BIO 1010 Chapter 2 - Chemistry

BIO 1010 Chapter 2 - Chemistry
Why study chemistry? Chemicals are the stuff that make up our bodies and those of other organisms They make up the physical environment as well The ordering of atoms into molecules represents the lowest level of biological organization Therefore, to understand life, it is important to understand the basic concepts of chemistry Life results from the ordering of atoms into molecules and the interactions of these molecules within cells. Chemicals are at the base level of biological hierarchy They are arranged into higher and higher levels of structural organization Arrangement eventually leads to formation of living organisms Corpse flower

Devil’s Gardens Why are there 800 year old stands of only lemon trees in the middle of the Amazon rain forest? Local legend holds that the devils are cultivated by an evil spirit called Chuyachaqui The lemon trees provide hollow stems for lemon ant nest sites In order to preserve their home, the ants inject a chemical into the other trees -The chemical, formic acid, causes the leaves of the tree to die and fall off -Long term exposure to the ants and their chemical kills the neighboring trees *Take any biological system apart and you eventually end up at the chemical level*

Some Basic Chemistry Take any biological system apart, and you eventually end up at the chemical level. Chemical reactions are always occurring in the human body and our environment.

1 2 3 4 5 6 10 9 7 8 Biosphere Ecosystems Communities Populations
Figure 1.2-3 1 Biosphere 2 Ecosystems 3 Communities 4 Populations 5 Organisms 6 Organ Systems and Organs Figure 1.2 Zooming in on life (step 3) 10 Molecules and Atoms 9 Organelles 7 Tissues Atom Nucleus 8 Cells 4

Are Chemicals Bad or Good for you?
Misconceptions and Concerns about chemicals Appreciation for chemical nature of our bodies and our world Potential harms and benefits Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.”

Several chemicals are added to food for a variety of reasons
Help preserve it Make it more nutritious Make it look better Figure 2.2B Nutrition facts from a fortified cereal. Ask students to identity a few elements Talk about Iron experiment below- if they are bored they can try to following experiment  Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example.

Matter: Elements and Compounds
Matter is anything that occupies space and has mass. Matter is found on Earth in three physical states: solid, liquid, and gas. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 7

Matter is composed of chemical elements. An element is a substance that cannot be broken down into other substances by chemical reactions. There are 92 naturally occurring elements on Earth. All of the elements are listed in the periodic table. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 8

Elements & Atoms Element: a substance that can neither be broken down nor converted to another substance by chemical reactions. Pure substance that contains only one type of atom

C 6 12 Atomic number (number of protons) Element symbol Mass number
Figure 2.1a Atomic number (number of protons) 6 C Element symbol 12 Mass number (number of protons plus neutrons) H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Figure 2.1 Abbreviated periodic table of the elements (part 1) Up close view of an element 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. Cs Ba La Hf Ta W Re Os Ir Pt Au Hg TI Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg Cn Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr 10

Thermometers, dental fillings and batteries
Figure 2.1b Thermometers, dental fillings and batteries Mercury in the air settles into water. It can pass through the food chain and build up in fish, shellfish and animals that eat fish. Exposure to high levels can damage the brain and kidneys. Can damage the kidneys and the nervous system, and interfere with development of the brain in the very young children Mercury Figure 2.1 Abbreviated periodic table of the elements (part 2) Metallic mercury is a shiny, silver-white, odorless liquid. If heated, it is a colorless, odorless gas. It also combines with other elements to form powders or crystals. Mercury is in many products. Metallic mercury is used in thermometers, dental fillings and batteries. Mercury salts may be used in skin creams and ointments. It's also used in many industries. Mercury in the air settles into water. It can pass through the food chain and build up in fish, shellfish and animals that eat fish. The nervous system is sensitive to all forms of mercury. Exposure to high levels can damage the brain and kidneys. Pregnant women can pass the mercury in their bodies to their babies. Mercury poisoning can damage the kidneys and the nervous system, and interfere with the development of the brain in unborn children and very young children.” 11

Figure 2.1c Figure 2.1 Abbreviated periodic table of the elements (part 3) Deficiencies of copper can cause premature hair graying, sterility and premature wrinkling of the skin. Deficiencies of copper can cause premature hair graying, sterility and premature wrinkling of the skin. 12

Houses built before ‘78, present toys, pipes, faucets
Figure 2.1d Affects nearly all system: Hearing loss, kidney problems, lower IQ, children are more vulnerable because they tend to place things in their mouths. Houses built before ‘78, present toys, pipes, faucets Lead Figure 2.1 Abbreviated periodic table of the elements (part 4) Lead exposure can affect nearly every system in the body It builds up slowly, sometimes signs are not obvious, it can cause hearing loss, kidney problems lower IQ, children are more vulnerable because they tend to place things in their mouthts. Houses built before 1978 contained lead. Even though now gasoline and paint does not contain lead, it is still present in new toys, plumbing, pipes, faucets to name a few. 13

Twenty-five elements are essential to people. Four elements make up about 96% of the weight of most cells: oxygen, carbon, hydrogen, and nitrogen. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 14

ELEMENTS, ATOMS, AND MOLECULES
What is matter? There are 92 chemical elements in nature Life requires 25 essential elements; some are called trace elements. Living organisms are composed of matter, which is anything that occupies space and has mass (weight) Matter is found on earth in three physical states: solid, liquid and gas Matter is composed of chemical elements (a substance that cannot be broken down to other substances There are 92 elements in nature—only a few exist in a pure state Life requires 25 essential elements; some are called trace elements Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.”

Trace elements: less than 0.01%
Figure 2.2 Carbon (C): 18.5% Oxygen (O): 65.0% Calcium (Ca): 1.5% Phosphorus (P): 1.0% Potassium (K): 0.4% Sulfur (S): 0.3% Sodium (Na): 0.2% Chlorine (Cl): 0.2% Magnesium (Mg): 0.1% Hydrogen (H): 9.5% Trace elements: less than 0.01% Figure 2.2 Chemical composition of the human body by weight Boron (B) Manganese (Mn) Chromium (Cr) Molybdenum (Mo) Nitrogen (N): 3.3% Cobalt (Co) Selenium (Se) Copper (Cu) Silicon (Si) Fluorine (F) Tin (Sn) Iodine (I) Vanadium (V) Iron (Fe) Zinc (Zn) 16

Trace elements are required in only very small amounts and essential for life. An iodine deficiency causes goiter. Fluorine is added to dental products and drinking water and helps to maintain healthy bones and teeth. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 17

Trace elements are common additives to food and water
Without iron, your body cannot transport oxygen An iodine deficiency prevents production of thyroid hormones, resulting in goiter -Deficiencies of copper can cause premature hair graying, sterility and premature wrinkling of the skin. Although the body needs trace elements in very small amounts, we cannot live without them Several chemicals are added to food for a variety of reasons Help preserve it Make it more nutritious Make it look better Check out the “Nutrition Facts” label on foods and drinks you purchase Goiter in a Malaysian woman, a symptom of iodine deficiency. Another example:. Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires.

Figure 2.3 Figure 2.3 The relationship between diet and goiter 19

Elements can combine to form compounds. Compounds are substances that contain two or more elements in a fixed ratio. Common compounds include NaCl (table salt) and H2O (water). Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 20

Compound—a substance consisting of two or more different elements combined in a fixed ratio
+ Figure 2.3 The emergent properties of the edible compound sodium chloride. Compound—a substance consisting of two or more different elements combined in a fixed ratio There are many compounds that consist of only two elements Table salt (sodium chloride or NaCl) is an example Sodium is a metal, and chloride is a poisonous gas However, when chemically combined, an edible compound emerges 1. The text notes the unique properties of pure sodium, pure chlorine, and the compound sodium chloride formed when the two bond together. Consider challenging your students to think of other simple examples of new properties that result when a compound is formed (for example, water, formed from hydrogen and oxygen, and rust, formed from iron and oxygen). Sodium (metal) Chlorine (poisonous gas) Sodium Chloride (Kitchen salt)

Atoms Each element consists of one kind of atom.
An atom is the smallest unit of matter that still retains the properties of an element. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 22

Atoms consist of protons, neutrons, and electrons
cloud Electrons are kept in orbit by the attraction between the negatively charged electrons and positively charged protons Nucleus 2e– Atomic number: 2 (number of protons) Mass number: 4 (sum of the number of protons and neutrons in its nucleus Helium atom An atom is the smallest unit of matter that still retains the properties of a element Atoms are made of over a hundred subatomic particles, but only three are important for biological compounds Proton—has a single positive electrical charge Electron—has a single negative electrical charge Neutron—is electrically neutral 2 Protons Mass number = 4 2 Neutrons 2 Electrons

The Structure of Atoms Atoms are composed of subatomic particles.
A proton is positively charged. An electron is negatively charged. A neutron is electrically neutral. Most atoms have protons and neutrons packed tightly into the nucleus. The nucleus is the atom’s central core. Electrons orbit the nucleus. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 24

C 6 12 Atomic Number The atomic number determines which element it is.
(# of Protons in an atom) Mass number (atomic weight) ( # Protons + # Neutrons) 12 Atomic Symbol (Carbon) Have students turn to appendix B in their books Write some Elements with Mass and atomic numbers on the board and have students say how many protons, electrons and neutrons each has: (6/13 C), 11/22 Na+, 17/34 Cl-, 8/16 O2-) Write out element abbreviations with atomic and mass numbers for B, Li, Ca, N, Ne assuming the protons and neutrons are the same # protons = atomic number # electrons = atomic number (in an uncharged atom) # neutrons = mass number – atomic number

The Structure of Atoms Elements differ in the number of subatomic particles in their atoms. The number of protons, the atomic number, determines which element it is. Mass is a measure of the amount of material in an object. An atom’s mass number is the sum of the number of protons and neutrons in its nucleus. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 26

Isotopes Isotopes are alternate mass forms of an element. Isotopes
have the same number of protons and electrons but differ in their number of neutrons. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 27

Atomic Weight (Mass Number) May Vary
99% ~1% Minute quantities Naturally occurring carbon Stable Isotopes: nuclei stay intact forever, more or less Radioactive Isotopes: unstable, nucleus decays, giving off particles and energy Stable isotopes Radioactive: the nucleus decays Isotope: a variant form of an atom with same number of protons but different numbers of neutrons

Another example of isotopes
Atomic number: Mass number:

Table 2.1 Table 2.1 Isotopes of Carbon 30

Isotopes The nucleus of a radioactive isotope decays spontaneously, giving off particles and energy. Radioactive isotopes have many uses in research and medicine. They can be used to determine the fate of atoms in living organisms. They are used in PET scans to diagnose heart disorders and some cancers. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 31

Radioisotopes in Medicine & Research
when radioactive compounds are used in metabolic processes, they act as radioactive tracers Normal Alzheimer’s Living cells cannot distinguish between isotopes of the same element Therefore, when radioactive compounds are used in metabolic processes, they act as tracers Radioactivity can be detected by instruments With instruments, the fate of radioactive tracers can be monitored in living organisms Biologists use radioactive tracers in research Radioactive 14C was used to show the route of 14CO2 in formation of sugar during plant photosynthesis An imaging instrument that uses positron-emission tomography (PET) detects the location of injected radioactive materials

Use of Carbon Isotopes in Determining Ancient Maya Land Use
C isotope studies can be used to identify areas with histories of vegetation change How the Maya were able to sustain such large populations is still a question today. Jungle: C3 Plants Discriminate more against CO2 containing 13C isotope Maize: C4 Plants 1.1% of all natural carbon on earth C13 isotope. Plants discriminate against c13 because it is slightly heavier than c12 (2.3%). Plants that utilize the Hatch-Slack pathway (C4) for carbon reduction, discriminate less than those that use the Calvin cycle (C3) Once C13 is incorporated into the plant, it can be deposited into the soil by plant decomposition Taking soil cores (up to several meters down) and analyzing the amount of C13 isotope present in the soil at different levels gives an idea to where maize was cultivated by the ancient Maya.

Why is the energy emitted in radioactive decay hazardous? .
Uncontrolled exposure to radioactive isotopes can cause damage to some molecules in a living cell, especially DNA Chemical bonds are broken by the emitted energy, which causes abnormal bonds to form One of the most serious environmental threats of radioactivity is radioactive fallout from nuclear accidents. A 1979 accident in the United States at the Three Mile Island nuclear power plant led to strict regulations about safety. In the Soviet Union on Saturday, April 26, 1986, the Number 4 reactor at the Chernobyl nuclear power plant blew apart and blasted through the surrounding concrete containment structure. Two workers died immediately in the blast, but over the following months 30 others died, most of them firefighters who battled to prevent the fire from spreading to the reactor in the next building. All died from radiation burns and radiation sickness following exposure to strong gamma and beta radiation.

Radioactive Isotopes Can Harm
Uncontrolled exposure to radioactive isotopes can harm living organisms by damaging DNA. The 1986 Chernobyl nuclear accident released large amounts of radioactive isotopes. Naturally occurring radon gas may cause lung cancer. nuclear power plant blew apart and blasted through the surrounding concrete containment in chernobyl russia, two workers died immediately in the blast, but over the following months 30 others died, most of them firefighters who battled to prevent the fire from spreading to the reactor in the next building. All died from radiation burns and radiation sickness following exposure to strong gamma and beta radiation.

Electron arrangement determines the chemical properties of an atom
Only electrons are involved in chemical activity Electrons occur in energy levels called electron shells The number of electrons in the outermost shell determines the chemical properties of an atom Information about the distribution of electrons is found in the periodic table of the elements

Electron Arrangement and the Chemical Properties of Atoms
Electrons orbit the nucleus of an atom in specific electron shells. The farther an electron is from the nucleus, the greater its energy. The number of electrons in the outermost shell determines the chemical properties of an atom. Student Misconceptions and Concerns 1.The dangers posed by certain chemicals in our food and broader environment often mislead people to associate “Chemicals” with harm. People might not want “Chemicals” added to their food or in their environment. Students often fail to appreciate the chemical nature of our bodies and our world and the potential harm or benefits of naturally occurring chemistry. They often fail to understand why “Natural” does not necessarily mean “good.” (Consider presenting a long list of naturally occurring toxins to make this point.) Your class may benefit from a class discussion of these misconceptions about our attitudes toward “Chemicals.” 2. Students with limited backgrounds in chemistry and physics might struggle with basic concepts of mass, weight, compounds, elements, and isotopes. It may also be early in the semester when mature study habits have not yet developed. Consider passing along basic studying advice and tips to help students master these early chemistry concepts. In-class quizzes (graded or not) or a few homework problems will also provide reinforcing practice. 3.The half-lives of many radioactive substances, especially those used for dating fossils, might lead some students to expect very long periods of decay for any radioactive substance. This might even be alarming if students are someday asked to consume a radioactive substance for a medical test. However, some medically significant isotopes have relatively short half-lives. Radioactive iodine-131 is often used to diagnose or treat certain thyroid problems. Its half-life of 8 days means that it will decay quickly. Teaching Tips 1.Students might be interested in the following aside: One of the challenges of raising captive, exotic animals is meeting the unique dietary requirements of the species. A zoo might have trouble keeping a particular animal because science has not fully determined what trace elements the animal requires. 2.Many breakfast cereals are fortified with iron. A person can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imagining device should clearly reveal the iron particles stuck to the magnet. This short practical demonstration can help connect an abstract concept to a concrete example. 3. Students with limited backgrounds in chemistry will benefit from a discussion of Figure 2.8 and the differences and limitations of diagrams of atomic structures. The contrasts in Figure 2.8 are a good beginning of such a discussion. 4.Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton is as massive as a bowling ball, an electron will be the mass of a Life Saver candy. (This is calculated by considering a 15-pound bowling ball, a Life Saver with a mass of 0.12 ounces, and the textbook’s formula that an electron is 1/2,000 the mass of a proton.) 5. Here is an analogy regarding the size of a helium atom. If a helium atom were the size of Yankee Stadium, the nucleus would be about the size of a fly in center field, and the two electrons would be like tiny gnats buzzing around the stadium. This analogy helps to relate the great distances between parts of an atom. Consider modifying the analogy to any local stadium in your region. 6. After sharing teaching tips 4 and 5 above, consider asking your students to compare the mass of the gnat orbiting Yankee Stadium to the mass of the fly in center field. If a proton or neutron is about 2,000 times more massive than an electron, how does the mass of a helium nucleus compare to the mass of one of its electrons? 7.Have your students try to calculate the number of covalent bonds possible for a variety of atoms. Then provide the students with a list of elements and the number of outer electrons for each and have them make predictions about the chemical formula for many types of molecules. 8. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) 9. As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 37

When all the orbitals are filled, the element is inactive (inert)
The number of electrons in the valence shell (outermost shell) determines the chemical properties of the atom An atom may have one, two, or three electron shells The number of electrons in the outermost shell determines the chemical properties of the atom The first shell is full with two electrons, whereas the second and third will hold up to eight electrons Noble gases have filled outer electron shells and are chemically unreactive Elements whose outer shells contain unfilled orbitals are chemically reactive When all the orbitals are filled, the element is inactive (inert)

An uncharged atom of gold has an atomic number of 79 and a mass number of 197. This atom has _________ protons, _______ neutrons, and __________ electrons. a. 79 …118 …79 b. 118 …79 …118 c. 276 …118 …79 d. 79 …276 …79 Answer a

The most abundant element found in the human body by weight is _________.
a. oxygen b. carbon c. fluorine d. hydrogen Answer: a

Isotopes of atoms differ in their number of __________.
a. neutrons b. electrons c. protons d. atomic nuclei Answer: a

Intro to Atoms, Ions etc http://www.youtube.com/watch?v=UhXS-iR2L3U

Chemical Bonding and Molecules
Chemical reactions enable atoms to give up or acquire electrons, completing their outer shells. Chemical reactions usually result in atoms staying close together and being held together by attractions called chemical bonds. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. © 2013 Pearson Education, Inc. 43

Ionic bonds are attractions between ions of opposite charge
Complete transfer of electrons. Attraction between ions of opposite charge Transfer of electron The octet rule: atoms are most stable when their outer most energy shells are either full or empty Na Sodium atom Cl Chlorine atom Figure 2.7A Formation of an ionic bond, producing sodium chloride. An ion is an atom or molecule with an electrical charge resulting from gain or loss of electrons When an electron is lost, a positive charge results; when one is gained, a negative charge results Two ions with opposite charges attract each other When the attraction holds the ions together, it is called an ionic bond Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. Ions: atoms electrically charged as a result of gaining or losing electrons

Sodium chloride (NaCl)
Transfer of electron + – Na Sodium atom Cl Chlorine atom Na+ Sodium ion Cl– Chloride ion Sodium chloride (NaCl) Figure 2.7A Formation of an ionic bond, producing sodium chloride. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players. Na+ Cl–

Covalent bonds join atoms into molecules through electron sharing
A covalent bond results when atoms share outer-shell electrons A molecule is formed when atoms are held together by covalent bonds Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (covalent, ionic, and hydrogen). Evaluating the accuracy of potential analogies requires careful analysis of the chemical bonding relationships and practices critical thinking skills. Small groups might provide immediate critiques before passing along analogies for the entire class to consider. The following is one example. Ionic and covalent bonds are different types of relationships. Consider this analogy. A woman taking out a loan has a specific relationship to her bank. She owes the bank money, something she got from the bank. A man shares an office with another man. Both look out the same window and answer the same phone. Ionic bonds are like a bank loan, in which something is borrowed. Covalent bonds are like sharing an office, with items (electrons) shared by both members of the relationship. After presenting this analogy, ask your students to modify the office analogy to represent a polar covalent bond. (Perhaps one man in the office sits closer to the window and the phone.) As noted in the text, chemical reactions do not create or destroy matter. Instead, they rearrange the structure and form new relationships. This is much like shuffling and dealing cards. When playing poker, cards are not created nor destroyed. Instead, new combinations are formed as the cards are dealt to the players.

Covalent Bonds The number of covalent bonds an atom can form is equal to the number of additional electrons it needs to fill its valence shell A single bond forms when two electrons are shared between two atoms A double bond forms when four electrons are shared between two atoms A triple bond forms when ______ electrons are shared between two atoms

Electrons are Unequally Shared in Polar Covalent Bonds
Polar molecule: opposite charges on opposite ends Water molecule Fig 2.8 Oxygen draws shared electrons towards itself Oxygen becomes “a little” negative Hydrogens become “a little” positive Results in Hydrogen Bonds between neighboring molecules Weak electrical attractions Water is a polar molecule because the electrons of the covalent bonds are not shared equally between the oxygen and hydrogen. Atoms in a covalently bonded molecule continually compete for shared electrons The attraction (pull) for shared electrons is called electronegativity More electronegative atoms pull harder In molecules of only one element, the pull toward each atom is equal, because each atom has the same electronegativity The bonds formed are called nonpolar covalent bonds

Hydrogen bonds are weak bonds important in the chemistry of life
Water molecules are electrically attracted to oppositely charged regions on neighboring molecules Because the positively charged region is always a hydrogen atom, the bond is called a hydrogen bond Hydrogen bond Figure 2.10 Hydrogen bonds between water molecules. Hydrogen bonds, though weak, without them, we would not have life as we know it. Next session devoted to the special properties of water, of which hydrogen bonding plays a significant role.

Chemical bonds and attractive forces
Type Chemical basis Strength Example Covalent bonds Atoms share electron pairs Strong Hydrocarbons, methane Ionic bonds Atoms donate one or more electrons to other atom of opposite charge Moderate Sodium chloride Sodium iodide Hydrogen bonds Atoms with partial negative charge attract hydrogen atoms Weak Water, DNA

WATER’S LIFE-SUPPORTING PROPERTIES

WATER AND LIFE Life on Earth began in water and evolved there for 3 billion years. Modern life remains tied to water. Your cells are composed of 70–95% water. The abundance of water is a major reason Earth is habitable. Student Misconceptions and Concerns 1. Students are unlikely to have carefully considered the four special properties of water that are apparent in our lives. However, these properties are of great biological significance and are often familiar. The connections between these properties and personal experiences can invest great meaning into a discussion of water’s properties. A homework assignment asking for three examples of each of these properties in each student’s experiences will require reflection and potentially produce meaningful illustrations. Similarly, quizzes or exam questions matching examples of each property to a list of the properties requires high-level evaluative analysis. 2. Students at all levels struggle with the distinction between heat and temperature. Students might also expect that all ice is about the same temperature, 0ºC. Redefining and correcting misunderstandings often takes more class time and energy than introducing previously unknown concepts. Teaching Tips 1. Here is a way to have your students think about the “sticky” nature of water in their lives. Ask them to consider the need for a towel after a shower or a bath. Once we get out of the shower or bath, we have left the source of water. So why do we need the towel? The reason we need a towel to dry off is that water is still clinging to our bodies because water molecules are polar. The molecules on cell surfaces are also polar, so our skin and the water both “stick” to each other. 2. Some students may be intrigued if you tell them that you too can stand on the surface of water—when it is frozen. Thus, it is necessary to note a liquid water surface when talking about surface tension. 3. Have students compare the seasonal ranges of temperatures of Anchorage and Fairbanks, Alaska. (Many websites provide weather information about various cities.) These two northern cities have large differences in their annual temperature ranges. Make the point that the coastal location of Anchorage moderates the temperature. 4. Several versions of the following analogy are easy to relate to students. (a) Ask students how the average on an exam would be affected if the brightest students did not take the test. (b) The authors note that the performance of a track team would drop if the fastest sprinters didn’t compete. In both analogies, removing the top performers lowers the average just as the evaporation of the most active water molecules cools the evaporative surface. 5. “It is not the heat, it is the humidity.” The efficiency of evaporative cooling is affected by humidity. As humidity rises, the rate of evaporation decreases, making it more difficult to cool our heat-generating bodies on a warm and humid summer day. 6. Ask your students if the ocean levels would change if ice did not float or if all the floating ice in the oceans were to melt. They can try this experiment to find out. Place several large chunks of ice in a glass and fill the glass up completely with water to the top rim. Thus, the ice cubes should be sticking up above the top of the filled glass. Will the glass overflow when the ice melts? (No.) This phenomenon is important when we consider the potential consequences of global warming. If floating glaciers melt, ocean levels will not be affected. However, if the ice over land melts, we can expect increased ocean levels. 7. Challenge students to think of other consequences from the expansion of water when it forms ice. (These include the ability to widen cracks in rocks, roads, and sidewalks!) 8. A simple demonstration of a solute dissolving in a solvent can focus students’ attention on the process when discussing solutions. Using colored water and white sugar or salt may make it easier to see and reference while you are discussing the process. Such simple visual aids add life to a lecture. (You might also add corn oil to the top of the solution to demonstrate the properties of hydrophobic substances, and challenge your class to explain why oil and water do not mix.) 9. Discussions of pH are enhanced by lab activities that permit students to test the pH of everyday items (foods and household solutions). If students do not have opportunities to conduct such tests in lab, consider testing a few items during your class (pH paper or a basic pH meter will, of course, be necessary). 10. The Environmental Protection Agency website includes good information about acid precipitation and teaching ideas. 11. The SETI (Search for Extraterrestrial Intelligence) Institute’s Mission is “to explore, understand and explain the origin, nature, and prevalence of life in the universe” ( Your students might enjoy exploring this respected scientific organization. © 2013 Pearson Education, Inc. 52

Water’s Life-Supporting Properties
The polarity of water molecules and the hydrogen bonding that results explain most of water’s life-supporting properties. Water molecules stick together. Water has a strong resistance to change in temperature. Frozen water floats. Water is a common solvent for life. Student Misconceptions and Concerns 1. Students are unlikely to have carefully considered the four special properties of water that are apparent in our lives. However, these properties are of great biological significance and are often familiar. The connections between these properties and personal experiences can invest great meaning into a discussion of water’s properties. A homework assignment asking for three examples of each of these properties in each student’s experiences will require reflection and potentially produce meaningful illustrations. Similarly, quizzes or exam questions matching examples of each property to a list of the properties requires high-level evaluative analysis. 2. Students at all levels struggle with the distinction between heat and temperature. Students might also expect that all ice is about the same temperature, 0ºC. Redefining and correcting misunderstandings often takes more class time and energy than introducing previously unknown concepts. Teaching Tips 1. Here is a way to have your students think about the “sticky” nature of water in their lives. Ask them to consider the need for a towel after a shower or a bath. Once we get out of the shower or bath, we have left the source of water. So why do we need the towel? The reason we need a towel to dry off is that water is still clinging to our bodies because water molecules are polar. The molecules on cell surfaces are also polar, so our skin and the water both “stick” to each other. 2. Some students may be intrigued if you tell them that you too can stand on the surface of water—when it is frozen. Thus, it is necessary to note a liquid water surface when talking about surface tension. 3. Have students compare the seasonal ranges of temperatures of Anchorage and Fairbanks, Alaska. (Many websites provide weather information about various cities.) These two northern cities have large differences in their annual temperature ranges. Make the point that the coastal location of Anchorage moderates the temperature. 4. Several versions of the following analogy are easy to relate to students. (a) Ask students how the average on an exam would be affected if the brightest students did not take the test. (b) The authors note that the performance of a track team would drop if the fastest sprinters didn’t compete. In both analogies, removing the top performers lowers the average just as the evaporation of the most active water molecules cools the evaporative surface. 5. “It is not the heat, it is the humidity.” The efficiency of evaporative cooling is affected by humidity. As humidity rises, the rate of evaporation decreases, making it more difficult to cool our heat-generating bodies on a warm and humid summer day. 6. Ask your students if the ocean levels would change if ice did not float or if all the floating ice in the oceans were to melt. They can try this experiment to find out. Place several large chunks of ice in a glass and fill the glass up completely with water to the top rim. Thus, the ice cubes should be sticking up above the top of the filled glass. Will the glass overflow when the ice melts? (No.) This phenomenon is important when we consider the potential consequences of global warming. If floating glaciers melt, ocean levels will not be affected. However, if the ice over land melts, we can expect increased ocean levels. 7. Challenge students to think of other consequences from the expansion of water when it forms ice. (These include the ability to widen cracks in rocks, roads, and sidewalks!) 8. A simple demonstration of a solute dissolving in a solvent can focus students’ attention on the process when discussing solutions. Using colored water and white sugar or salt may make it easier to see and reference while you are discussing the process. Such simple visual aids add life to a lecture. (You might also add corn oil to the top of the solution to demonstrate the properties of hydrophobic substances, and challenge your class to explain why oil and water do not mix.) 9. Discussions of pH are enhanced by lab activities that permit students to test the pH of everyday items (foods and household solutions). If students do not have opportunities to conduct such tests in lab, consider testing a few items during your class (pH paper or a basic pH meter will, of course, be necessary). 10. The Environmental Protection Agency website includes good information about acid precipitation and teaching ideas. 11. The SETI (Search for Extraterrestrial Intelligence) Institute’s Mission is “to explore, understand and explain the origin, nature, and prevalence of life in the universe” ( Your students might enjoy exploring this respected scientific organization. © 2013 Pearson Education, Inc. 53

Biology and Society: More Precious than Gold
A drought is a period of abnormally dry weather that changes the environment and one of the most devastating disasters. © 2013 Pearson Education, Inc. 54

Cohesion: attraction between molecules (water= hydrogen bonds)
Figs 2.12, 2.13 Hydrogen bonding causes molecules to stick together, a property called cohesion Cohesion is much stronger for water than other liquids This is useful in plants that depend upon cohesion to help transport water and nutrients up the plant Cohesion is related to surface tension—a measure of how difficult it is to break the surface of a liquid Hydrogen bonds are responsible for surface tension Water molecules stick together as a result of hydrogen bonding Water displays emergent properties resulting from the hydrogen bond that orders water molecules. Cohesion is just one of these properties. Hydrogen bonds give water a high surface tension Capillary Action Surface Tension

Biology and Society: More Precious than Gold
Droughts can cause severe crop damage, shortages of drinking water, dust storms, famine, habitat loss, and mass migration. © 2013 Pearson Education, Inc. 56

More Precious than Gold
Biology and Society: More Precious than Gold Throughout human history, droughts have helped wipe out societies and even whole civilizations. Droughts are catastrophic because life cannot exist without water. © 2013 Pearson Education, Inc. 57

Water resistance to temperature change stabilizes ocean temperatures.
Evaporative cooling: when a substance evaporates, the surface of the liquid remaining behind cools down Heat: amount of energy associated with the movement of atoms and molecules in a body of matter Temperature: measures the intensity of the heat, average speed of molecules Resistance to Temperature Change -When water heats up, it takes energy to break hydrogen bonds, heat is absorbed only raising water by a few degrees -When water cools, hydrogen bonds reform, releasing heat (energy), water can release a lot of heat to surroundings while only dropping by a few degrees.

Water’s hydrogen bonds moderate temperature
Because of hydrogen bonding, water has a greater ability to resist temperature change than other liquids. The efficiency of evaporative cooling is affected by humidity. Figure 2.12 Evaporative cooling. Evaporative cooling is a process in which water molecules with the greatest energy leave. Perspiration is an example of a benefit from this process. When water molecules leave, the water molecule

Water Expands When Frozen
Fewer molecules than an equal volume of liquid Fig 2.15

What is a solution? Example What is a solvent? Example
What is an aqueous solution? What is a solute? Example What is a Solution?: liquid consisting of a homogenous mixture of two or more substances Aqueous solution: water is the solvent Solvent: The dissolving agent Solute: substance that is dissolved

Ionic bond between Na+ and Cl- holds ions together in a solid crystal
Water is a polar solvent: A polar molecule has opposite charges on opposite ends Ionic bond between Na+ and Cl- holds ions together in a solid crystal Dissolve in water: The chloride anion (-) attracts the (+) pole of water -Why don’t the Cl- anions pull the hydrogens or the Na+ cations pull the oxygen off water? -Water cannot dissolve hydrophilic, nonpolar substances Dissolved ions cannot re-associate into a solid The sodium cation (+) attracts the (-) pole of water

Water is the solvent of life
Figure 2.14 A crystal of salt (NaCl) dissolving in water. Ion in solution Salt crystal

The chemistry of life is sensitive to acidic and basic conditions
Ionization of Water: Water dissociates hydrogen ions H+ hydroxide ions OH- Bases: release OH- (or accept H+): decrease [H+] Sodium Hydroxide  Na+ & OH - Acids: release H+ (or accept OH-): increase [H+] Hydrochloric acid: HCl  H+ & Cl- (in your stomach) Some chemicals accept hydrogen ions and remove them from solution These chemicals are called bases For example, sodium hydroxide (NaOH) provides OH– that combines with H+ to produce H2O (water) This reduces the H+ concentration A few water molecules can break apart into ions Some are hydrogen ions (H+) Some are hydroxide ions (OH–) Both are extremely reactive A balance between the two is critical for chemical processes to occur in a living organism Fig 2.17

Basic solution Neutral solution Acidic solution
Figure 2.17a OH− OH− OH− H OH− H H OH− OH− H H H OH− OH− H OH− OH− H H OH− H H OH− H Figure 2.17 The pH scale (detail) Some chemicals accept hydrogen ions and remove them from solution These chemicals are called bases For example, sodium hydroxide (NaOH) provides OH– that combines with H+ to produce H2O (water) This reduces the H+ concentration A few water molecules can break apart into ions Some are hydrogen ions (H+) Some are hydroxide ions (OH–) Both are extremely reactive A balance between the two is critical for chemical processes to occur in a living organism Basic solution Neutral solution Acidic solution 65

The pH Scale A pH scale (pH = potential of hydrogen) is used to describe whether a solution is acidic or basic. pH ranges from 0 (most acidic) to 14 (most basic) A solution that is neither acidic or basic is neutral (pH = 7) pH Scale: pH ranges from 0 (most acidic) to 14 (most basic) A solution that is neither acidic or basic is neutral

Acid rain Normal rain: pH 5.6 Water reacting with pollutants (SO2, NO)
When we burn fossil fuels (gasoline and heating oil), air-polluting compounds and CO2 are released into the atmosphere Sulfur and nitrous oxides react with water in the air to form acids These fall to Earth as acid precipitation, which is rain, snow, or fog with a pH lower than 5.6 (normal rain is slightly acidic due to dissolved carbonic acid) Additional CO2 in the atmosphere contributes to the “greenhouse” effect and alters ocean chemistry by combining with seawater and producing acid. 1. The Environmental Protection Agency (EPA) website, includes useful information about acid rain and related science experiments. 2. In Module 2.16, the authors note that threats to water quality are addressed in Chapter 38. If your course does not include Chapter 38, consider covering some of those threats in your discussion of acid precipitation. Normal rain: pH 5.6

Buffers Act as H+ reservoirs
Buffers minimize changes in pH Buffers Act as H+ reservoirs Take up H+ ions when they are abundant, release them when they are scarce. Keep proton concentration steady. A buffer that maintains pH 7: accepts protons if pH is < 7, releases protons if pH is > 7

The Carbonic-Acid-Bicarbonate Buffer in Blood
When we exercise, we increase the H+ concentration in our blood stream: [H+] increases Blood pH decreases Our blood contains a buffering system When we exercise, we increase: Heart rate, systolic blood pressure, cardiac output, breathing rate More active metabolism: Produces CO2 and H+ in muscles Uses up oxygen Anaerobic processes that produce lactic acid into blood stream Blood contains more H+ and blood pH decreases

Chemical reactions make and break bonds, changing the composition of matter
You learned that the structure of atoms and molecules determines the way they behave Remember that atoms combine to form molecules Hydrogen and oxygen can react to form water 2 H2 O2 2 H2O The formation of water from hydrogen and oxygen is an example of a chemical reaction The reactants (H2 and O2) are converted to H2O, the product Organisms do not make water, but they do carry out a large number of chemical reactions that rearrange matter Photosynthesis is an example where plants drive a sequence of chemical reactions that produce glucose Reactants Product

You should now be able to
Describe the importance of chemical elements to living organisms Explain the formation of compounds Describe the structure of an atom Distinguish between ionic, hydrogen, and covalent bonds Define a chemical reaction and explain how it changes the composition of matter List and define the life-supporting properties of water Explain the pH scale and the formation of acid and base solutions Copyright © 2009 Pearson Education, Inc.

BIO 1010 Chapter 2 - Chemistry

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