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Introduction to Organic Chemistry, Student Solutions Manual
Pathophysiology for the Health Professions
Study Guide for Pathophysiology for the Health Professions
1 1-1 Chapter 1: Matter and Energy Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 1-2 Chapter 1 Topics: 1. Matter and its Classification 2. Physical and Chemical Changes and Properties of Matter 3. Energy and Energy Changes 4. Scientific Inquiry 1-3 1.1 Scientific Notation -- Also called exponential notation 1.2 Significant Figures 1.3 Units and Conversions Chapter 1 Math Toolboxes: 1-4 1.1 Matter and Its Classification ? Matter is anything that occupies space and has mass. ? Forms of energy are NOT matter. Heat and light, for example, do not occupy space and have no mass. 1-5 Composition of Matter ? We classify matter so that we can understand it better. ? One way to classify matter is as pure substances or mixtures. Matter Pure Substances Mixtures Always has the same composition 1-6 Types of Pure Substances ? Two kinds of pure substances ? Elements ? An element is a substance that cannot be broken down into simpler substances even by a chemical reaction. ? Elements are separated further into metals and nonmetals. ? Compounds ? A compound is a substance composed of two or more different elements combined in definite proportions. 2 1-7 Figure 1.3 1-8 Atoms ? An atom is the smallest unit of an element that retains the chemical properties of that element. Figure 1.10 1-9 Elements and their Symbols ? Element symbols often consist of one or two letters of the element?s name. ? Examples: carbon: C calcium: Ca ? How do we explain that Fe is the symbol for iron? 1-10 Elements and their Symbols ? Know the names and symbols of the most common elements indicated (same as on handout) 1-11 Pure Substances Pure Substances Elements Compounds Metals Nonmetals 1-12 Activity: Classification of Matter Metals generally have a shiny appearance; most are a silver color. Metals conduct electricity and heat. Most are malleable. Figure 1.4 3 1-13 Figure 1.3 Metal elements are to the left of the stair-step line on the periodic table 1-14 ? Examples of elements ? Examples of compounds 1-15 Compounds ? A compound is a pure substance composed of two or more different elements combined chemically in definite proportions. ? A compound has properties that are different from those of its component elements. Compound: Iron pyrite, Fe 2 S 3 Mixture of Elements: Iron, Fe, and Sulfur, S Figure 1.5 1-16 Compounds ? Pure sand is the compound silicon dioxide, SiO 2 . Figure 1.1 1-17 Activity: Elements and Compounds ? Identify each of the following as an element or compound. 1. He 2. H 2 O 3. sodium chloride 4. copper 5. Cl 2 1. Element 2. Compound 3. Compound 4. Element 5. Element 1-18 Mixtures ? A mixture is a combination of two or more elements or compounds . ? Mixtures differ from pure compounds in that their components can be separated by physical processes. ? Examples: ? Pencil lead ? Lemonade ? Air 4 1-19 Mixtures can be separated by physical processes ? The components of salt water can be separated by evaporating the water?. ? Magnets work to separate Fe from S Figure 1.6 1-20 Mixtures ? Mixtures can be further classified as homogeneous and heterogeneous. ? Homogeneous mixtures have the same composition throughout. ? Heterogeneous mixtures do not. Mixtures ? consist of 2 or more pure substances Homogeneous Mixtures (solutions) - have uniform composition throughout Heterogeneous Mixtures ? do not have uniform composition throughout 1-21 Mixtures ? Classify each of the following mixtures as homogeneous or heterogeneous: ? Salt water ? Lemonade ? Tap water ? Air ? Cake mix ? Oil and water 1-22 Representations of Matter ? Macroscopic ? we can see with our eyes ? Molecular level ? a magnification to a level that shows atoms ? Symbolic ? shorthand using element symbols ? Atom ? the smallest unit of an element; represented as single sphere. ? Molecule ? two or more bound atoms 1-23 Molecular-Level Representations of Matter ? Helium Atoms Figure 1.9 1-24 Atoms and Molecules Molecules are composed of two or more atoms bound together in a discrete arrangement. ? The atoms bound together in a molecule can be from the same element or from different elements. ? The atoms in each molecule are the same, so this represents an element. O 2 Molecules 5 1-25 Molecular-Level Representations of Matter ? Water molecules Figure 1.10 1-26 Different Ways to Represent Water Figure 1.12igure 1.12 1-27 Activity: Classify each of the following as an element, compound, or mixture. Figure from p. 48 1-28 Classify each as a pure substance or mixture. 1-29 States of Matter ? A different way to classify matter is by its physical state: solid, liquid, or gas. ? How do the atoms and molecules of solids, liquids, and gases behave differently? Change of State 1-30 Gases can be compressed Figure 1.15 6 1-31 Solid and Liquid States of Iron Figure 1.14 1-32 Symbols Used in Chemistry ? Symbols for physical states ? are written in parenthesis after the elemental symbol or chemical formula. 1-33 NaCl(aq) ? Describe what the symbol NaCl(aq) represents, and classify as an element, compound, or mixture. 1-34 1.2 Physical and Chemical Changes and Properties of Matter ? A physical property is a characteristic that we can observe without changing the composition of a substance. ? Examples ? Color ? Odor ? Mass ? Volume ? Density ? Temperature quantitative properties qualitative properties 1-35 Math Toolbox 1.3 ? Units and Conversions Metric Base Units and Derived Units ? Length: meter (m) ? Mass: kilogram (kg) ? Time: second (s) ? Temperature: kelvins (K) ? Numbers of particles: mole (mol) 1-36 Mass ? Mass: ? measures the quantity of matter. ? has common units of grams (g). 7 1-37 SI Units for Mass ? A milligram (mg) is 1/1000 of a gram: ? The mass of tiny salt crystals usually has units of milligrams (mg). ? (~ 0.2 mg) ? A kilogram (kg) is 1000 times that of a gram: ? The mass of a person usually has units of kilograms (kg). ? (~ 50-100 kg) Figure 1.17 1-38 Unit Relationships (Mass) (See Math Toolbox 1.3 & back cover of text) Prefixes (Table 1.4) Mass measurements giga- G 10 9 1 Gg = 10 9 g mega- M 10 6 1 Mg = 10 6 g kilokilo- k 10 3 1 kg = 10 3 g centi- c 10 -2 1 cg = 10 -2 g milli- m 10 -3 1 mg = 10 -3 g micro- ? 10 -6 1 ?g = 10 -6 g nano- n 10 -9 1 ng = 10 -9 g pico- p 10 -12 1 pg = 10 -12 g 1-39 Unit Relationships (Length) (See Math Toolbox 1.3 & back cover of text) Prefixes (Table 1.4) Length measurements giga- G 10 9 1 Gm = 10 9 m mega- M 10 6 1 Mm = 10 6 m kilo- k 10 3 1 km = 10 3 m centi- c 10 -2 1 cm = 10 -2 m milli- m 10 -3 1 mm = 10 -3 m micro- ? 10 -6 1 ?m = 10 -6 m nano- n 10 -9 1 nm = 10 -9 m pico- p 10 -12 1 pm = 10 -12 m 1-40 Math Toolbox 1.1 ? Scientific Notation ? Powers of Ten (Slide Show). ? 0.000523 = 5.23×10 ?4 1-41 Activity: Mass Unit Conversions (See Math Toolbox 1.3) ? Convert 12.0 grams to milligrams ? (1 g = 1000 mg) ? Convert 12.0 grams to ounces ? (1 oz = 28.34 g) 1-42 Activity: Mass Unit Conversions (See Math Toolbox 1.3) ? Convert 12.0 grams to milligrams. ? (1 g = 1000 mg) ? Convert 12.0 grams to ounces ? (1 oz = 28.34 g) 12,000 mg or 1.20×10 4 mg 0.423 oz 8 1-43 Significant Figures (See Math Toolbox 1.2) ? Significant Figures ? All non-zero digits are significant. (435 g) ? A zero that falls between two significant digits is significant. (405 g; 40.5 g) ? Zeros to the right of a sig. digit and to the right of a decimal pt. are significant. (5.00 g) ? Zeros to the left of the first significant digit are not significant. (0.151 g; 0.00405 g) ? If a number is >1, the zeros to the right of the last nonzero digit may or may not be significant. Use scientific notation to specify. 1-44 Volume ? Volume: ? amount of space a substance occupies ? If a solid with rectangular shape: V = length × width × height ? Common units are centimeters cubed (cm 3 ) or milliliters (mL) Figure 1.18 1-45 Volume ? We can measure the volume of a cube by measuring the length of one of its sides, and then cubing the length. If the length of a side is 2.0 cm, what is the volume of this cube? 2.0 centimeters 1-46 Density ? The density of a substance is the ratio of its mass to volume: ? If the mass of the cube is 11.2 grams, what is its density? 2.0 centimeters mass Density = volume 1-47 Density ? If another piece of the substance from the last slide has a volume of 20.0 mL, what is its mass? mass Density = volume 1-48 Mass From Density ? V = 20.0 mL = 20.0 cm 3 ? D = 1.4 g/cm 3 mass D = volume 9 1-49 Density ? Which liquid is the least dense? ? Compare the density of the Dead Sea water and of the person. Figure 1.19 © J. Birk 1-50 Density = mass/volume ? Why is regular soda more dense? Figure from p. 20 1-51 Why is ice less dense than liquid water? Figure from p. 21 1-52 Temperature ? Temperature: ? SI units are degrees Celsius (°C) and kelvins (K) T K = T °C + 273.15 T °F = 1.8(T °C ) + 32 Figure 1.22 1-53 Temperature Scales ? T K = T °C + 273.15 ? T °F = 1.8(T °C ) + 32 ? Boiling Point of Water: ? 212°F, 100°C, 373.15 K ? Freezing Point of Water: ? 32°F, 0°C, 273.15 K ? Lowest Possible Temperature: ? ?273.15°C, 0.00 K 1-54 Chemical vs. Physical Changes In a chemical change, a new substance (element or compound) is formed. When heated, iron reacts with sulfur to form a compound, iron(III) sulfide. Before heating, the mixing process was just a physical change. 10 1-55 Physical Changes ? A physical change ? is a process that changes the physical properties of a substance without changing its chemical composition. ? Physical state transitions are physical changes because they do not change the identity of the substance. Figure 1.24 1-56 Vaporization or Evaporation Figure 1.23 1-57 Sublimation of Dry Ice (CO 2 ) ? CO 2 (s) ??CO 2 (g) Figure 1.13 1-58 Water vapor condenses from the air onto the cold surface of the glass. Figure 1.16 Gas Liquid 1-59 Physical State Transitions Figure 1.24 1-60 Chemical Changes ? A chemical change is a process where one or more substances are converted into one or more new substances. ? Also called a chemical reaction. ? Examples: ? Pennies tarnishing ? Burning gasoline ? The reaction of hydrogen and oxygen to form water 11 1-61 Chemical Changes ? A chemical change ? is a process in which one or more substances are converted into one or more new substances. ? Evidence of a chemical change includes: ? bubbling ? a permanent color change ? a sudden change in temperature ? a flame Figure 1.25 1-62 Chemical Changes Sodium metal reacts with chlorine gas to form sodium chloride VD02_007.mov 1-63 Is Boiling Water a Chemical or Physical Change? Figure 1.27B 1-64 Activity: Physical and Chemical Changes 1-65 Activity: Physical and Chemical Changes 1-66 Activity: Physical vs. Chemical Changes ? Classify each of the following as a physical or chemical change: ? Evaporation of water ? Burning of natural gas ? Melting a metal ? Converting H 2 and O 2 to H 2 O 12 1-67 Activity: Physical vs. Chemical Changes ? Classify each of the following as a physical or chemical change: ? Evaporation of water ? Burning of natural gas ? Melting a metal ? Converting H 2 and O 2 to H 2 O Physical Chemical Physical Chemical 1-68 Chemical Properties ? Chemical properties are descriptions of the ability of a substance to undergo a chemical change. ? Examples: ? Hydrogen burns easily with oxygen ? Helium is unreactive ? Iron rusts ? Silver tarnishes ? Gold is very unreactive 1-69 1.3 Energy and Energy Changes Two main forms of energy are: ? Kinetic energy: the energy of motion ? The kinetic energy of a sample?s particles will increase as temperature is increased. ? Potential energy: energy possessed by an object because of its position ? As a ball is raised up in the air, its potential energy increases. ? Very reactive substances have high potential energy. 1-70 Which pair of molecules has more kinetic energy? Figure from p. 50 1-71 Kinetic and Potential Energy Figure 1.29 1-72 Energy and Energy Changes ? Energy changes accompany chemical and physical changes. ? Examples: ? When wood burns with oxygen, energy in the form of heat is released. ? When ammonium nitrate dissolves in water in a cold pack, energy in the form of heat is absorbed. marks Microsoft PowerPoint - Ch.1 2e_OL.ppt
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