Week 2 Homework from Chemistry
1.1 Matter and Its Classification 5
Phosphorus Copper Bromine Nickel Lead
Gold Carbon Aluminum Sulfur Tin
FIGURE 1.4 Some elements. Which of these are metals? (Phosphorus, Bromine, Nickel, Lead, Aluminum, Sulfer & Tin): ©McGraw-Hill Education/Stephen Frisch; (Copper): ©Jim Birk; (Gold): ©Digital Vision/Getty Images; (Carbon): ©Photodisc/Getty Images.
Carbon Magnesium
Sulfur Aluminum
EXAMPLE 1.1 Metals and Nonmetals
Which of the elements pictured are metals? Why do you think so?
Solution: Notice that three of the elements—iron, aluminum, and magnesium—have a luster; that is, they shine. They are metals. If you could handle and test the substances, you could use other properties, such as electrical conductivity, to distinguish between metals and nonmetals.
Consider This 1.1 What if you were given element properties and had to determine the identity of the element from a list of possibilities? Suppose an element is rather dull in ap- pearance, a poor conductor of electricity, and a gas at room temperature. Would this element be zinc, platinum, or chlorine?
Practice Problem 1.1 Identify the nonmetals in Figure 1.4. Explain the characteristics you considered in making your decision.
Further Practice: Questions 1.31 and 1.32 at the end of the chapter
IronIron
(Iron): ©Sinclair Stammers/Science Source; (Carbon): ©Photodisc/Getty Images; (Magnesium, Sulfur & Aluminum): ©McGraw-Hill Education/ Stephen Frisch.
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To avoid having to write out the name of an element every time we refer to it, we use a system of symbols. An element symbol is a shorthand version of an element’s longer name. Often, the symbol is one or two letters of the element’s name (C for carbon, He for helium, Li for lithium). The first letter is uppercase, and the second letter, if present, is lowercase. When the names of two elements start with the same two first letters (magnesium and manganese, for example), the symbol uses the first letter and a later letter to distinguish them (Mg for magnesium, Mn for manganese).
For a few elements, the symbols are based on their Latin names or on names from other languages. These are listed in Table 1.1. Some recently synthesized ele- ments have been named for famous scientists or locations. See the periodic table and list of elements on the first page of this book for a complete list of the modern names and symbols.
To become familiar with the periodic table, you should learn the names and symbols for the first 36 elements, as well as the symbols for silver, tin, gold, mercury, and lead. Your instructor may ask you to learn others.
TABLE 1.1 Symbols of Selected Elements
English Name
Original Name
Symbol
English Name
Original Name
Symbol
copper cuprum Cu potassium kalium K
gold aurum Au silver argentum Ag
iron ferrum Fe sodium natrium Na
lead plumbum Pb tin stannum Sn
mercury hydrargyrum Hg tungsten wolfram W
Compounds A compound, sometimes called a chemical compound, is a pure substance composed of two or more elements combined in definite proportions. A compound has properties different from those of its component elements. For example, iron pyrite can be broken down into its component elements, iron and sulfur, but its characteristics are different from both (Figure 1.5). Anna and Bill
Iron pyrite
IronSulfur
FIGURE 1.5 Iron pyrite is composed of the elements iron and sulfur. Iron is magnetic and can be separated from sulfur when the two exist as elements mixed together. Iron pyrite, a compound of iron and sulfur, is not magnetic. (left): ©McGraw-Hill Education/Doug Sherman; (right): ©McGraw-Hill Education/ Stephen Frisch
EXAMPLE 1.2 Element Symbols
Potassium is a soft, silver-colored metal that reacts vigorously with water. Write the symbol for the element potassium.
Solution: The symbol for potassium is K. In the periodic table, potassium is element 19 in group (column) IA (1) of the periodic table.
Consider This 1.2 What if you instinctively identified the element symbol as P or Po? Why are these symbols incorrect for potassium?
Practice Problem 1.2 (a) Lead is a soft, dull, silver-colored metal. Write the symbol for the element
lead. (b) The symbol for a common element used to make jewelry is Ag. What is the
name of this element?
Further Practice: Questions 1.39 and 1.40 at the end of the chapter
1.1 Matter and Its Classification 7
saw many compounds during their trek around campus. These, and all com- pounds, can be chemically separated into their component elements. Sand is a compound of silicon and oxygen. Water, as discussed earlier, is composed of hydrogen and oxygen. The cheese on their pizza contains many complex com- pounds, but each of the compounds contains carbon, hydrogen, oxygen, nitrogen, and a few other elements.
Chemists represent compounds with formulas based on the symbols for the elements that are combined in the compound. (Chemical formulas are not the same as the mathematical formulas that may be familiar to you, such as A = πr2 for the area of a circle.) A chemical formula describes the composition of a compound, using the symbols for the elements that make up the compound. Sub- script numbers show the relative proportions of the elements in the compound. If no subscript number is given for an element in a formula, then you may assume that the element has a relative proportion of one. For example, water is known to consist of one unit of oxygen and two units of hydrogen. This compound is rep- resented by the formula H2O. Sodium chloride, the chemical compound com- monly called table salt, contains equal portions of the elements sodium and chlorine. Its formula is therefore NaCl. We will discuss formulas in detail in Chapter 3.
Mixtures Some forms of matter, such as pencil lead, do not have the same compo- sition in every sample. (Pencil lead isn’t the element lead. It is a mixture of graphite and clay.) A mixture consists of two or more elements or compounds. It is possible to separate mixtures into their component pure substances. The separation can be done physically, using procedures such as grinding, dissolving, or filtering. Chemi- cal processes are not needed to separate mixtures.
We can illustrate the difference between pure substances and mixtures by look- ing at salt water. Water that has been purified is a pure substance that is composed of hydrogen and oxygen, always in the same proportions. Salt water, on the other hand, is water mixed with salt and many other substances in varying proportions. For example, the Great Salt Lake in Utah is approximately 10% salt, while the Dead Sea is about 30% salt. In either case, we can readily separate salt from water by evaporating the water (Figure 1.6).
Mixtures differ in uniformity of composition. A homogeneous mixture has a uniform composition throughout and is often called a solution. Most solutions that we commonly encounter are composed of compounds dissolved in water. They are often clear. For example, a well-mixed sample of salt water prepared in a kitchen is uniform in appearance. The salt dissolved in it is invisible. Furthermore, any microscopically small portion of the sample would have the same composition as any other. The particles in the mixture might not be arranged in exactly the same pattern, but each sample, regardless of size, would have the same components in the same proportions.
A mixture that is not uniform throughout—a mixture of salt and pepper, for instance—is a heterogeneous mixture. Different samples have their components present in different proportions. Which of the things that Bill and Anna had for lunch is a homogeneous mixture? Which is heterogeneous? How about your own lunch? How can you tell?
We have considered a number of classes and subclasses of matter: mixtures, homogeneous mixtures, heterogeneous mixtures, pure substances, compounds, elements, metals, and nonmetals. A method for classifying matter into these categories is outlined in Figure 1.7. Note in the figure that yes or no answers to several questions distinguish one type of matter from another. First, we ask if the material can be separated physically. If so, then it is a mixture. If not, it must be a pure substance. If this substance can be decomposed (broken down into simpler substances) by chemical reactions, it is a compound. If it cannot, it is an element.
FIGURE 1.6 To collect salt, water is diverted into large ponds. The water evaporates, leaving solid salt behind. ©Science Photo Library/Alamy Stock Photo
Graphite leaves a mark similar to that made by dragging a rod of lead along a surface, so it was called lead. A hardness number indicates the relative amounts of graphite and clay in a pencil lead. A number 2 pencil is fairly soft, while a number 6 pencil is quite hard. Which has more graphite?
Not all solutions are liquids. For example, consider air that has been filtered to remove suspended solid particles. Filtered air is a gaseous solution containing a mixture of primarily oxygen and nitrogen. Solid solutions also exist and are called alloys. For example, brass is a solution of zinc and copper.
8 Chapter 1 Matter and Energy
FIGURE 1.7 We can classify matter by answering the short series of questions in this flowchart.
Can it be physically separated?
Matter
Can it be decomposed chemically?
Pure substance
Is it homogeneous?
Mixture
No Yes
Element Compound
Does it conduct electricity?
Nonmetal Metal Graphite
No Yes
Heterogeneous mixture
Homogeneous mixture
No YesNo Yes
or
EXAMPLE 1.3 Elements, Compounds, and Mixtures
Which of the following pictures represent pure substances?
(fountain): ©vora/iStock/Getty Images; (pizza): ©Kevin Sanchez/Cole Group/Getty Images; (coin): ©Randy Allbritton/Getty Images; (balloons): ©Jules Frazier/Getty Images (drink): ©Brian Moeskau/Moeskau Photography
1.1 Matter and Its Classification 9
Solution: The copper on the outside of the coin and the helium inside the balloons are pure substances. (However, the helium and balloons considered together provide an example of a mixture.)
Consider This 1.3 Why isn’t the water in the fountain considered a pure substance?
Practice Problem 1.3 Which of the pictures represent mixtures? Which are heterogeneous? Which are homogeneous?
Further Practice: Questions 1.45 and 1.46 at the end of the chapter
FIGURE 1.8 A copper pipe consists of a regular array of copper atoms. ©Thinkstock/Getty Images
Copper atom Helium atom
FIGURE 1.9 Helium atoms are present inside the balloon. ©Jules Frazier/Getty Images
Representations of Matter Chemists and other scientists view the world on several different levels. So far we have considered matter on a macroscopic scale. That is, we’ve discussed matter and phenomena we can see with our eyes. But simple observation is limited. Sometimes we cannot classify things merely by looking at them as Anna and Bill did. What do we do then? Chemists try to make sense of the structure of matter and its behavior on a scale that is much, much smaller than what we can see with our eyes.
Consider the copper pipe at the construction site, for example. If we could enlarge the tiniest unit that makes up the pipe, what would we see? Experimental evidence tells us copper is made up of discrete, spherical entities that all appear to be identical (Figure 1.8). Chemists identify these entities as atoms. An atom is the smallest unit of an element that has the chemical properties of that element. For example, we can imagine the helium inside a balloon as many, many atoms of helium, which we represent symbolically as He. In Figure 1.9, each sphere repre- sents a single helium atom. Similarly, if we could magnify the structure of water, we would find two small hydrogen atoms bound separately to a single larger oxy- gen atom (Figure 1.10). Such a combination of elemental units is a molecule. Molecules are made up of two or more atoms bound together in a discrete arrange- ment. Several molecules of water, H2O, are shown in Figure 1.10, where the cen- tral red sphere represents an oxygen atom and the two smaller, white spheres stand for hydrogen atoms. (Some compounds do not exist as molecules. We will discuss them in Chapter 3.)
Although chemists generally use color coding to distinguish between atoms of different elements in representations, the atoms themselves do not have colors. Macroscopic samples of matter may have color, but these colors do not usually match those used to represent atoms. In accurate representations, the sizes of the spheres change to reflect the relative differences in the sizes of atoms of different elements.
10 Chapter 1 Matter and Energy
In addition to molecules of compounds, molecules can also be formed by the combination of atoms of only one element. For example, as shown in Figure 1.11, the oxygen we breathe consists of molecules of two oxygen atoms joined together. We represent oxygen molecules symbolically as O2.
Chemists use many different ways to represent matter. Some for water are shown in Figure 1.12. Element symbols with subscripts represent a ratio of ele- ments in a compound. One example is Figure 1.12B. To describe how the atoms are attached to one another, chemists often use lines and element symbols as shown in Figure 1.12C. In Figure 1.12D, spheres represent the atoms, and sticks show how they are connected. Figure 1.12E represents how the atoms fit together and their relative sizes. Macroscopic, molecular-level, and symbolic representations like these all have their advantages, and sometimes one is more convenient than another. You’ll use them all as you progress through this course.
FIGURE 1.10 Molecules containing hydrogen atoms and oxygen atoms make up the water in the fountain. Note: The molecular-level image does not include dissolved matter that is present in the fountain water. ©Glowimages/Getty Images
Oxygen atom
Hydrogen atom
FIGURE 1.11 Oxygen molecules are made up of two interconnected oxygen atoms and are represented symbolically as O2.
FIGURE 1.12 Different ways of representing water: (A) macroscopic, (B and C) symbolic, and (D and E) molecular. ©Royalty-Free/Corbis
A
H2OB
C OH H
D
E
EXAMPLE 1.4 Representations of Matter
(a) Which of these images best represents a mixture of elements?
(b) If image A represents nitrogen, write its formula.
A B C
D E
1.1 Matter and Its Classification 11
States of Matter Earlier we considered the classification of matter based on composition. Let’s look at a different way to classify matter: by its physical state. A physical state is a form that matter can take. The three most familiar to us are solid, liquid, and gas. Some substances, including some of those Anna and Bill observed, can be found in all three states under more or less ordinary conditions. Water, for example, can be a solid (ice), a liquid (flowing water), or a gas (water vapor) at environmental temperatures.
Other substances require extreme conditions to change from one state to an- other. For example, while carbon dioxide is a gas under normal conditions, it be- comes a solid, called dry ice, at very low temperatures (Figure 1.13).
How do we know if a substance is in the solid, liquid, or gaseous state? Each state has characteristics that we can observe with our eyes and characteristics that are detectable or measurable at the molecular level. These characteristics are summarized in Table 1.2.
Solution: (a) There are two mixtures represented in the images. Since the spheres (repre-
senting atoms) in image C have different colors and sizes, we can conclude that image C is a mixture of two elements. Image D is also a mixture, but it is a mixture of an element and a compound.
(b) The formula of the substance represented in image A is N2. Note that two atoms are connected in the molecule.
Consider This 1.4 Which combination of entities shown in the images would represent a mixture of compounds?
Practice Problem 1.4 (a) Which of the images represents an element that exists as a molecule? (b) If image E represents a compound of oxygen (red) and sulfur (yellow), what
is its formula? (Write the symbol for sulfur first.)
Further Practice: Questions 1.51 and 1.52 at the end of the chapter
FIGURE 1.13 Dry ice is the solid state of carbon dioxide. It converts from a gas to a solid at a very low temperature. ©Charles D. Winters/Science Source
ANIMATION: Three States of Matter
TABLE 1.2 Characteristics of the Physical States of Matter
Solid Liquid Gas
fixed shape shape of container (may or may not fill it)
shape of container (fills it)
its own volume its own volume volume of container
no volume change under pressure
slight volume change under pressure
large volume change under pressure
particles are fixed in place and tend to be in a regular (crystalline) array
particles are randomly arranged and free to move about until they bump into one another
particles are widely separated and move independently of one another
12 Chapter 1 Matter and Energy
Liquid iron
Solid iron
FIGURE 1.14 The liquid and solid states of iron. Notice that the liquid atoms are randomly arranged and are free to move around each other. In the solid, the atoms are fixed in a regular array. ©Arthur S. Aubry/Photodisc/Getty Images
A solid has a fixed shape that is not related to the shape of the container holding it. When you place an iron pipe in a box, the pipe does not change shape. Some solids can be made to change shape if enough force is applied. However, if you try to squeeze a solid to make it smaller, you’ll fail. A solid cannot be compressed be- cause its particles are arranged in a tightly packed, highly ordered structure that does not include much free space into which they might be squeezed. Note the closely packed particles in the solid state of iron shown in Figure 1.14.
A liquid is different from a solid in that it has no fixed shape. It takes the shape of the filled portion of its container, and it can be poured. Although they touch, the particles in a liquid are not arranged in ordered structures like those in a solid; they are free to move past one another. A liquid can be compressed slightly because its particles have a little free space between them. Note the differences between the liquid and solid states of iron shown in Figure 1.14.
A gas has no fixed shape; it adopts the shape of its container, expanding to fill the available space completely. A gas is easily compressed. When squeezed, gases can undergo large changes in volume. The particles of a gas are widely separated with much empty space between them. When a gas is compressed, the amount of space between the particles is reduced. This happens when pressure is applied, such as when a bicycle tire is filled with air, as shown in Figure 1.15. Another character- istic of gases is that they move through space quickly. When Bill and Anna smelled the pizza they had for lunch, they were detecting particles that migrated as gases from the source of the food to their noses. When gases cool sufficiently, they be- come liquids or even solids. This occurs, for example, when water vapor in the air liquefies on the surface of a cold glass. Note the differences between the liquid and gaseous states of water shown in Figure 1.16.
It is often convenient to show the physical state of a substance when represent- ing it symbolically. For example, solid, liquid, and gaseous water can be represented as H2O(s), H2O(l), and H2O(g), respectively. The symbol (aq) represents an aque- ous solution, a solution in which a substance is dissolved in water. A salt and water solution, for instance, can be written as NaCl(aq). These symbols for physical state are listed in Table 1.3.
Does NaCl(aq) represent a pure substance or a mixture?
Some solids, called amorphous solids, do not have the high order that most crystalline solids have.
1.2 Physical and Chemical Changes and Properties of Matter 13
1.2 Physical and Chemical Changes and Properties of Matter
Bill and Anna observed some of the properties of matter, including changes in mat- ter. Observations can be either qualitative, based on some quality of the matter; or quantitative, based on a numerical value. When making qualitative observations, Bill and Anna described color, shape, texture, shininess, and physical state. Quanti- tative observations are different. They are numbers or measurements, and they must be carefully made and carefully reported.
Since quantitative data used to describe matter can involve both very large and very small numbers, it is often useful to express such numbers in scientific or exponential notation. Math Toolbox 1.1 (located at the end of this chapter) provides a review of this notation. In addition, it is necessary to express numbers in such a way as to indicate how accurately the value is known and how precisely it has been measured. The use of sig- nificant figures to properly express numerical values is presented in Math Toolbox 1.2.
Physical Properties When reporting qualitative data, we can classify properties as either physical or chemical. When Bill and Anna observed the color, shape, texture, shininess, and physical state of things around them, they were noting their physical properties.
High pressure Compressed air
Low pressure Normal air
Water vapor in humid air
Condensed water on glass
N2
O2
FIGURE 1.15 At the same temperature, a gas under high pressure has particles closer together than at low pressure. Notice that the composition (1 O2:4 N2) does not change with an increase in pressure.
FIGURE 1.16 Water condenses from a gas to a liquid on a cold surface. Air molecules (e.g., oxygen and nitrogen) are not shown. ©Brian Moeskau/Moeskau Photography
TABLE 1.3 Symbols for Physical State
Physical State Symbol Example (bromine)
solid (s) Br2(s)
liquid (l) Br2(l)
gas (g) Br2(g)
aqueous (dissolved in water) (aq) Br2(aq)
MATH TOOLBOX
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14 Chapter 1 Matter and Energy
A physical property is a characteristic that we can observe or measure without changing the composition of a substance. Other examples of physical properties are odor, taste, hardness, mass, volume, density, magnetism, conductivity, and the tem- peratures at which a substance changes from one physical state to another. (Later in this section we’ll discuss chemical properties, which include reactivity and flam- mability.) Let’s take a close look at mass, volume, density, and temperature. These four properties are quantitative; they involve numerical values.
Mass Recall that mass is a measure of the quantity of matter. We usually measure the mass of an object by weighing it on a balance. In chemistry, masses are often reported in units of grams (g). Large masses, like people or elephants, may be re- ported in units of kilograms (kg); and small masses, such as salt crystals or impuri- ties in water, may be reported in units of milligrams (mg) or micrograms (μg), as shown in Figure 1.17. (Math Toolbox 1.3 summarizes the relationships among units such as these.) Sometimes the mass of something is reported in grams, but we might want to know the mass in another mass unit such as milligrams or kilograms. We can easily convert a measurement from one unit to another if we know the relationship between the units. Tables 1.4 and 1.5 summarize common relationships between metric and English units. Example 1.5 shows how to convert between mass units. (See Math Toolbox 1.3 for more information on unit conversions.)
FIGURE 1.17 A salt crystal has a mass of about 50 mg, while a person has a mass of about 70 kg. (left): ©Jim Birk; (right): ©Doug Menuez/Getty Images
Mass: 50 mg, 0.05 g, or 5 × 10−5 kg Mass: 7 × 107 mg, 7 × 104 g, or 70 kg
TABLE 1.4 Metric Conversions
Prefix Factor Symbol
giga 109 G
mega 106 M
kilo 103 k
deci 10−1 d
centi 10−2 c
milli 10−3 m
micro 10−6 μ
nano 10−9 n
pico 10−12 p
TABLE 1.5 Some English-Metric Conversions
English Unit Metric Unit
1 lb = 16 oz 453.6 g
1 in 2.54 cm (exactly)
1 yd 0.9144 m
1 mi 1.609 km 1 fluid oz 29.57 mL
1 qt 0.9464 L 1 gal 3.785 L 1 ft3 28.32 L
EXAMPLE 1.5 Units of Mass
Anna and Bill notice that there are 50.0 mg of sodium in the soda they bought to go with their lunch. How many grams of sodium are present in the can of soda? How many pounds?
Solution: One way to solve this problem uses the dimensional-analysis approach. Consult Math Toolbox 1.3 for details. The general approach to solving the first part of the problem can be summarized by the following diagram:
Mass in milligrams Mass in grams
?
MATH TOOLBOX
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MATH TOOLBOX
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1.2 Physical and Chemical Changes and Properties of Matter 15
The mass in milligrams has to be converted to the mass in grams. We need to find a relationship between these two quantities: 10−3 g = 1 mg (obtained from Table 1.4). Another way of expressing this relationship is 1 g = 1000 mg. Using this relationship, we get the following conversion:
Mass in milligrams Mass in grams
1000 mg = 1 g
We use the equivalence to set up possible conversion ratios:
1000 mg and 1 g
1 g 1000 mg
To convert milligrams to grams, we can multiply 50.0 mg by the ratio (conversion factor) that will allow like units to cancel:
Mass in g = 50.0 mg 1000 mg
1 g 0.0500 g=×
Note that the milligram units cancel to leave the appropriate unit of grams. Also notice that the answer is reported to three significant figures because the mea- sured quantity (50.0 mg) is reported to three significant figures and the other numbers in the calculation (1 g and 1000 mg) are exact quantities. (Consult Math Toolbox 1.2 for details about significant figures.) For convenience, we could re- port the answer in scientific notation: 5.00 × 10−2. (See Math Toolbox 1.1 for a discussion of scientific notation.)
The second part of the question asks you to convert milligrams to pounds:
Mass in milligrams
Weight in pounds
?
There isn’t a direct relationship between milligrams and pounds listed in Tables 1.4 and 1.5. However, Table 1.5 lists a relationship between pounds and grams: 1 lb = 453.6 g. We can convert the grams we found in the first part of this example to pounds using the relationship summarized in the following diagram:
Weight in pounds
453.6 g = 1 lb Mass in grams
The ratios for converting between grams and pounds are
1 lb and453.6 g
453.6 g 1 lb
To convert grams to pounds, we multiply 0.0500 g by the ratio (conversion factor) that will allow like units to cancel:
Weight in pounds = 0.0500 g × ×1 lb 453.6 g
= 1.10 10−4 lb
Does this answer make sense? Yes, it does. There are a lot of grams (453.6) in a pound, so we would expect the answer to be very small.
MATH TOOLBOX
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16 Chapter 1 Matter and Energy
Without a single conversion from milligrams to pounds, the problem we just solved involves multiple steps:
Mass in milligrams Mass in grams
1000 mg = 1 g Weight in pounds
453.6 g = 1 lb
The sequence of steps can be summarized as:
1 g Weight in pounds = 50.0 mg × × 1 lb 453.6 g
= 1.10 × 10−4 lb 1000 mg
Consider This 1.5 Why wouldn’t an answer of 22.7 lb for the second part of the question make sense?
Practice Problem 1.5 Anna and Bill see an aluminum recycling truck pass by on their way to class. If there are 765 lb of aluminum in the truck, how many grams are there? How many kilograms?
Further Practice: Questions 1.67 and 1.68 at the end of the chapter Student data indicate you may struggle with unit conversions. Access the SmartBook to view additional Learning Resources on this topic.
Student Hot Spot
Volume Volume is the amount of space a substance occupies. We can determine the volume of a rectangular solid such as a cube by measuring its length, width, and height and then multiplying them. For example, the volume of a cube that is 2.0 centimeters (cm) on each side is 8.0 cubic centimeters (cm3): Volume of a cube = length × width × height Volume = 2.0 cm × 2.0 cm × 2.0 cm = 8.0 cm3
Notice that the units are cm3, consistent with a three-dimensional quantity. One cubic centimeter is equal to 1 mL (1 cm3 = 1 mL), so the volume of 8.0 cm3 could also be reported as 8.0 mL.
The volumes of liquids are usually measured in units of liters (L) or milliliters (mL), as shown in Figure 1.18. Larger volumes, such as big bottles of soda, are usu- ally reported in liters. A 1-L bottle of soda contains 1000 mL. Example 1.6 shows how to convert between volume units.
If you need to determine the volume of a sphere, the relationship between volume and radius is 4
3 πr3V = .
FIGURE 1.18 Some 500-mL, 1-L, and 250-mL containers. ©Brian Moeskau/Moeskau Photography
1.2 Physical and Chemical Changes and Properties of Matter 17
EXAMPLE 1.6 Units of Volume
For lunch, Anna and Bill had 12-ounce (oz) cans of soda. What is the volume of a 12.0-oz can of soda in units of milliliters? What is its volume in units of liters?
Solution: To solve this problem using the dimensional-analysis approach (see Math Toolbox 1.3), we determine if there is a relationship between fluid ounces and milliliters:
Volume in ounces
Volume in milliliters
?
To convert fluid ounces to milliliters, we use the following relationship from Table 1.5: 1 oz = 29.57 mL.
Volume in ounces
Volume in milliliters
1 oz = 29.57 mL
We use the equivalence to set up possible conversion ratios:
1 oz 29.57 mL
29.57 mL 1 ozand
To convert ounces to milliliters, we can multiply 12.0 oz by the ratio (conversion factor) that will allow like units to cancel:
1 oz 29.57 mLVolume in milliliters = 12.0 oz × = 355 mL
The answer is reported to three significant figures, because the quantity we’re given (12.0 oz) has three significant figures. Consult Math Toolbox 1.2 for details.
The second part of this problem asks you to convert milliliters to liters:
Volume in milliliters
Volume in liters
?
To convert volume in milliliters to volume in liters, we use the following relation- ship from Table 1.4: 1 mL = 10−3 L or 1000 mL = 1 L.
Volume in milliliters
Volume in liters
1000 mL = 1 L
This is a unit conversion similar to the conversion we just did for mass. The ratios for converting between milliliters and liters are
1000 mL 1 L
1 L 1000 mLand
To convert from milliliters to liters, we can multiply 355 mL by the conversion factor that allows like units to cancel:
1000 mL 1 LVolume in L = 355 mL × = 0.355 L
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18 Chapter 1 Matter and Energy
FIGURE 1.19 The densities of antifreeze, corn oil, dish detergent, maple syrup, shampoo, and water in g/mL are 1.13, 0.93, 1.03, 1.32, 1.01, and 1.00, respectively. Which layer is which substance? ©Richard Megna/Fundamental Photographs
Au Al
FIGURE 1.20 Gold (Au) has a greater density than aluminum (Al) because gold has a greater mass per unit volume.
Without a single conversion from ounces to liters, the problem we just solved involves multiple steps:
Volume in ounces
Volume in milliliters
1 oz = 29.57 mL Volume in liters
1000 mL = 1 L
The sequence of steps can be summarized as:
1000 mL 1 L
1 oz 29.57 mLVolume in liters = 12.0 oz × × = 0.355 L
Consider This 1.6 What if Anna and Bill shared a 1-L bottle of water over lunch? How many ounces is this?
Practice Problem 1.6 Anna and Bill saw some balloons outside the bookstore. The volume of gas inside one of the helium balloons was 4.60 L. What is the volume of gas in units of milliliters? In units of cubic centimeters? In units of gallons (4 qt = 1 gal)?
Further Practice: Questions 1.71 and 1.72 at the end of the chapter
TABLE 1.6 Densities of Some Common Substances
Substance Physical State Density (g/mL)*
helium gas 0.000178 oxygen gas 0.00143 cooking oil liquid 0.92 water liquid 1.00 mercury liquid 13.6 gold solid 19.3 copper solid 8.92 zinc solid 7.14 ice solid 0.92
*At room temperature and at normal atmospheric pressure, except gases at 0 degrees Celsius (°C) and water at 4°C.
Density The density of an object is the ratio of its mass to its volume. While mass and volume both depend on the size of the object or sample, density does not. Den- sity is an unvarying property of a substance no matter how much of it is present, as long as temperature and pressure are constant. For example, the density of water at 4°C is 1.00 g/mL. It doesn't matter if we have 10 mL or 10 L; the ratio of mass to volume would be the same, 1.00 g/mL. However, if the temperature increases, the water would expand to a larger volume, while the mass stays the same. The density of liquid water would decrease as the temperature increases. The densities of water and a few other substances are listed in Table 1.6.
As Anna and Bill noted when they observed the fountain, a copper coin sinks in water. It sinks because copper (and the other metals in a penny) have a greater den- sity than water. Conversely, air bubbles, just like other gases, rise to the top of water because gases are less dense than liquids. Oil floats on water for this same reason.
The density column in Figure 1.19 shows a variety of liquids with different densities. Which liquid has the greatest density? Which is the least dense?
If we compare equal volumes of two different substances, such as aluminum and gold, as shown in Figure 1.20, the substance with the greater mass has the greater
1.2 Physical and Chemical Changes and Properties of Matter 19
density. How, though, can we compare densities if we do not have equal volumes? The mathematical relationship of mass, volume, and density reveals the answer:
Density = mass volume
For example, a 1.0-cm3 sample of copper has a mass of 8.9 g. An 8.0-cm3 sample of copper has a mass of 71 g. A 27-cm3 sample of copper has a mass of 240 g. In all these samples (Figure 1.21), the mass of copper divided by its volume is 8.9 g/cm3. This is the density of copper.
If we know the mass and volume of an object, we can determine its density by substituting directly into the density equation. For example, suppose we have a cube of an unknown metal with a mass of 178 g and an edge length of 2.92 cm. The vol- ume of the cube is 24.9 cm3:
Volume = 2.92 cm × 2.92 cm × 2.92 cm = 24.9 cm3
The density can then be calculated by taking the ratio of the mass to volume:
Density = mass volume
Density = 178 g
24.9 cm3 = 7.15 g/cm3
Consulting Table 1.6, we see that the unknown metal could be zinc. Additionally, if we know the density of a substance and its mass in our sample, we
can determine its volume. For example, suppose we want to know the volume occupied by 100 g of copper. Should the volume be greater than or less than 100 cm3? There are many approaches to this problem. One way is to rearrange the density equation to solve for volume. Another way is to solve for the unknown volume in a set of equivalent ratios because density is a ratio of mass and volume that is constant for a given sub- stance at a particular temperature. Both of these methods are shown in Example 1.7.
8.9 g 1.0 cm3
240 g 27.0 cm3
71 g 8.0 cm3
FIGURE 1.21 The density of copper is 8.9 g/cm3. All three samples have the same ratio of mass to volume.
MATH TOOLBOX
1.3
These samples of metals have the same mass. Which has the greater density?
©Jim Birk
EXAMPLE 1.7 Density, Volume, and Mass
What is the volume of 100.0 g of copper? The density of copper is 8.9 g/cm3.
Solution: We need to carry out the following conversion:
Volume in millilitersMass in grams
?
The relationship between mass and volume is given by density:
Volume in millilitersMass in grams
Density = mass volume
First, we rearrange the density equation to get volume on one side by itself. This manipulation involves cross multiplication, which is described in Math Tool- box 1.3 (Ratio Approach). In the expression for density there is an implied 1:
Density = mass volume
= mass volume
Density 1
A can of diet cola floats in water, but a can of regular cola sinks. Suggest a reason why. How can you use this information to quickly select your preferred type of soft drink from a cooler filled with ice water at a party?
©Brian Moeskau/Moeskau Photography
ANIMATION: Density of Liquids and Solids
20 Chapter 1 Matter and Energy
Cross multiplying this density expression, we get: Density × volume = mass × 1
Since we are trying to find the volume, we want to isolate it on one side of the equation. We can do this by dividing both sides by the density. (We’ll also drop the “× 1” because any quantity times 1 is that quantity.)
= mass density
Density × volume Density
Now we have an expression that solves for the volume:
Volume = mass density
Then, we substitute the known values of mass and density into the equation and solve for the value of volume:
Volume 100.0 g 8.9 g/cm3
11 cm3= =
In a second approach to this problem, consider that since the density of copper is always the same, the ratio of mass to volume is the same for both what we know and what we don’t:
8.9 g 1 cm3
100.0 g= x cm3
Cross multiply to solve for x:
x cm3 = 8.9 g
= 11 cm3(1 cm3) × (100.0 g )
In both approaches, the gram units cancel to give the expected volume unit of cm3. There is yet another approach to solving this problem that involves using density
as a conversion factor:
Volume = 100.0 g × = 11 cm3 8.9 g 1 cm3
Does the answer make sense? Yes. The density tells us that 8.9 g of copper occupy a volume of 1 cm3. The mass given, 100.0 g, is over 10 times greater than 8.9 so we would expect it to occupy a volume that is over 10 times greater than 1 cm3.
Consider This 1.7 What if the 100.0-g piece of metal were gold? Would you expect its volume to be greater or less than 11 cm3? How much greater or less?
Practice Problem 1.7 Solve the following problems. (a) The density of pure gold is 19.3 g/cm3. What is the volume of 1.00 g of pure gold? (b) 14-Carat gold is a homogeneous mixture of metals containing 58% gold by
mass. The other 42% is a mixture of silver and copper. Silver and copper are both less dense than gold. Which of the following could be the mass of 1.00 cm3 of 14-carat gold: 16.0 g, 19.3 g, or 23.0 g?
Further Practice: Questions 1.77 and 1.78 at the end of the chapter
The density 8.9 g/cm3 can be expressed in fractional form as: 8.9 g 1 cm3
Water is unique among liquids because its solid form (ice) floats on its liquid form. This results from the relatively open structure adopted by water molecules in the solid state. What would happen to fish during the winter if water were like other substances that sink in their solid form?
Why do substances have different densities? Gases, in general, have very low densi- ties because gas particles spread out and occupy large volumes. Metals tend to have high densities because their atoms pack together efficiently. Because ice floats on water, we
1.2 Physical and Chemical Changes and Properties of Matter 21
can infer that water in its solid form must have a lesser density than water in its liquid form. Example 1.8 shows how to use molecular pictures to predict relative densities.
Ice Liquid water
Helium
Carbon dioxide
Helium
Carbon dioxide
EXAMPLE 1.8 Explanations for Density
How do the molecular diagrams of ice and water help explain why ice is less dense than water?
Solution: In ice, the H2O molecules have more space between them than in liquid water. The total volume occupied by a given number of molecules is greater in ice. Because density is a ratio of mass to volume, the larger volume accounts for the lower density.
Consider This 1.8 What if hexane, with a density of 0.659 g/cm3, were carefully added to a glass of ice water? Where would the hexane be relative to the ice and liquid water?
Practice Problem 1.8 Helium balloons rise in air, which is a mixture of oxygen and nitrogen molecules, so we know helium is less dense than air. Look at the molecular-level diagrams of helium and carbon dioxide. Predict whether a helium balloon rises or falls in an atmosphere of carbon dioxide.
Further Practice: Questions 1.81 and 1.82 at the end of the chapter
Temperature Bill and Anna weren’t happy with their lunches. The pizza was cold and their sodas were warm. When we make such comparisons, we are observing relative temperatures. Temperature is a measure of how hot or cold something is relative to some standard. We measure temperature with a thermometer.
In the United States, we often use the Fahrenheit scale to measure body tem- perature and air temperature. Fahrenheit is rarely used in science. Two other tem- perature scales are standard: the Celsius scale and the Kelvin scale. The relationships between the three temperature scales, Fahrenheit (°F), Celsius (°C), and Kelvin (K), are shown in Figure 1.22.
Another property of matter that is independent of sample size is the temperature at which the substance changes from one physical state to another. The boiling point is the temperature at which the liquid form of a substance changes to the gaseous form.
ANIMATION: Unique Properties of Water
Temperatures are written differently for the different scales. While Celsius and Fahrenheit use the superscript ° to indicate degrees, the Kelvin scale does not. The unit is written as K (the capital letter), but temperatures are measured in kelvins (lowercase).
22 Chapter 1 Matter and Energy
At the melting point, the substance changes from a solid to a liquid. Between these two temperatures, the substance is normally in its liquid state. For example, on the Celsius scale, the boiling point of water is 100°C. Water melts (or freezes, depending on its original state) at 0°C. On the Kelvin scale, these values are 373.15 K and 273.15 K, respectively. On the Fahrenheit scale, they are 212°F and 32°F, respectively.
There are no negative values on the Kelvin scale. It is an absolute temperature scale because its zero point is the lowest possible temperature observable in the uni- verse. This value is absolute zero, which is equivalent to −273.15°C. The tempera- ture increments on the Kelvin scale are the same as those on the Celsius scale. The difference in temperature between the boiling point of water and the freezing point of water is 100 in both the Celsius (100°C − 0°C) and Kelvin (373.15 K − 273.15 K) scales, while the difference is 180 on the Fahrenheit scale (212°F − 32°F). Because the temperature in kelvins is always 273.15 greater than the temperature in degrees Celsius, we can easily convert between them:
TK = T°C + 273.15 When converting between the Fahrenheit and Celsius scales, the calculation is
more complicated, because the degree increments are not equal: T°F = 1.8(T°C) + 32
The equation can be rearranged, solving for degrees Celsius:
T°C = 1.8
T°F − 32
FIGURE 1.22 The Fahrenheit, Celsius, and Kelvin temperature scales.
Liquid water boils/ water vapor condenses
212°F
Room temperature
Lowest possible temperature:
Ice melts/ liquid water freezes
Fahrenheit Celsius Kelvin
77°F
32°F
−460°F −273.15°C 0 K
100°C
25°C
0°C
373.15 K
298.15 K
273.15 K
EXAMPLE 1.9 Units of Temperature
The melting point of copper is 1083°C. Above what temperature, in kelvins and degrees Fahrenheit, is copper a liquid?
Solution: Copper becomes a liquid above its melting point. In units of kelvin this temperature is
TK = T°C + 273.15
We substitute the value of the Celsius temperature into the expression and solve for the temperature in kelvins:
TK = 1083 + 273.15 = 1356 K
1.2 Physical and Chemical Changes and Properties of Matter 23
Physical Changes A process that changes the physical properties of a substance without changing its chemical composition is a physical change. For example, we can change liquid water to water vapor by heating it. This change from a liquid to a gas, called boiling or vaporization, is a physical change, since both forms involve the same chemical substance, water (H2O).
To represent such changes, we can refine the symbolic representations we de- veloped for elements and for compounds. We write the chemical formula for the initial condition and composition of the matter we are considering, then an arrow, and finally the chemical formula for the final condition and composition. The arrow is used to show that a change has occurred and in which direction. Using this sym- bolism, the change of water from a liquid to a gas would be represented as
H2O(l)
reaction
reaction
resonance
equilibrium
H2O(g) The molecular and symbolic representations in Figure 1.23 show that the water
molecules do not themselves change, but their physical state does. All the processes that change water from one physical state into another are summarized in Figure 1.24.
Another example of a physical change is the separation of different substances in a mixture. For example, a magnet divides magnetic materials from nonmagnetic mate- rials without changing their identities, as shown earlier in Figure 1.5. A filter separates solid materials from liquid substances without changing either one chemically.
To convert degrees Celsius to degrees Fahrenheit, we use the equation: T°F = 1.8(T°C) + 32
Substituting the temperature in degrees Celsius, we get: T°F = 1.8(1083) + 32 = 1981°F
Consider This 1.9 What is the physical state of copper at 1000°C?
Practice Problem 1.9 (a) The boiling point of acetylene is −28.1°C. Below what temperature, in kel-
vins and degrees Fahrenheit, is acetylene a liquid? (b) The boiling point of helium is 4 K. Below what temperature, in degrees Cel-
sius, is helium a liquid? (c) Human body temperature is normally 98.6°F. What is this temperature on the
Celsius and Kelvin scales?
Further Practice: Questions 1.85 and 1.86 at the end of the chapter
The most common type of heating we encounter involves things we do in the kitchen. When we heat something up, we’re typically not changing the chemical composition. However, the process of cooking or baking often involves a combination of physical and chemical changes. Of course, if you were to overcook something and burn it, that would be a chemical change.
FIGURE 1.23 Molecular-level and symbolic representations of the evaporation of water.
H2O(l) H2O(g)
24 Chapter 1 Matter and Energy
Chemical Changes Remember the pennies in the fountain that Anna and Bill observed? Some were shiny and others looked dingy and brown. They might describe these less-shiny pennies as “tarnished.” The pennies have undergone a chemical change, a process in which one or more substances are converted into one or more new substances. When pennies tarnish, some of the copper and zinc metal atoms in them combine with oxygen, forming compounds called metal oxides. The compounds are chemi- cally different from either of the elements that formed them.
Suppose we clean a tarnished penny. Is the process a physical or a chemical change? It can be either. If you simply rub off the metal oxide coating with an eraser, the change is physical. Most penny collectors, however, prefer a chemical change that removes less metal. Rubbing ketchup on a penny is a great way to make it shiny. The vinegar in the ketchup reacts chemically with the metal oxides, freeing them from the surface of the penny. When the penny is rinsed, the result of the chemical change is easy to see.
Anna and Bill observed other examples of chemical change during their cam- pus walk. When gasoline-powered cars burn fuel, a chemical change occurs. The gasoline reacts with oxygen to form carbon dioxide and water vapor. This chemical change releases the energy that runs the car. Chemical changes that involve burning are often accompanied by the release of energy. Anna and Bill also observed vehi- cles that run on alternative fuels. In hydrogen-powered vehicles, the hydrogen fuel combines with oxygen to form water vapor—and to release a lot of energy. The molecular-level and symbolic representations for this chemical change are shown in Figure 1.25. A chemical change is often called a chemical reaction. What are some examples of chemical reactions that you can observe around you?
Chemical Properties The copper and zinc in a penny, the gasoline in a car, and the hydrogen in an alternative- fuel vehicle all share a common chemical property: They react with oxygen. However, they differ in how they react and what products they form. Only the latter two release sufficient energy rapidly enough to make their use as fuels possible.
FIGURE 1.24 The physical states of solid, liquid, and gas can all change into one another either directly or by going through two changes of state. The names of these processes are shown here next to arrows that designate the direction of the change.
Solid
Melting
Condensation Vaporization
Sublimation
Deposition
Freezing
Liquid
Gas
Many metals combine with oxygen to form a metal oxide compound at the surface of the metal. When this occurs with iron, we call it rust.
1.2 Physical and Chemical Changes and Properties of Matter 25
A chemical property of a substance is defined by what it is composed of and what chemical changes it can undergo. For example, let’s compare hydrogen and helium. Although they have similar physical properties (colorless gases, similar densities), their chemical properties are very different. While hydrogen reacts with many other elements and compounds, helium is considered inert (Figure 1.26). It has not yet been shown to react with any other element or compound. Other terms used to describe chemical properties include reactivity and flammability.
FIGURE 1.25 A chemical change occurs when the atoms in H2 and O2 rearrange to form H2O. 2H2(g) + O2(g) 2H2O(g)
O2 H2
H2O
FIGURE 1.26 (A) The Hindenburg was a giant, rigid balloon filled with hydrogen gas. In 1937, it was destroyed when its hydrogen caught fire. (B) Today, blimps are filled with helium, an inert gas that will not explode. (a): ©Bettmann/Getty Images; (b): ©David R. Frazier/Alamy Stock Photo
A B
EXAMPLE 1.10 Physical and Chemical Changes
Which of the following are physical changes and which are chemical changes? (a) evaporation (b) burning methane gas to form carbon dioxide and water (c) using a magnet to separate metal and plastic paperclips (d) rusting (the conversion of iron to iron oxide)
Solution: (a) Evaporation is a physical change because it involves only a change of state. (b) Burning methane gas is a chemical change because new substances form. (c) Separating components of a mixture is a physical change. (d) Rusting is a chemical change because a new substance forms.
Consider This 1.10 Is digesting the pizza Anna and Bill ate for a lunch a chemical or a physical change?
26 Chapter 1 Matter and Energy
Practice Problem 1.10 Which of the following are physical properties and which are chemical properties? (a) boiling point of ethanol (b) ability of propane to burn (c) tendency for silver to tarnish (d) density of aluminum
Further Practice: Questions 1.91 and 1.92 at the end of the chapter
Sometimes simple observation cannot tell us whether a change is chemical or physical. For example, bubbles appear when baking soda and vinegar mix. Bubbles also appear when water boils, but the changes that produce the bubbles are different in these two cases. Baking soda and vinegar release bubbles because a chemical change takes place. They react to form carbon dioxide gas. However, when we warm water in a pan on the stove, small bubbles rise due to the release of dissolved air (mostly oxygen and nitrogen gas) from the water (before the water starts to boil). This process is only a physical change. If we could look at the nitrogen and oxygen molecules, as shown in Figure 1.27A, we would see that they are the same whether they are dissolved in water or not. When these molecules are dissolved in water, they form a homogeneous mixture with it. Heating the water merely separates the oxygen and nitrogen molecules from the water molecules. If we continue to heat the water to boiling, larger bubbles form and then rise from the bottom, as shown in Figure 1.27B. These bubbles are gaseous water, or water vapor. The result of the physical change can be represented symbolically as
H2O(l)
reaction
reaction
resonance
equilibrium
polarity
H2O(g)
FIGURE 1.27 Water contains small amounts of dissolved nitrogen and oxygen gases. (A) When heated, these molecules go to the gaseous state in bubbles that rise to the surface. (B) When the water begins to boil, it no longer contains dissolved gases, and the bubbles contain gaseous water. ©Brian Moeskau/Moeskau Photography
O2 H2ON2
A B
Air
Air dissolved in water
Water vapor
Pure water
Water molecules surround solute particles in aqueous solutions. For clarity in seeing the particles, this book shows water molecules faded in the background as in Figure 1.27A.
1.3 Energy and Energy Changes 27
Before
After
Before
After
O2
CO2
H2O
CH4
Before
After
O2
CO2
H2O
CH4
Before
After
EXAMPLE 1.11 Physical and Chemical Changes
Do the following molecular-level images represent a chemical change or a physi- cal change?
Solution: The substances after the change have a different composition than the substances before the change. Therefore, this is a chemical change.
Consider This 1.11 How do atom attachments differ for each of the elements in the change shown in the example?
Practice Problem 1.11 Do the following molecular-level images represent a chemical change or a physi- cal change?
Further Practice: Questions 1.95 and 1.96 at the end of the chapter
1.3 Energy and Energy Changes Physical and chemical changes involve energy. Energy is hard to define, but we see and feel evidence of it when something moves or changes temperature. At the con- struction site, Anna and Bill saw a worker pushing a wheelbarrow up a ramp. If released at the top of the ramp, the wheelbarrow would roll back down, converting energy from one form to another in the process. This release of energy is related to the spontaneous process of rolling down the ramp. (A spontaneous process is one that doesn’t have to be forced to occur after it gets started.) But returning the wheelbarrow to the top of the ramp is not spontaneous. It requires a continuous energy input. Similarly, chemical and physical changes are usually accompanied by energy changes. Some chemical reactions are spontaneous. They happen on their own. Others need a continuous energy input from an external source. Con- sider the reaction (Figure 1.28) of hydrogen and oxygen gases to form water vapor, represented by this equation:
2H2(g) + O2(g)
reaction
2H2O(g)
FIGURE 1.28 In the reaction shown here, a balloon was filled with appropriate amounts of hydrogen and oxygen gas. When a lit candle touched the balloon, the hydrogen and oxygen reacted explosively to form water vapor. ©Charles D. Winters
28 Chapter 1 Matter and Energy
This reaction is spontaneous and explosive. It releases a tremendous amount of en- ergy. But the opposite reaction—the breakdown of water into hydrogen and oxygen gases—is not spontaneous. It occurs only if sufficient energy is continuously added, such as by passing electricity through liquid water. This process, called electrolysis, was shown in Figure 1.2, and can be described symbolically as follows: electrolysis
2H2O(l)
reaction
reaction
resonance
equilibrium
conversion
polarity
Δ
2H2(g) + O2(g)
But what is energy? Energy is the capacity to do work or to transfer heat. Work, usually taken to mean mechanical work, occurs when a force acts over a distance. For example, work is done when the construction worker pushes the wheelbarrow up a ramp. Work is done when compressed gases, resulting from the combustion of a fuel, push the piston in the cylinder of an automobile engine. Not all reactions can be made to do work directly, but heat energy can be harnessed to do work. For example, boiling water produces steam, which turns the turbines in power plants. The turbines spin copper coils inside a magnetic field in a generator to produce an electric current.
Energy takes many different forms, and it can be converted from one form to another. Scientists describe two types of energy: kinetic energy and potential en- ergy. Kinetic energy is the energy of motion. The wheelbarrow rolling down a ramp possesses kinetic energy. Potential energy is energy possessed by an object because of its position. The wheelbarrow resting at the top of the ramp has potential energy. If it did not contain this stored energy, it could not release energy when it rolled down the ramp. Any object in a position to be rolled, dropped, or otherwise allowed to move spontaneously has potential energy that will be converted to kinetic energy once the motion starts.
To distinguish kinetic from potential energy, consider the volleyball game Anna and Bill watched while they ate lunch. After they ate, Anna and Bill joined the game. When Anna served the ball, kinetic energy was transferred from her hand to the volleyball (Figure 1.29). As it ascended, the ball transferred some of its kinetic
FIGURE 1.29 Kinetic energy is transferred from the server to the volleyball when it is served. As the volleyball rises in the air, its kinetic energy is converted to potential energy. When it descends, potential energy is converted to kinetic energy.
Kinetic energy of the ball is almost all converted to potential energy.
Kinetic energy is converted to potential energy as the ball ascends.
Kinetic energy is transferred from the server’s hand to the volleyball.
Potential energy is converted to kinetic energy as the ball descends.
Some of the kinetic energy of the ball is transferred to the ground.
1.3 Energy and Energy Changes 29
energy to the surrounding air molecules, but most of its kinetic energy was con- verted to potential energy. Nearly all of the kinetic energy was converted to potential energy when the ball reached the top of its ascent. As the volleyball descended, its potential energy was converted to kinetic energy. Most of its kinetic energy was transferred to the ground when Bill missed the ball. The ball’s remaining kinetic energy allowed it to bounce back up, where again its kinetic energy was converted to potential energy during its ascent.
Other forms of energy—chemical, mechanical, electric, and heat energy, for example—are really just forms of kinetic or potential energy. For example, chemical compounds can release chemical energy, the energy associated with a chemical re- action. Chemical energy is potential energy arising from the positions of the atoms and molecules in the compounds. A compound releases its potential energy when it undergoes a spontaneous chemical reaction that forms substances with less poten- tial energy. For example, the explosive material TNT (trinitrotoluene) contains con- siderable potential energy that is released as kinetic energy when it reacts. Chemical compounds can also have kinetic energy. Molecules move faster as the temperature rises. The motion of molecules or atoms is associated with heat energy, or the ki- netic energy that increases with increasing temperature. The fast-moving gases produced by the explosion of TNT have high kinetic energy.
We will use the lengths of trails behind atoms or molecules to depict their relative speeds. Longer trails correspond to higher speeds.
A B
EXAMPLE 1.12 Molecular Motion and Kinetic Energy
Which of these two samples of argon gas has more kinetic energy?
Solution: The atoms that are moving faster have the greater kinetic energy. Thus, the atoms in A have the greater kinetic energy.
Consider This 1.12 How would the trail lengths have to be changed if the kinetic energy of A de- creased and B increased?
Practice Problem 1.12 Which of the two samples of argon gas is at a lower temperature?
Further Practice: Questions 1.103 and 1.104 at the end of the chapter
Electric energy is associated with the passage of electricity, generally through metals. The electric energy from a battery arises from a chemical reaction. The chemi- cal energy stored in the compounds that make up the battery is converted to electric energy. Electric current passed through the filament of a lightbulb causes the metal to glow red and increases the motion of the atoms. This is a conversion of electric energy to kinetic energy. A lightbulb also gives off light energy. Nuclear energy involves both light and heat. Energy is released when one element is converted to another, as in a nuclear reactor or in the Sun, where hydrogen atoms fuse to form helium.
30 Chapter 1 Matter and Energy
EXAMPLE 1.13 Forms of Energy
Identify examples of potential and kinetic energy in this picture.
Solution: Anything that might move in the picture has potential energy. Kinetic energy is evident in the moving people and vehicles.
Consider This 1.13 What are some examples of electric energy in the picture?
Practice Problem 1.13 Identify three additional forms of energy in the photograph.
Further Practice: Questions 1.107 and 1.108 at the end of the chapter
©Grant V. Faint/Getty Images
Common units of energy are calories, Calories (with a capital C), and joules. The energy input you need per day is about 2000 Calories, 2 million calories, or 8 million joules. We will study ways of measuring energy changes in Chapter 6.
All these forms of energy can be converted into one another. For example, the welder at the construction site starts a gasoline engine that runs a generator that makes the electricity the welder uses to join two pieces of metal together. The chemical energy in the gasoline is converted to mechanical energy that turns the generator. The mechanical energy is converted to electric energy by the generator. The electric energy is converted to heat in an arc that is formed between the welding rod and the metal to be welded. This heat melts the metal and creates the weld. Some of the electric energy is also converted to light. What energy conversions can you observe going on around you right now?
1.4 Scientific Inquiry We began this chapter by describing some of the things that Anna and Bill saw around their campus. To classify the items, they observed similarities and differences in properties. They were making observations to help them understand nature.
Observation is one of the tools of scientific inquiry, but it is not the only one. The scientific method is an approach to asking questions and seeking answers that employs a variety of tools, techniques, and strategies. Although the scientific method is often explained as a series of steps and procedures, it is more accurately described as a way of looking at the world that differs from nonscience forms of inquiry. Scientists, like all humans, use intuition. They generalize about the world, sometimes with insufficient data. Chemists, especially, make inferences about atoms and molecules from data obtained from instruments that aren’t quite capable of showing these tiny particles.