Matter and Energy Chapter 3 PDF

3: MATTER AND ENERGY  CHAPTER OVERVIEW 3: Matter and Energy 3.1: In Your Room 3.2: What is Matter? 3.3: Classifying Matter According to Its State—Solid, Liquid, and Gas 3.4: Classifying Matter According to Its Composition 3.5: Differences in Matter- Physical and Chemical Properties 3.6...

3: MATTER AND ENERGY  CHAPTER OVERVIEW 3: Matter and Energy 3.1: In Your Room 3.2: What is Matter? 3.3: Classifying Matter According to Its State—Solid, Liquid, and Gas 3.4: Classifying Matter According to Its Composition 3.5: Differences in Matter- Physical and Chemical Properties 3.6: Changes in Matter - Physical and Chemical Changes 3.7: Conservation of Mass - There is No New Matter 3.8: Energy 3.9: Energy and Chemical and Physical Change 3.10: Temperature - Random Motion of Molecules and Atoms 3.11: Temperature Changes - Heat Capacity 3.12: Energy and Heat Capacity Calculations 3.E: Matter and Energy (Exercises) 3: Matter and Energy is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 1 3.1: In Your Room Matter is any substance that has mass and takes up space. Matter includes atoms and anything made up of atoms, but not other energy phenomena or waves such as light or sound. While this simple definition is easily applied, the way people view matter is often broken down into two characteristic length scales: the macroscopic and the microscopic. Figure 3.1.1 : A typical American university and college dormitory room in 2002 (CC BY-SA 3.0; Raul654). The macroscopic scale is the length scale on which objects or phenomena are large enough to be visible almost practically with the naked eye, without magnifying optical instruments. Everything that one can see, touch, and handle in the dorm room of Figure 3.1.1 is within the macroscopic scale. To describe each of these objects, only a few macroscopic properties are required. However, each of these items can be decomposed into smaller microscopic scale properties. The microscopic scale is the scale of objects and events smaller than those that can easily be seen by the naked eye, requiring a lens or microscope to see them clearly. All of the everyday objects that we can bump into, touch, or squeeze are ultimately composed of atoms. This ordinary atomic matter is in turn made up of interacting subatomic particles—usually a nucleus of protons and neutrons, and a cloud of orbiting electrons. Because of this, a large number of variables are needed to describe such a system which complicates the characterization.  Matter vs. Mass Matter should not be confused with mass, as the two are not the same in modern physics. Matter is a physical substance of which systems may be composed, while mass is not a substance, but rather a quantitative property of matter and other substances or systems. Contributions & Attributions Wikipedia 3.1: In Your Room is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.1.1 https://chem.libretexts.org/@go/page/47450 3.2: What is Matter?  Learning Objectives Define matter and explain how it is composed of building blocks known as "atoms". We are all familiar with matter. The definition of Matter is anything that has mass and volume (takes up space). For most common objects that we deal with every day, it is fairly simple to demonstrate that they have mass and take up space. You might be able to imagine, however, the difficulty for people several hundred years ago to demonstrate that air had mass and volume. Air (and all other gases) are invisible to the eye, have very small masses compared to equal amounts of solids and liquids, and are quite easy to compress (change volume). Without sensitive equipment, it would have been difficult to convince people that gases are matter. Today, we can measure the mass of a small balloon when it is deflated and then blow it up, tie it off, and measure its mass again to detect the additional mass due to the air inside. The mass of air, under room conditions, that occupies a one quart jar is approximately 0.0002 pounds. This small amount of mass would have been difficult to measure in times before balances were designed to accurately measure very small masses. Later, scientists were able to compress gases into such a small volume that the gases turned into liquids, which made it clear that gases are matter. Figure 3.2.1 : Everything from an ant, to a truck, to the earth, and even the entire galaxy is composed of matter. Images used with permission from Wikipedia (CC_SA-BY-3.0; credit High Contrast). Even though the universe consists of "things" as wildly different as ants and galaxies, the matter that makes up all of these "things" is composed of a very limited number of building blocks. These building blocks are known as atoms, and so far, scientists have discovered or created a grand total of 118 different types of atoms. Scientists have given a name to each different type of atom. A substance that is composed of only one type of atom is called an element. At this point, what should amaze you is that all forms of matter in our universe are made with only 118 different building blocks. In some ways, it's sort of like cooking a gourmet, five- course meal using only three ingredients! How is it possible? To answer that question, you have to understand the ways in which different elements are put together to form matter. The most important method that nature uses to organize atoms into matter is the formation of molecules. Molecules are groups of two or more atoms that have been bonded together. There are millions of different ways to bond atoms together, which means that there are millions of different possible molecules. Each of these molecules has its own set of chemical properties, and it's these 3.2.1 https://chem.libretexts.org/@go/page/47452 properties with which chemists are most concerned. You will learn a lot more about atoms and molecules, including how they were discovered, in a later part of the textbook. Summary All matter has mass and occupies space. All physical objects are made of matter. Matter itself is composed of tiny building blocks known as "atoms". There are only 118 different types of atoms known to man. Frequently, atoms are bonded together to form "molecules". Contributions & Attributions Wikipedia 3.2: What is Matter? is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.2.2 https://chem.libretexts.org/@go/page/47452 3.3: Classifying Matter According to Its State—Solid, Liquid, and Gas  Learning Objectives To describe the solid, liquid and gas phases. Water can take many forms. At low temperatures (below 0 C), it is a solid. When at "normal" temperatures (between 0 C and o o 100 C), it is a liquid. While at temperatures above 100 C, water is a gas (steam). The state that water is in depends upon the o o temperature. Each state has its own unique set of physical properties. Matter typically exists in one of three states: solid, liquid, or gas. Figure 3.3.1 : Matter is usually classified into three classical states, with plasma sometimes added as a fourth state. From left to right: quartz (solid), water (liquid), nitrogen dioxide (gas). The state that a given substance exhibits is also a physical property. Some substances exist as gases at room temperature (oxygen and carbon dioxide), while others, like water and mercury metal, exist as liquids. Most metals exist as solids at room temperature. All substances can exist in any of these three states. Figure 3.3.2 shows the differences among solids, liquids, and gases at the molecular level. A solid has definite volume and shape, a liquid has a definite volume but no definite shape, and a gas has neither a definite volume nor shape. Figure 3.3.2 : A Representation of the Solid, Liquid, and Gas States. (a) Solid O2 has a fixed volume and shape, and the molecules are packed tightly together. (b) Liquid O2 conforms to the shape of its container but has a fixed volume; it contains relatively densely packed molecules. (c) Gaseous O2 fills its container completely—regardless of the container’s size or shape—and consists of widely separated molecules.  Plasma: A Fourth State of Matter Technically speaking, a fourth state of matter called plasma exists, but it does not naturally occur on earth, so we will omit it from our study here. A plasma globe operating in a darkened room. (CC BY-SA 3.0; Chocolateoak). 3.3.1 https://chem.libretexts.org/@go/page/47454 Solids In the solid state, the individual particles of a substance are in fixed positions with respect to each other because there is not enough thermal energy to overcome the intermolecular interactions between the particles. As a result, solids have a definite shape and volume. Most solids are hard, but some (like waxes) are relatively soft. Many solids composed of ions can also be quite brittle. Solids are defined by the following characteristics: Definite shape (rigid) Definite volume Particles vibrate around fixed axes If we were to cool liquid mercury to its freezing point of −39 C, and under the right pressure conditions, we would notice all of o the liquid particles would go into the solid state. Mercury can be solidified when its temperature is brought to its freezing point. However, when returned to room temperature conditions, mercury does not exist in solid state for long, and returns back to its more common liquid form. Solids usually have their constituent particles arranged in a regular, three-dimensional array of alternating positive and negative ions called a crystal. The effect of this regular arrangement of particles is sometimes visible macroscopically, as shown in Figure 3.3.3. Some solids, especially those composed of large molecules, cannot easily organize their particles in such regular crystals and exist as amorphous (literally, “without form”) solids. Glass is one example of an amorphous solid. Figure 3.3.3 : (left) The periodic crystalline lattice structure of quartz SiO in two-dimensions. (right) The random network 2 structure of glassy SiO in two-dimensions. Note that, as in the crystal, each Silicon atom is bonded to 4 oxygen atoms, where the 2 fourth oxygen atom is obscured from view in this plane. Images used with permission (public domain). Liquids If the particles of a substance have enough energy to partially overcome intermolecular interactions, then the particles can move about each other while remaining in contact. This describes the liquid state. In a liquid, the particles are still in close contact, so liquids have a definite volume. However, because the particles can move about each other rather freely, a liquid has no definite shape and takes a shape dictated by its container. Liquids have the following characteristics: No definite shape (takes the shape of its container). Has definite volume. Particles are free to move over each other, but are still attracted to each other. A familiar liquid is mercury metal. Mercury is an anomaly. It is the only metal we know of that is liquid at room temperature. Mercury also has an ability to stick to itself (surface tension)—a property that all liquids exhibit. Mercury has a relatively high surface tension, which makes it very unique. Here you see mercury in its common liquid form. 3.3.2 https://chem.libretexts.org/@go/page/47454  What does the LIQUID METAL Mercury l… l… Video 3.3.1 : Mercury boiling to become a gas. If we heat liquid mercury to its boiling point of 357 C o under the right pressure conditions, we would notice all particles in the liquid state go into the gas state. Gases If the particles of a substance have enough energy to completely overcome intermolecular interactions, then the particles can separate from each other and move about randomly in space. This describes the gas state, which we will consider in more detail elsewhere. Like liquids, gases have no definite shape, but unlike solids and liquids, gases have no definite volume either. The change from solid to liquid usually does not significantly change the volume of a substance. However, the change from a liquid to a gas significantly increases the volume of a substance, by a factor of 1,000 or more. Gases have the following characteristics: No definite shape (takes the shape of its container) No definite volume Particles move in random motion with little or no attraction to each other Highly compressible Table 3.3.1 : Characteristics of the Three States of Matter Characteristics Solids Liquids Gases shape definite indefinite indefinite volume definite definite indefinite relative intermolecular interaction strong moderate weak strength relative particle positions in contact and fixed in place in contact but not fixed not in contact, random positions  Example 3.3.1 What state or states of matter does each statement, describe? a. This state has a definite volume, but no definite shape. b. This state has no definite volume. c. This state allows the individual particles to move about while remaining in contact. Solution a. This statement describes the liquid state. b. This statement describes the gas state. c. This statement describes the liquid state. 3.3.3 https://chem.libretexts.org/@go/page/47454   Exercise 3.3.1 What state or states of matter does each statement describe? a. This state has individual particles in a fixed position with regard to each other. b. This state has individual particles far apart from each other in space. c. This state has a definite shape. Answer a: solid Answer b: gas Answer c: solid Summary Three states of matter exist—solid, liquid, and gas. Solids have a definite shape and volume. Liquids have a definite volume, but take the shape of the container. Gases have no definite shape or volume. 3.3: Classifying Matter According to Its State—Solid, Liquid, and Gas is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.3.4 https://chem.libretexts.org/@go/page/47454 3.4: Classifying Matter According to Its Composition  Learning Objectives Explain the difference between a pure substance and a mixture. Explain the difference between an element and a compound. Explain the difference between a homogeneous mixture and a heterogeneous mixture. One useful way of organizing our understanding of matter is to think of a hierarchy that extends down from the most general and complex to the simplest and most fundamental (Figure 3.4.1). Matter can be classified into two broad categories: pure substances and mixtures. A pure substance is a form of matter that has a constant composition (meaning that it is the same everywhere) and properties that are constant throughout the sample (meaning that there is only one set of properties such as melting point, color, boiling point, etc. throughout the matter). A material composed of two or more substances is a mixture. Elements and compounds are both examples of pure substances. A substance that cannot be broken down into chemically simpler components is an element. Aluminum, which is used in soda cans, is an element. A substance that can be broken down into chemically simpler components (because it has more than one element) is a compound. For example, water is a compound composed of the elements hydrogen and oxygen. Today, there are about 118 elements in the known universe. In contrast, scientists have identified tens of millions of different compounds to date. Figure 3.4.1 : Relationships between the Types of Matter and the Methods Used to Separate Mixtures Ordinary table salt is called sodium chloride. It is considered a substance because it has a uniform and definite composition. All samples of sodium chloride are chemically identical. Water is also a pure substance. Salt easily dissolves in water, but salt water cannot be classified as a substance because its composition can vary. You may dissolve a small amount of salt or a large amount into a given amount of water. A mixture is a physical blend of two or more components, each of which retains its own identity and properties in the mixture. Only the form of the salt is changed when it is dissolved into water. It retains its composition and properties. 3.4.1 https://chem.libretexts.org/@go/page/47456 A homogeneous mixture is a mixture in which the composition is uniform throughout the mixture. The salt water described above is homogeneous because the dissolved salt is evenly distributed throughout the entire salt water sample. Often it is easy to confuse a homogeneous mixture with a pure substance because they are both uniform. The difference is that the composition of the substance is always the same. The amount of salt in the salt water can vary from one sample to another. All solutions are considered homogeneous because the dissolved material is present in the same amount throughout the solution. A heterogeneous mixture is a mixture in which the composition is not uniform throughout the mixture. Vegetable soup is a heterogeneous mixture. Any given spoonful of soup will contain varying amounts of the different vegetables and other components of the soup.  Phase A phase is any part of a sample that has a uniform composition and properties. By definition, a pure substance or a homogeneous mixture consists of a single phase. A heterogeneous mixture consists of two or more phases. When oil and water are combined, they do not mix evenly, but instead form two separate layers. Each of the layers is called a phase.  Example 3.4.1 Identify each substance as a compound, an element, a heterogeneous mixture, or a homogeneous mixture (solution). a. filtered tea b. freshly squeezed orange juice c. a compact disc d. aluminum oxide, a white powder that contains a 2:3 ratio of aluminum and oxygen atoms e. selenium Given: a chemical substance Asked for: its classification Strategy: A. Decide whether a substance is chemically pure. If it is pure, the substance is either an element or a compound. If a substance can be separated into its elements, it is a compound. B. If a substance is not chemically pure, it is either a heterogeneous mixture or a homogeneous mixture. If its composition is uniform throughout, it is a homogeneous mixture. Solution a. A) Tea is a solution of compounds in water, so it is not chemically pure. It is usually separated from tea leaves by filtration. B) Because the composition of the solution is uniform throughout, it is a homogeneous mixture. b. A) Orange juice contains particles of solid (pulp) as well as liquid; it is not chemically pure. B) Because its composition is not uniform throughout, orange juice is a heterogeneous mixture. c. A) A compact disc is a solid material that contains more than one element, with regions of different compositions visible along its edge. Hence, a compact disc is not chemically pure. B) The regions of different composition indicate that a compact disc is a heterogeneous mixture. d. A) Aluminum oxide is a single, chemically pure compound. e. A) Selenium is one of the known elements.  Exercise 3.4.1 Identify each substance as a compound, an element, a heterogeneous mixture, or a homogeneous mixture (solution). a. white wine b. mercury c. ranch-style salad dressing d. table sugar (sucrose) 3.4.2 https://chem.libretexts.org/@go/page/47456  Answer a: homogeneous mixture (solution) Answer b: element Answer c: heterogeneous mixture Answer d: compound  Example 3.4.2 How would a chemist categorize each example of matter? a. saltwater b. soil c. water d. oxygen Solution a. Saltwater acts as if it were a single substance even though it contains two substances—salt and water. Saltwater is a homogeneous mixture, or a solution. b. Soil is composed of small pieces of a variety of materials, so it is a heterogeneous mixture. c. Water is a substance. More specifically, because water is composed of hydrogen and oxygen, it is a compound. d. Oxygen, a substance, is an element.  Exercise 3.4.2 How would a chemist categorize each example of matter? a. coffee b. hydrogen c. an egg Answer a: a homogeneous mixture (solution), assuming it is filtered coffee Answer b: element Answer c: heterogeneous mixture Summary Matter can be classified into two broad categories: pure substances and mixtures. A pure substance is a form of matter that has a constant composition and properties that are constant throughout the sample. Mixtures are physical combinations of two or more elements and/or compounds. Mixtures can be classified as homogeneous or heterogeneous. Elements and compounds are both examples of pure substances. Compounds are substances that are made up of more than one type of atom. Elements are the simplest substances made up of only one type of atom. Vocabulary Element: a substance that is made up of only one type of atom. Compound:a substance that is made up of more than one type of atom bonded together. Mixture: a combination of two or more elements or compounds which have not reacted to bond together; each part in the mixture retains its own properties. 3.4.3 https://chem.libretexts.org/@go/page/47456 Contributions & Attributions Stephen Lower, Professor Emeritus (Simon Fraser U.) Chem1 Virtual Textbook 3.4: Classifying Matter According to Its Composition is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew, Henry Agnew, Stephen Lower, & Stephen Lower. 3.4.4 https://chem.libretexts.org/@go/page/47456 3.5: Differences in Matter- Physical and Chemical Properties  Learning Objectives To separate physical from chemical properties. All matter has physical and chemical properties. Physical properties are characteristics that scientists can measure without changing the composition of the sample under study, such as mass, color, and volume (the amount of space occupied by a sample). Chemical properties describe the characteristic ability of a substance to react to form new substances; they include its flammability and susceptibility to corrosion. All samples of a pure substance have the same chemical and physical properties. For example, pure copper is always a reddish-brown solid (a physical property) and always dissolves in dilute nitric acid to produce a blue solution and a brown gas (a chemical property). Physical Property A physical property is a characteristic of a substance that can be observed or measured without changing the identity of the substance. Silver is a shiny metal that conducts electricity very well. It can be molded into thin sheets, a property called malleability. Salt is dull and brittle and conducts electricity when it has been dissolved into water, which it does quite easily. Physical properties of matter include color, hardness, malleability, solubility, electrical conductivity, density, melting point, and boiling point. For the elements, color does not vary much from one element to the next. The vast majority of elements are colorless, silver, or gray. Some elements do have distinctive colors: sulfur and chlorine are yellow, copper is (of course) copper-colored, and elemental bromine is red. However, density can be a very useful parameter for identifying an element. Of the materials that exist as solids at room temperature, iodine has a very low density compared to zinc, chromium, and tin. Gold has a very high density, as does platinum. Pure water, for example, has a density of 0.998 g/cm3 at 25°C. The average densities of some common substances are in Table 3.5.1. Notice that corn oil has a lower mass to volume ratio than water. This means that when added to water, corn oil will “float.” Table 3.5.1 : Densities of Common Substances Substance Density at 25°C (g/cm3) blood 1.035 body fat 0.918 whole milk 1.030 corn oil 0.922 mayonnaise 0.910 honey 1.420 Hardness helps determine how an element (especially a metal) might be used. Many elements are fairly soft (silver and gold, for example) while others (such as titanium, tungsten, and chromium) are much harder. Carbon is an interesting example of hardness. In graphite, (the "lead" found in pencils) the carbon is very soft, while the carbon in a diamond is roughly seven times as hard. 3.5.1 https://chem.libretexts.org/@go/page/47458  Figure 3.5.1 : Pencil (left) and Diamond ring (right). Both are a form of carbon, but exhibit very different physical properties. Melting and boiling points are somewhat unique identifiers, especially of compounds. In addition to giving some idea as to the identity of the compound, important information can be obtained about the purity of the material. Chemical Properties Chemical properties of matter describe its potential to undergo some chemical change or reaction by virtue of its composition. The elements, electrons, and bonds that are present give the matter potential for chemical change. It is quite difficult to define a chemical property without using the word "change". Eventually, after studying chemistry for some time, you should be able to look at the formula of a compound and state some chemical property. For example, hydrogen has the potential to ignite and explode given the right conditions—this is a chemical property. Metals in general have the chemical property of reacting with an acid. Zinc reacts with hydrochloric acid to produce hydrogen gas—this is a chemical property. Figure 3.5.2 : Heavy rust on the links of a chain near the Golden Gate Bridge in San Francisco; it was continuously exposed to moisture and salt spray, causing surface breakdown, cracking, and flaking of the metal. (CC BY-SA 3.0; Marlith). A chemical property of iron is its capability of combining with oxygen to form iron oxide, the chemical name of rust (Figure 3.5.2). The more general term for rusting and other similar processes is corrosion. Other terms that are commonly used in descriptions of chemical changes are burn, rot, explode, decompose, and ferment. Chemical properties are very useful in identifying substances. However, unlike physical properties, chemical properties can only be observed as the substance is in the process of being changed into a different substance. Table 3.5.2 : Contrasting Physical and Chemical Properties Physical Properties Chemical Properties Gallium metal melts at 30 oC. Iron metal rusts. Mercury is a very dense liquid. A green banana turns yellow when it ripens. Gold is shiny. A dry piece of paper burns. 3.5.2 https://chem.libretexts.org/@go/page/47458   Example 3.5.1 Which of the following is a chemical property of iron? a. Iron corrodes in moist air. b. Density = 7.874 g/cm3 c. Iron is soft when pure. d. Iron melts at 1808 K. Solution "Iron corrodes in moist air" is the only chemical property of iron from the list.  Exercise 3.5.1A Which of the following is a physical property of matter? a. corrosiveness b. pH (acidity) c. density d. flammability Answer c  Exercise 3.5.1B Which of the following is a chemical property? a. flammability b. melting point c. boiling point d. density Answer a Summary A physical property is a characteristic of a substance that can be observed or measured without changing the identity of the substance. Physical properties include color, density, hardness, and melting and boiling points. A chemical property describes the ability of a substance to undergo a specific chemical change. To identify a chemical property, we look for a chemical change. A chemical change always produces one or more types of matter that differ from the matter present before the change. The formation of rust is a chemical change because rust is a different kind of matter than the iron, oxygen, and water present before the rust formed. 3.5: Differences in Matter- Physical and Chemical Properties is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.5.3 https://chem.libretexts.org/@go/page/47458 3.6: Changes in Matter - Physical and Chemical Changes  Learning Objectives Label a change as chemical or physical. List evidence that can indicate a chemical change occurred. Change is happening all around us all of the time. Just as chemists have classified elements and compounds, they have also classified types of changes. Changes are classified as either physical or chemical changes. Chemists learn a lot about the nature of matter by studying the changes that matter can undergo. Chemists make a distinction between two different types of changes that they study—physical changes and chemical changes. Physical Change Physical changes are changes in which no bonds are broken or formed. This means that the same types of compounds or elements that were there at the beginning of the change are there at the end of the change. Because the ending materials are the same as the beginning materials, the properties (such as color, boiling point, etc.) will also be the same. Physical changes involve moving molecules around, but not changing them. Some types of physical changes include: Changes of state (changes from a solid to a liquid or a gas and vice versa). Separation of a mixture. Physical deformation (cutting, denting, stretching). Making solutions (special kinds of mixtures). As an ice cube melts, its shape changes as it acquires the ability to flow. However, its composition does not change. Melting is an example of a physical change. A physical change is a change to a sample of matter in which some properties of the material change, but the identity of the matter does not. When liquid water is heated, it changes to water vapor. However, even though the physical properties have changed, the molecules are exactly the same as before. We still have each water molecule containing two hydrogen atoms and one oxygen atom covalently bonded. When you have a jar containing a mixture of pennies and nickels and you sort the mixture so that you have one pile of pennies and another pile of nickels, you have not altered the identity of the pennies or the nickels—you've merely separated them into two groups. This would be an example of a physical change. Similarly, if you have a piece of paper, you don't change it into something other than a piece of paper by ripping it up. What was paper before you started tearing is still paper when you are done. Again, this is an example of a physical change. Figure 3.6.1 : Ice melting is a physical change. When liquid water (H O ) freezes into a solid state (ice), it appears changed; 2 however, this change is only physical, as the composition of the constituent molecules is the same: 11.19% hydrogen and 88.81% oxygen by mass. (Public Domain; Moussa). Physical changes can further be classified as reversible or irreversible. The melted ice cube may be refrozen, so melting is a reversible physical change. Physical changes that involve a change of state are all reversible. Other changes of state include vaporization (liquid to gas), freezing (liquid to solid), and condensation (gas to liquid). Dissolving is also a reversible physical change. When salt is dissolved into water, the salt is said to have entered the aqueous state. The salt may be regained by boiling off the water, leaving the salt behind. 3.6.1 https://chem.libretexts.org/@go/page/47460 Chemical Change Chemical changes occur when bonds are broken and/or formed between molecules or atoms. This means that one substance with a certain set of properties (such as melting point, color, taste, etc) is turned into a different substance with different properties. Chemical changes are frequently harder to reverse than physical changes. One good example of a chemical change is burning a candle. The act of burning paper actually results in the formation of new chemicals (carbon dioxide and water) from the burning of the wax. Another example of a chemical change is what occurs when natural gas is burned in your furnace. This time, on the left there is a molecule of methane, CH , and two molecules of oxygen, 4 O ; on the right are two molecules of water, H O , and one molecule of carbon dioxide, CO. In this case, not only has the 2 2 2 appearance changed, but the structure of the molecules has also changed. The new substances do not have the same chemical properties as the original ones. Therefore, this is a chemical change. Figure 3.6.2 : Burning of wax to generate water and carbon dioxide is a chemical reaction. (CC-SA-BY-3.0; Andrikkos ) We can't actually see molecules breaking and forming bonds, although that's what defines chemical changes. We have to make other observations to indicate that a chemical change has happened. Some of the evidence for chemical change will involve the energy changes that occur in chemical changes, but some evidence involves the fact that new substances with different properties are formed in a chemical change. Observations that help to indicate chemical change include: Temperature changes (either the temperature increases or decreases). Light given off. Unexpected color changes (a substance with a different color is made, rather than just mixing the original colors together). Bubbles are formed (but the substance is not boiling—you made a substance that is a gas at the temperature of the beginning materials, instead of a liquid). Different smell or taste (do not taste your chemistry experiments, though!). A solid forms if two clear liquids are mixed (look for floaties—technically called a precipitate).  Example 3.6.1 Label each of the following changes as a physical or chemical change. Give evidence to support your answer. a. Boiling water. b. A nail rusting. c. A green solution and colorless solution are mixed. The resulting mixture is a solution with a pale green color. d. Two colorless solutions are mixed. The resulting mixture has a yellow precipitate. Solution a. Physical: boiling and melting are physical changes. When water boils, no bonds are broken or formed. The change could be written: H O (l) → H O (g) 2 2 b. Chemical: The dark grey nail changes color to form an orange flaky substance (the rust); this must be a chemical change. Color changes indicate chemical change. The following reaction occurs: Fe + O → Fe O 2 2 3 3.6.2 https://chem.libretexts.org/@go/page/47460  c. Physical: because none of the properties changed, this is a physical change. The green mixture is still green and the colorless solution is still colorless. They have just been spread together. No color change occurred or other evidence of chemical change. d. Chemical: the formation of a precipitate and the color change from colorless to yellow indicate a chemical change.  Exercise 3.6.1 Label each of the following changes as a physical or chemical change. a. A mirror is broken. b. An iron nail corroded in moist air c. Copper metal is melted. d. A catalytic converter changes nitrogen dioxide to nitrogen gas and oxygen gas. Answer a: physical change Answer b: chemical change Answer c: physical change Answer d: chemical change Separating Mixtures Through Physical Changes Homogeneous mixtures (solutions) can be separated into their component substances by physical processes that rely on differences in some physical property, such as differences in their boiling points. Two of these separation methods are distillation and crystallization. Distillation makes use of differences in volatility, a measure of how easily a substance is converted to a gas at a given temperature. A simple distillation apparatus for separating a mixture of substances, at least one of which is a liquid. The most volatile component boils first and is condensed back to a liquid in the water-cooled condenser, from which it flows into the receiving flask. If a solution of salt and water is distilled, for example, the more volatile component, pure water, collects in the receiving flask, while the salt remains in the distillation flask. 3.6.3 https://chem.libretexts.org/@go/page/47460  Figure 3.6.3 : The Distillation of a Solution of Table Salt in Water. The solution of salt in water is heated in the distilling flask until it boils. The resulting vapor is enriched in the more volatile component (water), which condenses to a liquid in the cold condenser and is then collected in the receiving flask. Parts of a distillation setup: Bunsen burner, salt water in distilling flask, condenser with cool water in and warm water out, pure water in receiving flask Mixtures of two or more liquids with different boiling points can be separated with a more complex distillation apparatus. One example is the refining of crude petroleum into a range of useful products: aviation fuel, gasoline, kerosene, diesel fuel, and lubricating oil (in the approximate order of decreasing volatility). Another example is the distillation of alcoholic spirits such as brandy or whiskey. This relatively simple procedure caused more than a few headaches for federal authorities in the 1920s during the era of Prohibition, when illegal stills proliferated in remote regions of the United States. Another example for using physical properties to separate mixtures is filtration (Figure 3.6.4). Filtration is any mechanical, physical or biological operation that separates solids from fluids (liquids or gases) by adding a medium through which only the fluid can pass. The fluid that passes through is called the filtrate. There are many different methods of filtration; all aim to attain the separation of substances. Separation is achieved by some form of interaction between the substance or objects to be removed and the filter. The substance that is to pass through the filter must be a fluid, i.e. a liquid or gas. Methods of filtration vary depending on the location of the targeted material, i.e. whether it is dissolved in the fluid phase or suspended as a solid. 3.6.4 https://chem.libretexts.org/@go/page/47460  Figure 3.6.4 : Filtration for the separation of solids from a hot solution. (CC BY-SA 4.0; Suman6395). Summary Chemists make a distinction between two different types of changes that they study—physical changes and chemical changes. Physical changes are changes that do not alter the identity of a substance. Chemical changes are changes that occur when one substance is turned into another substance. Chemical changes are frequently harder to reverse than physical changes. Observations that indicate a chemical change has occurred include color change, temperature change, light given off, formation of bubbles, formation of a precipitate, etc. Contributions & Attributions Boundless (www.boundless.com) 3.6: Changes in Matter - Physical and Chemical Changes is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.6.5 https://chem.libretexts.org/@go/page/47460 3.7: Conservation of Mass - There is No New Matter It may seem as though burning destroys matter, but the same amount, or mass, of matter still exists after a campfire as before. Look at Figure 3.7.1 below. It shows that when wood burns, it combines with oxygen and changes not only to ashes, but also to carbon dioxide and water vapor. The gases float off into the air, leaving behind just the ashes. Suppose you had measured the mass of the wood before it burned and the mass of the ashes after it burned. Also suppose you had been able to measure the oxygen used by the fire and the gases produced by the fire. What would you find? The total mass of matter after the fire would be the same as the total mass of matter before the fire. Figure 3.7.1 : Burning is a chemical process. The flames are caused as a result of a fuel undergoing combustion (burning). (CC BY- SA 2.5; Einar Helland Berger for fire and Walter Siegmund for ash). Law of Conservation of Mass The law of conservation of mass was created in 1789 by a French chemist, Antoine Lavoisier. The law of conservation of mass states that matter cannot be created or destroyed in a chemical reaction. For example, when wood burns, the mass of the soot, ashes, and gases equals the original mass of the charcoal and the oxygen when it first reacted. So the mass of the product equals the mass of the reactant. A reactant is the chemical reaction of two or more elements to make a new substance, and a product is the substance that is formed as the result of a chemical reaction (Video 3.7.1). Matter and its corresponding mass may not be able to be created or destroyed, but can change forms to other substances like liquids, gases, and solids. Demo - Conservation of Matter Video 3.7.1 : This is a nice little demonstration showing the Conservation of Mass in action. If you witness a 300 kg tree burn to the ground, there are only ashes left after the burn, and all of them together weigh 10 kg. It may make you wonder where the other 290 kg went. The missing 290 kg was released into the atmosphere as smoke, so the only thing left that you can see is the 10 kg of ash. If you know the law of conservation of mass, then you know that the other 290 kg has to go somewhere, because it has to equal the mass of the tree before it burnt down.  Example 3.7.1 If heating 10.0 grams of calcium carbonate (CaCO3) produces 4.4 g of carbon dioxide (CO2) and 5.6 g of calcium oxide (CaO), show that these observations are in agreement with the law of conservation of mass. 3.7.1 https://chem.libretexts.org/@go/page/47461  Solution Mass of the reactants = Mass of the products 10.0 g of CaCO = 4.4 g of CO + 5.6 g of CaO 3 2 10.0 g of reactant = 10.0 g of products Because the mass of the reactant is equal to the mass of the products, the observations are in agreement with the law of conservation of mass.  Exercise 3.7.1 Potassium hydroxide (KOH ) readily reacts with carbon dioxide (CO ) to produce potassium carbonate (K CO ) and water ( 2 2 3 H O ). How many grams of potassium carbonate are produced if 224.4 g of KOH reacts with 88.0 g of CO ? The reaction 2 2 also produces 36.0 g of water. Answer 276.4 g of potassium carbonate The Law is also applicable to both chemical and physical changes. For example, if you have an ice cube that melts into a liquid and you heat that liquid up, it becomes a gas. It will appear to have disappeared, but is still there. Summary Burning and other changes in matter do not destroy matter. The mass of matter is always the same before and after the changes occur. The law of conservation of mass states that matter cannot be created or destroyed. Contributions & Attributions Binod Shrestha (University of Lorraine) 3.7: Conservation of Mass - There is No New Matter is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.7.2 https://chem.libretexts.org/@go/page/47461 3.8: Energy  Learning Objectives Define heat and work. Distinguish between kinetic energy and potential energy. State the law of conservation of matter and energy. Just like matter, energy is a term that we are all familiar with and use on a daily basis. Before you go on a long hike, you eat an energy bar; every month, the energy bill is paid; on TV, politicians argue about the energy crisis. But what is energy? If you stop to think about it, energy is very complicated. When you plug a lamp into an electric socket, you see energy in the form of light, but when you plug a heating pad into that same socket, you only feel warmth. Without energy, we couldn't turn on lights, we couldn't brush our teeth, we couldn't make our lunch, and we couldn't travel to school. In fact, without energy, we couldn't even wake up because our bodies require energy to function. We use energy for every single thing that we do, whether we are awake or asleep. Ability to Do Work or Produce Heat When we speak of using energy, we are really referring to transferring energy from one place to another. When you use energy to throw a ball, you transfer energy from your body to the ball, and this causes the ball to fly through the air. When you use energy to warm your house, you transfer energy from the furnace to the air in your home, and this causes the temperature in your house to rise. Although energy is used in many kinds of different situations, all of these uses rely on energy being transferred in one of two ways. Energy can be transferred as heat or as work. When scientists speak of heat, they are referring to energy that is transferred from an object with a higher temperature to an object with a lower temperature, as a result of the temperature difference. Heat will "flow" from the hot object to the cold object until both end up at the same temperature. When you cook with a metal pot, you witness energy being transferred in the form of heat. Initially, only the stove element is hot—the pot and the food inside the pot are cold. As a result, heat moves from the hot stove element to the cold pot. After a while, enough heat is transferred from the stove to the pot, raising the temperature of the pot and all of its contents (Figure 3.8.1). Figure 3.8.1 : Energy is transferred as heat from the hot stove element to the cooler pot until the pot and its contents become just as hot as the element. The energy that is transferred into the pot as heat is then used to cook the food. Heat is only one way in which energy can be transferred. Energy can also be transferred as work. The scientific definition of work is force (any push or pull) applied over a distance. When you push an object and cause it to move, you do work, and you transfer some of your energy to the object. At this point, it's important to warn you of a common misconception. Sometimes we think that the amount of work done can be measured by the amount of effort put in. This may be true in everyday life, but it is not true in science. By definition, scientific work requires that force be applied over a distance. It does not matter how hard you push or how hard you pull. If you have not moved the object, you haven't done any work. So far, we've talked about the two ways in which energy can be transferred from one place, or object, to another. Energy can be transferred as heat, and energy can be transferred as work. But the question still remains—what IS energy? Kinetic Energy Machines use energy, our bodies use energy, energy comes from the sun, energy comes from volcanoes, energy causes forest fires, and energy helps us to grow food. With all of these seemingly different types of energy, it's hard to believe that there are really only 3.8.1 https://chem.libretexts.org/@go/page/47462 two different forms of energy: kinetic energy and potential energy. Kinetic energy is energy associated with motion. When an object is moving, it has kinetic energy. When the object stops moving, it has no kinetic energy. While all moving objects have kinetic energy, not all moving objects have the same amount of kinetic energy. The amount of kinetic energy possessed by an object is determined by its mass and its speed. The heavier an object is and the faster it is moving, the more kinetic energy it has. Kinetic energy is very common, and it's easy to spot examples of it in the world around you. Sometimes we even try to capture kinetic energy and use it to power things like our home appliances. If you are from California, you might have driven through the Tehachapi Pass near Mojave or the Montezuma Hills in Solano County and seen the windmills lining the slopes of the mountains (Figure 3.8.2). These are two of the larger wind farms in North America. As wind rushes along the hills, the kinetic energy of the moving air particles turns the windmills, trapping the wind's kinetic energy so that people can use it in their houses and offices. Figure 3.8.2 : A wind farm in Solano County harnesses the kinetic energy of the wind. (CC BY-SA 3.0 Unported; BDS2006 at Wikipedia) Potential Energy Potential energy is stored energy. It is energy that remains available until we choose to use it. Think of a battery in a flashlight. If left on, the flashlight battery will run out of energy within a couple of hours, and the flashlight will die. If, however, you only use the flashlight when you need it, and turn it off when you don’t, the battery will last for days or even months. The battery contains a certain amount of energy, and it will power the flashlight for a certain amount of time, but because the battery stores potential energy, you can choose to use the energy all at once, or you can save it and only use a small amount at a time. Any stored energy is potential energy. There are a lot of different ways in which energy can be stored, and this can make potential energy very difficult to recognize. In general, an object has potential energy because of its position relative to another object. For example, when a rock is held above the earth, it has potential energy because of its position relative to the ground. This is potential energy because the energy is stored for as long as the rock is held in the air. Once the rock is dropped, though, the stored energy is released as kinetic energy as the rock falls. Chemical Energy There are other common examples of potential energy. A ball at the top of a hill stores potential energy until it is allowed to roll to the bottom. When two magnets are held next to one another, they store potential energy too. For some examples of potential energy, though, it's harder to see how "position" is involved. In chemistry, we are often interested in what is called chemical potential energy. Chemical potential energy is energy stored in the atoms, molecules, and chemical bonds that make up matter. How does this depend on position? As you learned earlier, the world, and all of the chemicals in it are made up of atoms and molecules. These store potential energy that is dependent on their positions relative to one another. Of course, you can't see atoms and molecules. Nevertheless, scientists do know a lot about the ways in which atoms and molecules interact, and this allows them to figure out how much potential energy is stored in a specific quantity (like a cup or a gallon) of a particular chemical. Different chemicals have different amounts of potential energy because they are made up of different atoms, and those atoms have different positions relative to one another. Since different chemicals have different amounts of potential energy, scientists will sometimes say that potential energy depends not only on position, but also on composition. Composition affects potential energy because it determines which 3.8.2 https://chem.libretexts.org/@go/page/47462  molecules and atoms end up next to one another. For example, the total potential energy in a cup of pure water is different than the total potential energy in a cup of apple juice, because the cup of water and the cup of apple juice are composed of different amounts of different chemicals. At this point, you may wonder just how useful chemical potential energy is. If you want to release the potential energy stored in an object held above the ground, you just drop it. But how do you get potential energy out of chemicals? It's actually not difficult. Use the fact that different chemicals have different amounts of potential energy. If you start with chemicals that have a lot of potential energy and allow them to react and form chemicals with less potential energy, all the extra energy that was in the chemicals at the beginning, but not at the end, is released. Units of Energy Energy is measured in one of two common units: the calorie and the joule. The joule (J) is the SI unit of energy. The calorie is familiar because it is commonly used when referring to the amount of energy contained within food. A calorie (cal) is the quantity of heat required to raise the temperature of 1 gram of water by 1 C. For example, raising the temperature of 100 g of water from o 20 C to 22 C would require 100 × 2 = 200 cal. o o Calories contained within food are actually kilocalories (kcal). In other words, if a certain snack contains 85 food calories, it actually contains 85 kcal or 85, 000 cal. In order to make the distinction, the dietary calorie is written with a capital C. 1 kilocalorie = 1 Calorie = 1000 calories To say that the snack "contains" 85 Calories means that 85 kcal of energy are released when that snack is processed by your body. Heat changes in chemical reactions are typically measured in joules rather than calories. The conversion between a joule and a calorie is shown below. 1 J = 0.2390 cal or 1 cal = 4.184 J We can calculate the amount of heat released in kilojoules when a 400 Calorie hamburger is digested. 4.184 kJ 3 400 Cal = 400 kcal × = 1.67 × 10 kJ 1 kcal Summary Any time we use energy, we transfer energy from one object to another. Energy can be transferred in one of two ways: as heat, or as work. Heat is the term given to energy that is transferred from a hot object to a cooler object due to the difference in their temperatures. Work is the term given to energy that is transferred as a result of a force applied over a distance. Energy comes in two fundamentally different forms: kinetic energy and potential energy. Kinetic energy is the energy of motion. Potential energy is stored energy that depends on the position of an object relative to another object. Chemical potential energy is a special type of potential energy that depends on the positions of different atoms and molecules relative to one another. Chemical potential energy can also be thought of according to its dependence on chemical composition. Energy can be converted from one form to another. The total amount of mass and energy in the universe is conserved. Contributions & Attributions Wikibooks 3.8: Energy is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.8.3 https://chem.libretexts.org/@go/page/47462 3.9: Energy and Chemical and Physical Change  Learning Objectives Define endothermic and exothermic reactions. Describe how heat is transferred in endothermic and exothermic reactions. Determine whether a reaction is endothermic or exothermic through observations, temperature changes, or an energy diagram. So far, we have talked about how energy exists as either kinetic energy or potential energy and how energy can be transferred as either heat or work. While it's important to understand the difference between kinetic energy and potential energy and the difference between heat and work, the truth is, energy is constantly changing. Kinetic energy is constantly being turned into potential energy, and potential energy is constantly being turned into kinetic energy. Likewise, energy that is transferred as work might later end up transferred as heat, while energy that is transferred as heat might later end up being used to do work. Even though energy can change form, it must still follow one fundamental law: Energy cannot be created or destroyed, it can only be changed from one form to another. This law is known as the Law of Conservation of Energy. In a lot of ways, energy is like money. You can exchange quarters for dollar bills and dollar bills for quarters, but no matter how often you convert between the two, you will not end up with any more or any less money than you started with. Similarly, you can transfer (or spend) money using cash, or transfer money using a credit card; but you still spend the same amount of money, and the store still makes the same amount of money. A campfire is an example of basic thermochemistry. The reaction is initiated by the application of heat from a match. The reaction converting wood to carbon dioxide and water (among other things) continues, releasing heat energy in the process. This heat energy can then be used to cook food, roast marshmallows, or just keep warm when it's cold outside. An image of a campfire with colored flames, made by the burning of a garden hose in a copper pipe. (CC SA-BY 3.0; Jared) Exothermic and Endothermic Processes When physical or chemical changes occur, they are generally accompanied by a transfer of energy. The law of conservation of energy states that in any physical or chemical process, energy is neither created nor destroyed. In other words, the entire energy in the universe is conserved. In order to better understand the energy changes taking place during a reaction, we need to define two parts of the universe: the system and the surroundings. The system is the specific portion of matter in a given space that is being studied during an experiment or an observation. The surroundings are everything in the universe that is not part of the system. In practical terms for a laboratory chemist, the system is the particular chemicals being reacted, while the surroundings are the immediate vicinity within the room. During most processes, energy is exchanged between the system and the surroundings. If the system loses a certain amount of energy, that same amount of energy is gained by the surroundings. If the system gains a certain amount of energy, that energy is supplied by the surroundings. A chemical reaction or physical change is endothermic if heat is absorbed by the system from the surroundings. In the course of an endothermic process, the system gains heat from the surroundings and so the temperature of the surroundings decreases. The quantity of heat for a process is represented by the letter q. The sign of q for an endothermic process is positive because the system 3.9.1 https://chem.libretexts.org/@go/page/47463 is gaining heat. A chemical reaction or physical change is exothermic if heat is released by the system into the surroundings. Because the surroundings are gaining heat from the system, the temperature of the surroundings increases. The sign of q for an exothermic process is negative because the system is losing heat. Figure 3.9.1 : (A) Endothermic reaction. (B) Exothermic reaction. Endothermic reaction: surroundings get cooler and delta H is greater than 0, Exothermic reaction: surroundings get warmer and delta H is less than 0 During phase changes, energy changes are usually involved. For example, when solid dry ice vaporizes (physical change), carbon dioxide molecules absorb energy. When liquid water becomes ice, energy is released. Remember that all chemical reactions involve a change in the bonds of the reactants. The bonds in the reactants are broken and the bonds of the products are formed. Chemical bonds have potential energy or "stored energy". Because we are changing the bonding, this means we are also changing how much of this "stored energy" there is in a reaction. Energy changes are frequently shown by drawing an energy diagram. Energy diagrams show the stored/hidden energy of the reactants and products as well as the activation energy. If, on an energy diagram, the products have more stored energy than the reactants started with, the reaction is endothermic. You had to give the reaction energy. If, on the energy diagram, the products have less stored energy than the reactants started with, the reaction is exothermic.  Example 3.9.1 Label each of the following processes as endothermic or exothermic. a. water boiling b. gasoline burning c. ice forming on a pond Solution a. Endothermic—you must put a pan of water on the stove and give it heat in order to get water to boil. Because you are adding heat/energy, the reaction is endothermic. b. Exothermic—when you burn something, it feels hot to you because it is giving off heat into the surroundings. c. Exothermic—think of ice forming in your freezer instead. You put water into the freezer, which takes heat out of the water, to get it to freeze. Because heat is being pulled out of the water, it is exothermic. Heat is leaving.  Exercise 3.9.1 Label each of the following processes as endothermic or exothermic. a. water vapor condensing b. gold melting Answer (a) exothermic Answer (b) endothermic 3.9.2 https://chem.libretexts.org/@go/page/47463 Summary Phase changes involve changes in energy. All chemical reactions involve changes in energy. This may be a change in heat, electricity, light, or other forms of energy. Reactions that absorb energy are endothermic. Reactions that release energy are exothermic. 3.9: Energy and Chemical and Physical Change is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar- Agnew & Henry Agnew. 3.9.3 https://chem.libretexts.org/@go/page/47463 3.10: Temperature - Random Motion of Molecules and Atoms  Learning Objectives Identify the different between temperature and heat. Recognize the different scales used to measure temperature The concept of temperature may seem familiar to you, but many people confuse temperature with heat. Temperature is a measure of how hot or cold an object is relative to another object (its thermal energy content), whereas heat is the flow of thermal energy between objects with different temperatures. Temperature is a measure of the average kinetic energy of the particles in matter. In everyday usage, temperature indicates a measure of how hot or cold an object is. Temperature is an important parameter in chemistry. When a substance changes from solid to liquid, it is because there was in increase in the temperature of the material. Chemical reactions usually proceed faster if the temperature is increased. Many unstable materials (such as enzymes) will be viable longer at lower temperatures. Figure 3.10.1 : The glowing charcoal on the left represents high kinetic energy, while the snow and ice on the right are of much lower kinetic energy. Three different scales are commonly used to measure temperature: Fahrenheit (expressed as °F), Celsius (°C), and Kelvin (K). Thermometers measure temperature by using materials that expand or contract when heated or cooled. Mercury or alcohol thermometers, for example, have a reservoir of liquid that expands when heated and contracts when cooled, so the liquid column lengthens or shortens as the temperature of the liquid changes. Figure 3.10.2 : Daniel Gabriel Fahrenheit (left), Anders Celsius (center), and Lord Kelvin (right). The Fahrenheit Scale The first thermometers were glass and contained alcohol, which expanded and contracted as the temperature changed. The German scientist, Daniel Gabriel Fahrenheit used mercury in the tube, an idea put forth by Ismael Boulliau. The Fahrenheit scale was first developed in 1724 and tinkered with for some time after that. The main problem with this scale is the arbitrary definitions of temperature. The freezing point of water was defined as 32 F and the boiling point as 212 F. The Fahrenheit scale is typically not o o used for scientific purposes. The Celsius Scale The Celsius scale of the metric system is named after Swedish astronomer Anders Celsius (1701-1744). The Celsius scale sets the freezing point and boiling point of water at 0 C and 100 C respectively. The distance between those two points is divided into 100 o o equal intervals, each of which is one degree. Another term sometimes used for the Celsius scale is "centigrade" because there are 100 degrees between the freezing and boiling points of water on this scale. However, the preferred term is "Celsius". 3.10.1 https://chem.libretexts.org/@go/page/47464 The Kelvin Scale The Kelvin temperature scale is named after Scottish physicist and mathematician Lord Kelvin (1824-1907). It is based on molecular motion, with the temperature of 0 K, also known as absolute zero, being the point where all molecular motion ceases. The freezing point of water on the Kelvin scale is 273.15 K, while the boiling point is 373.15 K. Notice that there is no "degree" used in the temperature designation. Unlike the Fahrenheit and Celsius scales where temperatures are referred to as "degrees F " or "degrees C", we simply designate temperatures in the Kelvin scale as kelvins. Figure 3.10.1 : A Comparison of the Fahrenheit, Celsius, and Kelvin Temperature Scales. Because the difference between the freezing point of water and the boiling point of water is 100° on both the Celsius and Kelvin scales, the size of a degree Celsius (°C) and a kelvin (K) are precisely the same. In contrast, both a degree Celsius and a kelvin are 9/5 the size of a degree Fahrenheit (°F). (CC BY-SA-NC 3.0; anonymous) Converting Between Scales The Kelvin is the same size as the Celsius degree, so measurements are easily converted from one to the other. The freezing point of water is 0°C = 273.15 K; the boiling point of water is 100°C = 373.15 K. The Kelvin and Celsius scales are related as follows: T (in °C) + 273.15 = T (in K) (3.10.1) T (in K) − 273.15 = T (in °C) (3.10.2) Degrees on the Fahrenheit scale, however, are based on an English tradition of using 12 divisions, just as 1 ft = 12 in. The relationship between degrees Fahrenheit and degrees Celsius is as follows: where the coefficient for degrees Fahrenheit is exact. (Some calculators have a function that allows you to convert directly between °F and °C.) There is only one temperature for which the numerical value is the same on both the Fahrenheit and Celsius scales: −40°C = −40°F. The relationship between the scales is as follows: (°F − 32) °C = (3.10.3) 1.8 °F = 1.8 × (°C ) + 32 (3.10.4)  Example 3.10.1: Temperature Conversions A student is ill with a temperature of 103.5°F. What is her temperature in °C and K? Solution Converting from Fahrenheit to Celsius requires the use of Equation 3.10.3: (103.5°F − 32) °C = (3.10.1) 1.8 = 39.7 °C (3.10.2) Converting from Celsius to Kelvin requires the use of Equation 3.10.1: 3.10.2 https://chem.libretexts.org/@go/page/47464  K = 39.7 °C + 273.15 (3.10.3) = 312.9 K (3.10.4)  Exercise 3.10.1 Convert each temperature to °C and °F. a. the temperature of the surface of the sun (5800 K) b. the boiling point of gold (3080 K) c. the boiling point of liquid nitrogen (77.36 K) Answer (a) 5527 K, 9980 °F Answer (b) 2807 K, 5084 °F Answer (c) -195.79 K, -320.42 °F Summary Three different scales are commonly used to measure temperature: Fahrenheit (expressed as °F), Celsius (°C), and Kelvin (K). Contributions & Attributions 3.10: Temperature - Random Motion of Molecules and Atoms is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.10.3 https://chem.libretexts.org/@go/page/47464 3.11: Temperature Changes - Heat Capacity If a swimming pool and wading pool, both full of water at the same temperature, were subjected to the same input of heat energy, the wading pool would certainly rise in temperature more quickly than the swimming pool. The heat capacity of an object depends on both its mass and its chemical composition. Because of its much larger mass, the swimming pool of water has a larger heat capacity than the wading pool. Heat Capacity and Specific Heat Different substances respond to heat in different ways. If a metal chair sits in the bright sun on a hot day, it may become quite hot to the touch. An equal mass of water in the same sun will not become nearly as hot. We would say that water has a high heat capacity (the amount of heat required to raise the temperature of an object by 1 C). Water is very resistant to changes in o temperature, while metals in general are not. The specific heat of a substance is the amount of energy required to raise the temperature of 1 gram of the substance by 1 C. The symbol for specific heat is c , with the p subscript referring to the fact that o p specific heats are measured at constant pressure. The units for specific heat can either be joules per gram per degree (J/g C) or o calories per gram per degree (cal/g C) (Table 3.11.1). This text will use J/g C for specific heat. o o heat specific heat = o mass × cal/g C Notice that water has a very high specific heat compared to most other substances. Table 3.11.1 : Specific Heat Capacities Specific Heat Capacity Specific Heat Capacity Substance Substance at 25oC in J/g oC at 25oC in J/g oC H 2 gas 14.267 steam @ 100oC 2.010 He gas 5.300 vegetable oil 2.000 H O(l) 2 4.184 sodium 1.23 lithium 3.56 air 1.020 ethyl alcohol 2.460 magnesium 1.020 ethylene glycol 2.200 aluminum 0.900 ice @ 0oC 2.010 concrete 0.880 steam @ 100oC 2.010 glass 0.840 Water is commonly used as a coolant for machinery because it is able to absorb large quantities of heat (see table above). Coastal climates are much more moderate than inland climates because of the presence of the ocean. Water in lakes or oceans absorbs heat from the air on hot days and releases it back into the air on cool days. Figure 3.11.1 : This power plant in West Virginia, like many others, is located next to a large lake so that the water from the lake can be used as a coolant. Cool water from the lake is pumped into the plant, while warmer water is pumped out of the plant and back into the lake. 3.11.1 https://chem.libretexts.org/@go/page/47469 Summary Heat capacity is the amount of heat required to raise the temperature of an object by 1 C). o The specific heat of a substance is the amount of energy required to raise the temperature of 1 gram of the substance by 1 o. C 3.11: Temperature Changes - Heat Capacity is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.11.2 https://chem.libretexts.org/@go/page/47469 3.12: Energy and Heat Capacity Calculations  Learning Objectives To relate heat transfer to temperature change. Heat is a familiar manifestation of transferring energy. When we touch a hot object, energy flows from the hot object into our fingers, and we perceive that incoming energy as the object being “hot.” Conversely, when we hold an ice cube in our palms, energy flows from our hand into the ice cube, and we perceive that loss of energy as “cold.” In both cases, the temperature of the object is different from the temperature of our hand, so we can conclude that differences in temperatures are the ultimate cause of heat transfer. The specific heat of a substance can be used to calculate the temperature change that a given substance will undergo when it is either heated or cooled. The equation that relates heat (q) to specific heat (c ), mass (m), and temperature change (ΔT ) is shown p below. q = cp × m × ΔT The heat that is either absorbed or released is measured in joules. The mass is measured in grams. The change in temperature is given by ΔT = T − T , where T is the final temperature and T is the initial temperature. f i f i Every substance has a characteristic specific heat, which is reported in units of cal/g °C or cal/g K, depending on the units used to express ΔT. The specific heat of a substance is the amount of energy that must be transferred to or from 1 g of that substance to change its temperature by 1°. Table 3.12.1 lists the specific heats for various materials. Table 3.12.1 : Specific Heats of Some Common Substances Substance Specific Heat (J/g o C) Water (l) 4.18 Water (s) 2.06 Water (g) 1.87 Ammonia (g) 2.09 Ethanol (l) 2.44 Aluminum (s) 0.897 Carbon, graphite (s) 0.709 Copper (s) 0.385 Gold (s) 0.129 Iron (s) 0.449 Lead (s) 0.129 Mercury (l) 0.140 Silver (s) 0.233 The direction of heat flow is not shown in heat = mcΔT. If energy goes into an object, the total energy of the object increases, and the values of heat ΔT are positive. If energy is coming out of an object, the total energy of the object decreases, and the values of heat and ΔT are negative.  Example 3.12.1 A 15.0 g piece of cadmium metal absorbs 134 J of heat while rising from o 24.0 C to o 62.7 C. Calculate the specific heat of cadmium. Solution 3.12.1 https://chem.libretexts.org/@go/page/47467  Step 1: List the known quantities and plan the problem. Known Heat = q = 134 J Mass = m = 15.0 g o o o ΔT = 62.7 C − 24.0 C = 38.7 C Unknown c of cadmium =? J/g o p C The specific heat equation can be rearranged to solve for the specific heat. Step 2: Solve. q 134 J o cp = = = 0.231 J/g C o m × ΔT 15.0 g × 38.7 C Step 3: Think about your result. The specific heat of cadmium, a metal, is fairly close to the specific heats of other metals. The result has three significant figures. Since most specific heats are known (Table 3.12.1), they can be used to determine the final temperature attained by a substance when it is either heated or cooled. Suppose that a 60.0 g of water at 23.52 C was cooled by the removal of 813 J of heat. The o change in temperature can be calculated using the specific heat equation: q 813 J o ΔT = = = 3.24 C o cp × m 4.18 J/g C × 60.0 g Since the water was being cooled, the temperature decreases. The final temperature is: o o o Tf = 23.52 C − 3.24 C = 20.28 C  Example 3.12.2 What quantity of heat is transferred when a 150.0 g block of iron metal is heated from 25.0°C to 73.3°C? What is the direction of heat flow? Solution We can use heat = mcΔT to determine the amount of heat, but first we need to determine ΔT. Because the final temperature of the iron is 73.3°C and the initial temperature is 25.0°C, ΔT is as follows: ΔT = Tfinal − Tinitial = 73.3°C − 25.0°C = 48.3°C The mass is given as 150.0 g, and Table 7.3 gives the specific heat of iron as 0.108 cal/g °C. Substitute the known values into heat = mcΔT and solve for amount of heat: cal ∘ heat = (150.0 g) (0.108 ) (48.3 C) = 782 cal ∘ g⋅ C Note how the gram and °C units cancel algebraically, leaving only the calorie unit, which is a unit of heat. Because the temperature of the iron increases, energy (as heat) must be flowing into the metal.  Exercise 3.12.1 What quantity of heat is transferred when a 295.5 g block of aluminum metal is cooled from 128.0°C to 22.5°C? What is the direction of heat flow? Answer Heat leaves the aluminum block. 3.12.2 https://chem.libretexts.org/@go/page/47467   Example 3.12.2 A 10.3 g sample of a reddish-brown metal gave off 71.7 cal of heat as its temperature decreased from 97.5°C to 22.0°C. What is the specific heat of the metal? Can you identify the metal from the data in Table 3.12.1? Solution The question gives us the heat, the final and initial temperatures, and the mass of the sample. The value of ΔT is as follows: ΔT = Tfinal − Tinitial = 22.0°C − 97.5°C = −75.5°C If the sample gives off 71.7 cal, it loses energy (as heat), so the value of heat is written as a negative number, −71.7 cal. Substitute the known values into heat = mcΔT and solve for c: −71.7 cal = (10.3 g)(c)(−75.5°C) −71.7 cal c= ∘ (10.3 g)(−75.5 C) c = 0.0923 cal/g °C This value for specific heat is very close to that given for copper in Table 7.3.  Exercise 3.12.2 A 10.7 g crystal of sodium chloride (NaCl) has an initial temperature of 37.0°C. What is the final temperature of the crystal if 147 cal of heat were supplied to it? Answer Summary Specific heat calculations are illustrated. 3.12: Energy and Heat Capacity Calculations is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.12.3 https://chem.libretexts.org/@go/page/47467 3.E: Matter and Energy (Exercises) 3.1: In Your Room 3.2: What Is Matter? 1. What is matter? 2. What does weight mean? 3. In this chapter, we'll learn about atoms, which are the building blocks of all matter in the universe. As of 2011, scientists only know of 118 different types of atoms. How do you think it's possible to generate so many different forms of matter using only 118 types of building blocks? 4. Which do you think has more matter, a cup of water or a cup of mercury? Explain. 5. Decide whether each of the following statements is true or false. a. Mass and weight are two words for the same concept. b. Molecules are bonded together to form atoms. c. Alchemists couldn't make gold out of common metals because gold is an element. d. The symbol for Gold in the periodic table is Gd. 6. Would you have more mass on the moon or on Earth? 7. Would you have more weight on the moon or on Earth? The force of gravity is stronger on the Earth than it is on the moon. 8. Match the following terms with their meaning. Terms Definitions (a) Mass a. a measure of the total quantity of matter in an object (b) Volume b. a measure of how strongly gravity pulls on an object (c) Weight c. a measure of the space occupied by an object 9. For the following statements, circle all of the options that apply: Mass depends on… (a) the total quantity of matter (b) the temperature (c) the location (d) the force of gravity Volume depends on… (a) the total quantity of matter (b) the temperature (c) the object's shape (independent of size) (d) the object's size (independent of shape) Weight depends on… (a) the total quantity of matter (b) the temperature (c) the location (d) the force of gravity 3.3: Classifying Matter According to Its State: Solid, Liquid, and Gas 3.4 Classifying Matter According to Its Composition 3.5: Differences in Matter: Physical and Chemical Properties 3.6: Changes in Matter: Physical and Chemical Changes 3.E.1 https://chem.libretexts.org/@go/page/52963 3.7: Conservation of Mass: There is No New Matter 3.8: Energy 1. Classify each of the following as energy primarily transferred as heat, or energy primarily transferred as work: a. The energy transferred from your body to a shopping cart as you push the shopping cart down the aisle. b. The energy transferred from a wave to your board when you go surfing. c. The energy transferred from the flames to your hotdog when you cook your hotdog over a campfire. 2. Decide whether each of the following statements is true or false: a. When heat is transferred to an object, the object cools down. b. Any time you raise the temperature of an object, you have done work. c. Any time you move an object by applying force, you have done work. d. Any time you apply force to an object, you have done work. 3. Rank the following scenarios in order of increasing work: a. You apply 100 N of force to a boulder and successfully move it by 2 m. b. You apply 100 N of force to a boulder and successfully move it by 1 m. c. You apply 200 N of force to a boulder and successfully move it by 2 m. d. You apply 200 N of force to a boulder but cannot move the boulder. 4. In science, a vacuum is defined as space that contains absolutely no matter (no molecules, no atoms, etc.) Can energy be transferred as heat through a vacuum? Why or why not? 5. Classify each of the following energies as kinetic energy or potential energy: a. The energy in a chocolate bar. b. The energy of rushing water used to turn a turbine or a water wheel. c. The energy of a skater gliding on the ice. d. The energy in a stretched rubber band. 6. Decide which of the following objects has more kinetic energy: a. A 200 lb. man running at 6 mph or a 200 lb. man running at 3 mph. b. A 200 lb. man running at 7 mph or a 150 lb. man running at 7 mph. c. A 400 lb. man running at 5 mph or a 150 lb. man running at 3 mph. 7. A car and a truck are traveling along the highway at the same speed. a. If the car weighs 1500 kg and the truck weighs 2500 kg, which has more kinetic energy, the car or the truck? b. Both the car and the truck convert the potential energy stored in gasoline into the kinetic energy of motion. Which do you think uses more gas to travel the same distance, the car or the truck? 8. You mix two chemicals in a beaker and notice that as the chemicals react, the beaker becomes noticeably colder. Which chemicals have more chemical potential energy, those present at the start of the reaction or those present at the end of the reaction? 3.9: Energy and Chemical and Physical Change 3.10: Temperature: Random Motion of Molecules and Atoms 3.11: Temperature Changes: Heat Capacity 3.12: Energy and Heat Capacity Calculations 1. A pot of water is set on a hot burner of a stove. What is the direction of heat flow? 2. Some uncooked macaroni is added to a pot of boiling water. What is the direction of heat flow? 3. How much energy in calories is required to heat 150 g of H2O from 0°C to 100°C? 4. How much energy in calories is required to heat 125 g of Fe from 25°C to 150°C? 5. If 250 cal of heat were added to 43.8 g of Al at 22.5°C, what is the final temperature of the aluminum? 6. If 195 cal of heat were added to 33.2 g of Hg at 56.2°C, what is the final temperature of the mercury? 3.E.2 https://chem.libretexts.org/@go/page/52963  7. A sample of copper absorbs 145 cal of energy, and its temperature rises from 37.8°C to 41.7°C. What is the mass of the copper? 8. A large, single crystal of sodium chloride absorbs 98.0 cal of heat. If its temperature rises from 22.0°C to 29.7°C, what is the mass of the NaCl crystal? 9. If 1.00 g of each substance in Table 7.3 were to absorb 100 cal of heat, which substance would experience the largest temperature change? 10. If 1.00 g of each substance in Table 7.3 were to absorb 100 cal of heat, which substance would experience the smallest temperature change? 11. Determine the heat capacity of a substance if 23.6 g of the substance gives off 199 cal of heat when its temperature changes from 37.9°C to 20.9°C. 12. What is the heat capacity of gold if a 250 g sample needs 133 cal of energy to increase its temperature from 23.0°C to 40.1°C? Answers 1. Heat flows into the pot of water. 3. 15,000 cal 5. 49.0°C 7. 404 g 9. Mercury would experience the largest temperature change. 11. 0.496 cal/g °C 3.E: Matter and Energy (Exercises) is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 3.E.3 https://chem.libretexts.org/@go/page/52963