Chemistry: Atoms First 2e Chapter 1 PDF

CHEMISTRY: ATOMS FIRST 2e Chapter 1 ESSENTIAL IDEAS PowerPoint Image Slideshow WITH CONTRIBUTIONS BY JOE DEPASQUALE, NORTHEASTERN UNIVERSITY Chapter Outline 1.1 Chemistry in Context 1.2 Phases and Classification of Matter 1.3 Physical and Chemical Properties...

CHEMISTRY: ATOMS FIRST 2e Chapter 1 ESSENTIAL IDEAS PowerPoint Image Slideshow WITH CONTRIBUTIONS BY JOE DEPASQUALE, NORTHEASTERN UNIVERSITY Chapter Outline 1.1 Chemistry in Context 1.2 Phases and Classification of Matter 1.3 Physical and Chemical Properties 1.4 Measurements 1.5 Measurement Uncertainty, Accuracy, and Precision 1.6 Mathematical Treatment of Measurement Results Figure 1.1 Chemical substances and processes are essential for our existence, providing sustenance, keeping us clean and healthy, fabricating electronic devices, enabling transportation, and much more. (credit “left”: modification of work by “vxla”/Flickr; credit “left middle”: modification of work by “the Italian voice”/Flickr; credit “right middle”: modification of work by Jason Trim; credit “right”: modification of work by “gosheshe”/Flickr) Learning Objectives 1.1 Chemistry in Context Outline the historical development of chemistry Provide examples of the importance of chemistry in everyday life Describe the scientific method Differentiate among hypotheses, theories, and laws Provide examples illustrating macroscopic, microscopic, and symbolic domains Chemistry in Context Chemistry is the study of the composition, properties, and interactions of matter. Attempts to understand the behavior of matter extend back more than 2,500 years. Greeks: Matter consists of four elements: earth, air, fire, and water. Alchemists attempted to transform “base metals” into “noble metals.” Figure 1.2 This portrayal shows an alchemist’s workshop circa 1580. Although alchemy made some useful contributions to how to manipulate matter, it was not scientific by modern standards. (credit: Chemical Heritage Foundation) Chemistry the Central Science Chemistry is interconnected to a vast array of other STEM disciplines. Figure 1.3 Knowledge of chemistry is central to understanding a wide range of scientific disciplines. This diagram shows just some of the interrelationships between chemistry and other fields. Chemistry and Everyday Life Examples of chemistry in everyday life: Digesting food Synthesizing polymers for clothing, cookware, and credit cards Refining crude oil into gasoline and other products As you proceed through this course, you will discover: Many different examples of changes in the composition and structure of matter. How to classify these changes in matter and understand how they occur. The changes in energy that accompany these changes in matter. The Scientific Method Chemistry is a science based on observation and experimentation. Chemists often formulate a hypothesis: a tentative explanation of observations. The laws of science summarize a vast number of experimental observations, and describe or predict some facet of the natural world. Theory: A well-substantiated, comprehensive, testable explanation of a particular aspect of nature. The Domains of Chemistry Chemists study and describe the behavior of matter and energy in three different domains. 1) The macroscopic domain is familiar to us: It is the realm of everyday things that are large enough to be sensed directly by human sight or touch. 2) The microscopic domain of chemistry is almost always visited in the imagination. Micro also comes from Greek and means “small.” Some aspects of the microscopic domains are visible through a microscope. 3) The symbolic domain contains the specialized language used to represent components of the macroscopic and microscopic domains, such as chemical symbols. Figure 1.4 The scientific method follows a process similar to the one shown in this diagram. All the key components are shown, in roughly the right order. Scientific progress is seldom neat and clean: It requires open inquiry and the reworking of questions and ideas in response to findings. Figure 1.5 (a) Moisture in the air, icebergs, and the ocean represent water in the macroscopic domain. (b) At the molecular level (microscopic domain), gas molecules are far apart and disorganized, solid water molecules are close together and organized, and liquid molecules are close together and disorganized. (c) The formula H2O symbolizes water, and (g), (s), and (l) symbolize its phases. Note that clouds are actually comprised of either very small liquid water droplets or solid water crystals; gaseous water in our atmosphere is not visible to the naked eye, although it may be sensed as humidity. (credit a: modification of work by “Gorkaazk”/Wikimedia Commons) Learning Objectives 1.2 Phases and Classification of Matter Describe the basic properties of each physical state of matter: solid, liquid, and gas Distinguish between mass and weight Apply the law of conservation of matter Classify matter as an element, compound, homogeneous mixture, or heterogeneous mixture with regard to its physical state and composition Define and give examples of atoms and molecules Phases and Classification of Matter Matter: Anything that occupies space and has mass. The three most common states or phases of matter: 1) A solid is rigid and possesses a definite shape. 2) A liquid flows and takes the shape of its container. 3) A gas takes both the shape and volume of its container. Figure 1.6 The three most common states or phases of matter are solid, liquid, and gas. Plasma: A Fourth State of Matter Plasma: A gaseous state of matter that contains an appreciable amount of electrically charged particles. Plasma has unique properties distinct from ordinary gases. Plasma is found in certain high-temperature environments. Naturally: Stars, lightning Man-made: Television screens Figure 1.7 A plasma torch can be used to cut metal. (credit: “Hypertherm”/Wikimedia Commons) Mass vs. Weight Mass is a measure of the amount of matter in an object. Weight refers to the force that gravity exerts on an object. An object’s mass is the same on the earth and the moon but its weight is different. Law of Conservation of Matter Law of conservation of matter: There is no detectable change in the total quantity of matter present when matter converts from one type to another. This is true for both chemical and physical changes. Figure 1.8 (a) The mass of beer precursor materials is the same as the mass of beer produced: Sugar has become alcohol and carbonation. (b) The mass of the lead, lead oxide plates, and sulfuric acid that goes into the production of electricity is exactly equal to the mass of lead sulfate and water that is formed. Elements An element is a type of pure substance that cannot be broken down into simpler substances by chemical changes. The known elements are displayed in the periodic table. There are more than 100 known elements. Ninety of these occur naturally. Two dozen or so have been created in laboratories. Pure Substances and Mixtures Pure substances have constant composition. Elements: Pure substance that cannot be broken down into simpler substances by chemical changes. Consist of one type of element Examples: Gold (Au), Phosphorus (P), Oxygen (O) Compounds: Pure substances that can be broken down into simpler substances by chemical changes. Consist of two or more types of elements chemically bonded Examples: H2O, C6H12O6, AgCl The properties of compounds are different from the uncombined elements making up the compound. Figure 1.9 (a) The compound mercury(II) oxide, (b) when heated, (c) decomposes into silvery droplets of liquid mercury and invisible oxygen gas. (credit: modification of work by Paul Flowers) Pure Substances and Mixtures A mixture is composed of two or more types of matter that can be present in varying amounts and can be separated by physical changes. Evaporation is an example of a physical change. There are two types of mixtures: homogenous mixtures and heterogeneous mixtures. Two Type of Mixtures 1) A homogenous mixture exhibits a uniform composition and appears visually the same throughout. Another name for a homogenous mixture is a solution. 2) A heterogeneous mixture has a composition that varies from point to point. Figure 1.10 (a) Oil and vinegar salad dressing is a heterogeneous mixture because its composition is not uniform throughout. (b) A commercial sports drink is a homogeneous mixture because its composition is uniform throughout. (credit a “left”: modification of work by John Mayer; credit a “right”: modification of work by Umberto Salvagnin; credit b “left: modification of work by Jeff Bedford) Figure 1.11 Depending on its properties, a given substance can be classified as a homogeneous mixture, a heterogeneous mixture, a compound, or an element. Atoms and Molecules Atom: The smallest particle of an element that has the properties of that element and can enter into a chemical combination. Idea first proposed by Greek philosophers, Leucippus and Democritus, in the 5th century BCE. 19th century, John Dalton of England supported this hypothesis with quantitative measurements. Molecule: Consists of two or more atoms connected by strong forces known as chemical bonds. Figure 1.12 (a) This photograph shows a gold nugget. (b) A scanning-tunneling microscope (STM) can generate views of the surfaces of solids, such as this image of a gold crystal. Each sphere represents one gold atom. (credit a: modification of work by United States Geological Survey; credit b: modification of work by “Erwinrossen”/Wikimedia Commons) Figure 1.13 These images provide an increasingly closer view: (a) a cotton boll, (b) a single cotton fiber viewed under an optical microscope (magnified 40 times), (c) an image of a cotton fiber obtained with an electron microscope (much higher magnification than with the optical microscope); and (d and e) atomic-level models of the fiber (spheres of different colors represent atoms of different elements). (credit c: modification of work by “Featheredtar”/Wikimedia Commons) Figure 1.14 The elements hydrogen, oxygen, phosphorus, and sulfur form molecules consisting of two or more atoms of the same element. The compounds water, carbon dioxide, and glucose consist of combinations of atoms of different elements. Figure 1.15 The decomposition of water is shown at the macroscopic, microscopic, and symbolic levels. The battery provides an electric current (microscopic) that decomposes water. At the macroscopic level, the liquid separates into the gases hydrogen (on the left) and oxygen (on the right). Symbolically, this change is presented by showing how liquid H2O separates into H2 and O2 gases. Figure 1.16 A fuel cell generates electrical energy from hydrogen and oxygen via an electrochemical process and produces only water as the waste product. Figure 1.17 Almost one-third of naturally occurring elements are used to make a cell phone. (credit: modification of work by John Taylor) Learning Objectives 1.3 Physical and Chemical Properties Identify properties of and changes in matter as physical or chemical Identify properties of matter as extensive or intensive Physical and Chemical Properties The characteristics that enable us to distinguish one substance from another are called properties. A physical property is a characteristic of matter that is not associated with a change in its chemical composition. Examples: density, color, hardness, melting and boiling points, and electrical conductivity A physical change is a change in the state or properties of matter without any accompanying change in its chemical composition. Figure 1.18 (a) Butter undergoes a physical change when solid butter is heated and forms liquid melted butter. (b) Steam condensing inside a cooking pot is a physical change, as water vapor is changed into liquid water. (credit a: modification of work by “95jb14”/Wikimedia Commons; credit b: modification of work by “mjneuby”/Flickr) Physical and Chemical Properties The change of one type of matter into another type (or the inability to change) is a chemical property. Examples: flammability, toxicity, acidity, reactivity, and heat of combustion. Figure 1.19 (a) One of the chemical properties of iron is that it rusts; (b) one of the chemical properties of chromium is that it does not. (credit a: modification of work by Tony Hisgett; credit b: modification of work by “Atoma”/Wikimedia Commons) Figure 1.20 (a) Copper and nitric acid undergo a chemical change to form copper nitrate and brown, gaseous nitrogen dioxide. (b) During the combustion of a match, cellulose in the match and oxygen from the air undergo a chemical change to form carbon dioxide and water vapor. (c) Cooking red meat causes a number of chemical changes, including the oxidation of iron in myoglobin that results in the familiar red-to-brown color change. (d) A banana turning brown is a chemical change as new, darker (and less tasty) substances form. (credit b: modification of work by Jeff Turner; credit c: modification of work by Gloria Cabada-Leman; credit d: modification of work by Roberto Verzo) Figure 1.21 The National Fire Protection Agency (NFPA) hazard diamond summarizes the major hazards of a chemical substance. Extensive Properties and Intensive Properties Extensive property Depends on the amount of matter present. Examples: mass, volume, heat Intensive property Does not depend on the amount of matter present. Examples: density, temperature Figure 1.22 The periodic table shows how elements may be grouped according to certain similar properties. Note the background color denotes whether an element is a metal, metalloid, or nonmetal, whereas the element symbol color indicates whether it is a solid, liquid, or gas. Learning Objectives 1.4 Measurements Explain the process of measurement Identify the three basic parts of a quantity Describe the properties and units of length, mass, volume, density, temperature, and time Perform basic unit calculations and conversions in the metric and other unit systems Measurements Measurements provide the information that is the basis of most of the hypotheses, theories, and laws in chemistry. Every measurement provides three kinds of information: 1) The size or magnitude of the measurement: a number 2) A standard of comparison for the measurement: a unit 3) An indication of the uncertainty of the measurement. Units Without units, a number can be meaningless or confusing! In chemistry, we use an updated version of the metric system known as the International System of Units, or SI units. Used since 1964 Table 1.2 Base Units of the SI System Property Name of Unit Symbol of Unit length meter m mass kilogram kg time second s temperature kelvin K electric current ampere A amount of substance mole mol luminous Intensity candela cd Table 1.3 Common Unit Prefixes Fractional or multiple SI units are named using a prefix and the name of the base unit. Prefix Symbol Factor femto f 10–15 pico p 10–12 nano n 10–9 micro m 10–6 milli m 10–3 centi c 10–2 deci d 10–1 Table 1.3: Common Unit Prefixes Prefix Symbol Factor kilo k 103 mega M 106 giga G 109 tera T 1012 Common SI Base Units: Length The SI unit of length is the meter (m). The meter was originally intended to be 1/10,000,000 of the distance from the North Pole to the equator. A meter is now defined as the distance light travels in a vacuum in 1/299,792,458 of 1 second. A meter is about 3 inches longer than 1 yard. Figure 1.23 The relative lengths of 1 m, 1 yd, 1 cm, and 1 in. are shown (not actual size), as well as comparisons of 2.54 cm and 1 in., and of 1 m and 1.094 yd. Common SI Base Units: Mass The SI unit of mass is the kilogram (kg). A kilogram was originally defined as the mass of a liter of water. It is now defined by a certain cylinder of platinum-iridium alloy, which is kept in France. 1 kilogram is about 2.2 pounds. Figure 1.24 This replica prototype kilogram is housed at the National Institute of Standards and Technology (NIST) in Maryland. (credit: National Institutes of Standards and Technology) Common SI Base Units: Temperature The SI unit of temperature is the kelvin (K). No degree word nor symbol (°) is used with kelvin. The degree Celsius (°C) is also allowed in the SI system. Celsius degrees are the same magnitude as those of kelvin, but the two scales place their zeros in different places. Water freezes at 273.15 K (0 °C) and boils at 373.15 K (100 °C). Common SI Base Units: Time The SI unit of time is the second (s). Smaller and larger time intervals can be expressed with the appropriate prefixes. Alternatively, hours, days, and years can be used. Derived SI Units: Volume and Density We can derive many units from the seven SI base units. Volume: The measure of the amount of space occupied by an object. The standard SI unit for volume is the cubic meter (m3), which is derived from the SI base unit of length. Other units for volume are the liter (L) and milliliter (mL). 1 dm3 = 1 L 1 cm3 = 1 mL Figure 1.25 (a) The relative volumes are shown for cubes of 1 m3, 1 dm3 (1 L), and 1 cm3 (1 mL) (not to scale). (b) The diameter of a dime is compared relative to the edge length of a 1-cm3 (1- mL) cube. Density The density of a substance is the ratio of the mass of a sample of the substance to its volume. mass density = volume The standard SI unit for density is the kilogram per cubic meter (kg/m3). Commonly used density units based on state of matter: g/cm3 (solids, liquids) g/L (gases) Learning Objectives 1.5 Measurement Uncertainty, Accuracy, and Precision Define accuracy and precision Distinguish exact and uncertain numbers Correctly represent uncertainty in quantities using significant figures Apply proper rounding rules to computed quantities Measurement Uncertainty, Accuracy, and Precision Counting is the only type of measurement that is free from uncertainty. The result of a counting measurement is an example of an exact number. The numbers for defined quantities are also exact. 1 ft. is exactly 12 in. 1 in. is exactly 2.54 cm 1 g is exactly 0.001 kg 1.5 Measurement Uncertainty, Accuracy, and Precision Quantities derived from measurements other than counting are uncertain to varying extents. These numbers are not exact. There are always practical limitations of the measurement process used. A measured number must be reported in a way to indicate its uncertainty. In general, when recording a measurement, you are allowed to estimate one uncertain digit. Figure 1.26 To measure the volume of liquid in this graduated cylinder, you must mentally subdivide the distance between the 21 and 22 mL marks into tenths of a milliliter, and then make a reading (estimate) at the bottom of the meniscus. Significant Figures On the previous slide, if one recorded the volume in the graduated cylinder to be 21.6 mL: 2 and 1 are certain digits. 6 is an estimate. Someone else might perceive the volume to be 21.5 mL or 21.7mL. All the digits in a measurement, including the uncertain last digit, are called significant figures or significant digits. Frequently, we need to know the number of significant figures in a measurement reported by someone else. Significant Figures These numbers are always significant. Nonzero digits Captive zeros Trailing zeroes When they are to the right of the decimal place When in scientific notation These numbers are always not significant. Leading zeros Trailing zeros When they are to the left of the decimal place Significant Figures Significant Figures in Calculations Results calculated from measured numbers are at least as uncertain as the measurement itself. 1) When we add or subtract numbers, we should round the result to the same number of decimal places as the number with the least number of decimal places (the least precise value in terms of addition and subtraction). 2) When we multiply or divide numbers, we should round the result to the same number of digits as the number with the least number of significant figures (the least precise value in terms of multiplication and division). 3) If the digit to be dropped (the one immediately to the right of the digit to be retained) is less than 5, we “round down” and leave the retained digit unchanged; if it is more than 5, we “round up” and increase the retained digit by 1; if the dropped digit is 5, we round up or down, whichever yields an even value for the retained digit. Significant Figures in Calculations The following examples illustrate the application of this rule in rounding a few different numbers to three significant figures: 0.028675 rounds “up” to 0.0287 (the dropped digit, 7, is greater than 5) 18.3384 rounds “down” to 18.3 (the dropped digit, 3, is less than 5) 6.8752 rounds “up” to 6.88 (the dropped digit is 5, and the retained digit is even) 92.85 rounds “down” to 92.8 (the dropped digit is 5, and the retained digit is even) Example 1.7 Example 1.7 Accuracy and Precision A measurement is said to be precise if it yields very similar results when repeated in the same manner. A measurement is considered accurate if it yields a result that is very close to the true or accepted value. Figure 1.27 (a) These arrows are close to both the bull’s eye and one another, so they are both accurate and precise. (b) These arrows are close to one another but not on target, so they are precise but not accurate. (c) These arrows are neither on target nor close to one another, so they are neither accurate nor precise. Table 1.5 Volume (ML) of Cough Medicine Delivered by 10-oz (296 ML) Dispensers Dispenser #1 Dispenser # 2 Dispenser # 3 283.3 298.3 296.1 284.1 294.2 295.9 283.9 296.0 296.1 284.0 297.8 296.0 284.1 293.9 296.1 Dispenser #1 is precise, but not accurate. Dispenser #2 is more accurate, but less precise. Dispenser #3 is both accurate and precise. Learning Objectives 1.6 Mathematical Treatment of Measurement Results Explain the dimensional analysis (factor label) approach to mathematical calculations involving quantities Use dimensional analysis to carry out unit conversions for a given property and computations involving two or more properties Mathematical Treatment of Measurement Results A quantity of interest may not be easy (or even possible) to measure directly but instead must be calculated from other directly measured properties and appropriate mathematical relationships. The mathematical approach we will be using is known as dimensional analysis. Dimensional analysis is based on the premise that the units of quantities must be subjected to the same mathematical operations as their associated numbers. Conversion Factors and Dimensional Analysis A conversion factor or unit conversion factor is a ratio of two equivalent quantities expressed with different measurement units. Example: The lengths 2.54 centimeters and 1 inch are equivalent. 2.54 cm = 1 in. 2.54 cm 1 in. or 1 in. 2.54 cm Table 1.6 Common Conversion Factors Length Volume Mass 1 m = 1.0936 yd. 1 L = 1.0567 qt. 1 kg = 2.2046 lb 1 in. = 2.54 cm (exact) 1 qt. = 0.94635 L 1 lb = 453.59 g 1 km = 0.62137 mi 1 ft3 = 28.317 L 1 (avoirdupois) oz = 28.349 g 1 mi = 1609.3 m 1 tbsp = 14.1787 mL 1 (troy) oz = 31.103 g Conversion Factors and Dimensional Analysis We must use the form of the conversion factors that results in the original unit canceling out, leaving only the sought unit. Example: Convert 34 inches to centimeters. 2.54 cm 34 in. × = 86 cm 1 in. Conversion of Temperature Units Temperature refers to the hotness or coldness of a substance. Celsius scale Water freezes at 0 °C. Water boils at 100 °C. Fahrenheit scale Water freezes at 32 °F. Water boils at 212 °F. 100 °C covers the same temperature interval as 180 °F. Conversion of Temperature Units The SI unit of temperature is the kelvin (K). Unlike the Celsius and Fahrenheit scales, the kelvin scale is an absolute temperature scale. Zero kelvin corresponds to the lowest temperature that can theoretically be achieved. Kelvin scale Water freezes at 273.15 K. Water boils at 373.15 K. 100 °C covers the same temperature interval as 100 K. Mathematical Relationships Between Temperature Scales Fahrenheit and Celsius 9 F o  o To F =  o × To C  + 32 C 5 C  Kelvin and Celsius TK  To C + 273.15 Figure 1.28 The Fahrenheit, Celsius, and kelvin temperature scales are compared. Exercise 55 This OpenStax ancillary resource is © Rice University under a CC-BY 4.0 International license; it may be reproduced or modified but must be attributed to OpenStax, Rice University and any changes must be noted.