is an experimentally tested explanation of an observed behavior a. Must be consistent with existing experimental evidence b. Must accurately predict the results of future experiments c. Must explain future observations 6. Examine Figure 1.2 a. Scientist may not follow this order and they may not use all the steps b. Creativity and sufficient knowledge are required for the process c. Theories are continually reevaluated and possibly revised *See Sample Problem 1.1 and Practice Problem 1.1 DID YOU KNOW? Diabetes HEALTH LINK Science and Medicine (Examine Figure 1.4) II. Matter and Energy (Section 1.2) A. Matter Properties 1. Matter is anything that has mass and occupies space (i.e. body, air, book) 2. Physical properties are characteristics that can be determined without changing the chemical composition of matter (what it’s made of) a. Odor, color, malleability, taste, melting point, state/phases 3. Three physical states/phases a. Solids have fixed shapes and volumes b. Liquids have variable shapes and fixed volumes c. Gases have variable shapes and volumes 4. Examine Figure 1.5 5. Physical state is partially determined by the interactions of particles (atoms, molecules, etc.) 6. Examine Figure 1.6 B. Physical Changes 1. A physical change does not alter the chemical composition of matter (ice to water to steam) 1-1
2. Examine Figure 1.7 *See Sample Problem 1.2 and Practice Problem 1.2 C. Energy and Matter 1. Anytime matter is changed, work has been done 2. Energy is the ability to do work or transfer heat 3. Potential energy is stored energy and kinetic energy is the energy of motion 4. Potential energy (i.e. gasoline in a car) can be converted into kinetic energy (i.e. movement of a car’s piston), which produces work and heat *See Sample Problem 1.3 and Practice Problem 1.3 D. Heat and Changes 1. Heat plays a role in physical changes (i.e. heat added to water causing the motion of the particles/kinetic energy to increase turning it into steam) 2. Heat of fusion is energy put in during the melting process (solid to liquid) 3. Heat of vaporization is the energy put in during the boiling/vaporization process (liquid to gas) 4. These processes can be reversed with the removal of heat (i.e. condensation and freezing) 5. Examine Figure 1.8 6. Under certain conditions, some substances skip the liquid phase a. Sublimation is the conversion of a solid to a gas (i.e. dry ice) b. Deposition is the conversion of a gas to solid 7. Examine Figure 1.9 *See Sample Problem 1.4 and Practice Problem 1.4 III. Units of Measurement (Section 1.3) A. Measurement overview 1. Measurements consist of a number and a unit (i.e. 3 minutes, 3 hours, 3 miles) 2. The three systems/units of measurement are the metric units (used worldwide), the English units (common in US) and the SI units (international derived from metric system) 3. Examine Figure 1.10 B. Mass 1. Measure of the amount of matter in a sample 2. See Table 1.1 and 1.2 for common units and relationships 3. Weight (different from mass) is a measure of the force of gravity on an object’s mass C. Length 1. Measure of the distance between two points 2. One meter is equal to 3.281 feet therefore it is slightly longer than a yard D. Volume 1. Measure of the space occupied by an object 2. One liter is 1/1000 that of cubic meter 3. Examine Figure 1.11 DID YOU KNOW? Oil Spill E. Temperature 1. Three scales/units are degrees Fahrenheit (°F), degrees Celsius (°C), and Kelvin (K) a. °F is the English unit, water freezes at 32°F and boils at 212°F b. °C is the metric unit, water freezes at 0°C and boils at 100°C c. K is the SI unit based on absolute zero (coldest possible temperature: 0K/-273.15°C) 2. Examine Figure 1.12 F. Energy 1. The calorie (cal) is the metric and English unit for energy 2. 1 cal is the amount of energy required to raise 1 gram of water 1°C 3. The joule (J) is the SI unit; 4.184J = 1 cal 4. The food calorie (Cal) is used commonly; 1 Cal = 1000cal *See Sample Problem 1.5 and Practice Problem 1.5 1-2
IV. Scientific Notation, SI and Metric Prefixes (Section 1.4) A. Scientific Notation 1. Condenses very large or very small numbers 2. Expresses a number between 1 and 10 multiplied by a power of ten a. Positive exponents indicate how many times to multiply by 10 (shift decimal left) b. Negative exponents indicate how many times to divide by 10 (shift decimal right) 3. See Table 1.3 for scientific notation examples 4. Allows numbers to be compared without counting zeros *See Sample Problem 1.6 and Practice Problem 1.6 B. SI and Metric Prefixes 1. Prefixes are another way to express large or small numbers 2. The prefix indicates how the new unit relates to the original 3. See Table 1.4 for SI and metric prefixes and their relation to the original unit 4. Note there is more than one way to write equalities (1mL = 1x10-3L or 1000mL = 1L) 5. Prefixes can be applied to other units but it is uncommon (i.e. milliquarts) *See Sample Problem 1.7 and Practice Problem 1.7 V. Measurements and Significant Figures (Section 1.5) A. Accuracy and Precision 1. Accuracy is how close a measured value is to an accepted value (i.e. body temp is 98.6°F) 2. Precision is the reproducibility of a series of measurements (i.e. 98°F, 99°F, 97°F) 3. Measurements can be precise only, accurate only (on average), both precise and accurate, or neither precise and accurate 4. Examine Figure 1.14 B. Significant Figures 1. The quality of the equipment used to make measurements affects the accuracy and precision 2. Significant figures are defined as digits in a measurement that are reproducible when the measurement is repeated, plus the first uncertain digit (i.e. 5.671 g means the first three numbers 5.67 are certain while the last number 1 has some uncertainty) 3. Examine Figure 1.15 4. Reported zeros are sometimes significant and other times not (i.e. 2.00 has three significant figures while 0.009 only has one significant figure) 5. See Table 1.5 for rules involving significant figures 6. Significant figures only apply to measurements a. Exact numbers have an unlimited number of significant figures b. Come from exact counts (i.e. number of patients) and defined units (i.e. 1km = 1000m) *See Sample Problem 1.8 and Practice Problem 1.8 HEALTH LINK Body Mass Index (Examine Figure 1.16, Figure 1.17, and Table 1.6) C. Calculations with Significant Figures 1. Calculations should not change the degree of uncertainty in a value 2. When doing multiplication and division the answer should have the same number of significant figures as the quantity with the fewest (i.e. 5.3cm x 6.1cm should have only two: 32cm2) 3. The first digit to be removed is the digit immediately to right of the significant digit in a number *See Sample Problem 1.9 and Practice Problem 1.9 *See Sample Problem 1.10 and Practice Problem 1.10 4. When doing addition or subtractions the answer should have the same number of decimal places as the quantity with the fewest decimal places (i.e. 13.5g + 2.335g + 653g should have no places after the decimal: 669g) *See Sample Problem 1.11 and Practice Problem 1.11 HEALTH LINK Body Temperature (Examine Figure 1.18 and 1.19) VI. Conversion Factors and the Factor Label Method (Section 1.6) A. Unit Conversions 1-3
1. Many can be done in your head (i.e. 6 eggs are in a half dozen) 2. Others require a systematic approach called the factor label method B. Factor Label Method 1. Uses conversion factor to transform one unit into another 2. These are derived from the numerical relationship between two units (i.e. 2.205 lbs = 1kg) 3. All conversion factors are equal to 1 when divided 4. Review examples in text (What is the kilogram weight of 185 lb patient?) 5. Examine Figure 1.20 6. Make sure units cancel out, multiply the numerators, multiply the denominators, divide the two answers, and report final answer with proper units and significant figures DID YOU KNOW? Mars Climate Orbiter *See Sample Problem 1.12 and Practice Problem 1.12 *See Sample Problem 1.13 and Practice Problem 1.13 C. Temperature Conversions 1. To convert from °F to °C use one of the two equations (°F and °C are not the same size) a. °F = (1.8 x °C) + 32 b. °C = (°F – 32)/1.8 2. To convert from °C and K use one the two equations (°C and K are the same size and note there are no degrees in Kelvin) a. K = °C + 273.15 b. °C = K – 273.15 *See Sample Problem 1.14 and Practice Problem 1.14 VII. Density, Specific Gravity, and Specific Heat (Section 1.7) A. Density 1. The amount of mass contained in a given volume 2 Usually expressed in g/cm3 for solids, g/mL for liquids and g/L for gases 3. Density of water at 20°C is 1.00 g/mL 4. As temperature varies, density varies (i.e. how mercury in a thermometer works) 5. See Table 1.7 for the densities of common substances 6. For calculations density (D) = mass (m)/volume (V) 7. Review examples in text (What is the mass in grams of 2.15L of isopropyl alcohol?) *See Sample Problem 1.15 and Practice Problem 1.15 B. Specific Gravity 1. Relates the density of a substance to that of water 2. Specific Gravity = (Density of Substance)/(Density of Water) 3. Review example in text (What is the specific gravity of isopropyl alcohol?) 4. A substance with a specific gravity of less than 1 will float in water, while one that has a specific gravity of greater than 1 will sink 5. Specific gravity has a variety of applications (i.e. car batteries, alcohol content, urine analysis) 6. Examine Figure 1.21 HEALTH LINK Making Weight (Examine Figure 1.22) C. Specific Heat 1. The amount of heat required to raise the temperature of 1 gram of a substance by 1°C 2. Relates energy (cal), mass (g), and temperature (°C) 3. Review example in text (How many calories of heat are needed to increase the temperature of 2.5g of water by 5.0°C?) 4. See Table 1.8 for some specific heats *See Sample Problem 1.16 and Practice Problem 1.16 VIII. Measurements in General Chemistry, Organic Chemistry, and Biochemistry (Section 1.8) A. Sub-Disciplines of Chemistry 1-4
1. General chemistry is a study of the fundamental principles of chemistry (i.e. atoms, units, etc.) 2. Organic chemistry is a study of the chemistry of carbon 3. Biochemistry is a study of the chemistry of living things B. Examples of how chemistry relates to what goes on around you 1. A look at Cadmium (Examine Figure 1.23) *See Sample Problem 1.17 and Practice Problem 1.17 2. A look at Gasoline (Examine Figure 1.24 and 1.25) *See Sample Problem 1.18 and Practice Problem 1.18 3. A look at DNA (Examine Figure 1.26 and 1.27) *See Sample Problem 1.19 and Practice Problem 1.19
Chapter Summary Section 1.1 Science encompasses many subjects that greatly impact our lives. One such subject is chemistry, which is defined as the study of mater and its changes. Chemistry lays the groundwork for many other fields of study. Science, as a discipline, comes from the process known as the scientific method. The first step of the scientific method requires an observation. Next, a hypothesis is formed, which is a tentative and testable explanation of the observation. The hypothesis is then tested with an experiment. If the hypothesis survives repeated testing, then that hypothesis may become part of a theory. A theory should not be confused with a law, which is a statement that describes things consistently and reproducibly observed. Theories never become laws. Theories are an explanation and laws are a description. Section 1.2 Matter is anything that has mass and occupies space. The physical properties of matter can be determined without changing its chemical composition (what it’s made of). Some examples of physical properties include malleability, taste, melting point, and state (solid, liquid, or gas). A change in state is an example of a physical change (one that does not alter the chemical composition of matter). When matter is physically changed, energy is involved. Energy is the ability to do work or transfer heat. Energy can be found in two forms: potential energy (stored energy) and kinetic energy (energy of motion). Heat energy plays a major role in physical changes. For example, the heat of fusion is energy put in during the melting process (solid to liquid) and the heat of vaporization is the energy put in during the boiling/vaporization process (liquid to gas). The physical change of matter’s state can also sometimes skip the liquid phase. Sublimation is the conversion of a solid to a gas where deposition is the conversion of a gas to solid. Section 1.3 Scientific measurements consist of a number and a unit. The three systems of measurement units are the metric units, the English units, and the SI units. Five types of common measurements are mass, length, volume, temperature, and energy. Mass is measure of the amount of matter in a sample, which is different from weight (gravity dependent). Length is the measure of distance and volume is the measure of the space occupied by an object. There are three scales used to measure temperature: degrees Fahrenheit (English unit), degrees Celsius (metric unit), and Kelvin (SI unit based on absolute zero). The calorie is the metric and English unit for energy. One calorie is the amount of energy required to raise 1 gram of water 1°C. The joule (SI unit) and the food calorie are also used for energy measurements. Section 1.4 Scientific notation, which condenses very large or very small numbers, expresses a number between 1 and 10 multiplied by a power of ten. Positive exponents shift the decimal left (multiply by 10 “n” times); conversely, negative exponents shift the decimal right (divide by 10 “n” times). SI & metric prefixes are another way to express large or small numbers. The prefix indicates how the new unit relates to the original unit. Table 1.4 summarizes the common SI and metric prefixes. 1-5
Section 1.5 The accuracy of a measurement refers to how close a measured value is to a true value. Precision, on the other hand, is the measure of reproducibility. The quality of the equipment used to make measurements affects both accuracy and precision. Significant figures can be used to indicate the certainty of a measurement. Significant figures are defined as digits in a measurement that are reproducible when the measurement is repeated, plus the first uncertain digit. Significant figures only apply to measurements: exact numbers have an unlimited number of significant figures. Calculations should not change the degree of uncertainty in a value. When performing multiplication and division, the answer should have the same number of significant figures as the quantity with the fewest. When performing addition or subtractions, the answer should have the same number of decimal places as the quantity with the fewest decimal places. Section 1.6 When dealing with measurements it is often necessary to perform unit conversions. One approach to unit conversions is the factor label method, which uses conversion factors to transform one unit into another. Conversion factors are derived from the numerical relationship between two units, and when divided, they equal 1. Temperature conversions require equations that come from the relationships of the different temperature scales. For the conversion between degrees Fahrenheit and degrees Celsius the relationship is °F = (1.8 x °C) + 32 or °C = (°F – 32)/1.8. For the conversion between degrees Celsius and Kelvin the relationship is K = °C + 273.15 or °C = K – 273.15. Section 1.7 Another unit used to describe matter is density, which is the amount of mass contained in a given volume. Density is usually expressed in g/cm3 for solids, g/mL for liquids and g/L for gases. The density of water at 20°C is 1.00 g/mL(temperature affects density). Specific gravity relates the density of a substance to that of water. This relationship, which has many applications, is summarized by the equation Specific Gravity = (Density of Substance)/(Density of Water). A substance with a specific gravity of less than 1 will float in water, while one that has a specific gravity of greater than 1 will sink. Specific heat is yet another unit used to describe matter. Specific heat is the amount of heat required to raise the temperature of 1 gram of a substance by 1°C. It relates energy (cal), mass (g), and temperature (°C). Section 1.8 Chemistry has many sub-disciplines, three of which will be studied in this text. General chemistry is a study of the fundamental principles of chemistry. Organic chemistry is a study of the chemistry of carbon. Biochemistry is a study of the chemistry of living things. Some examples of how chemistry relates to what goes on around us include: a look at cadmium (General Chemistry), a look at gasoline (organic chemistry), and a look at DNA (biochemistry).
Lecture Suggestions Section 1.1 and 1.2 To introduce these section, ask students what makes a science textbook different from another subject’s textbook (history, English, etc.). Many students will say it is because a science book contains only facts. Explain that it is incorrect to describe science as simply a collection of facts. Describe how science is actually a process known as the scientific method, and the information in this textbook is the result of numerous scientists doing the scientific method over numerous generations. While the end results or “facts” are very important, it is the process (scientific method) that distinguishes science from other areas of study. Ask students to name ways in which science has directly affected their lives. 1-6