1. The Chemical Context of Life
Lec # 2
Instructor: Dr. Hatem Eideh
Course: General Biology
Reference:
Reece JB, Urry LA, Cain ML, Wasserman SA, Minorsky PV and Jackson RB (2011). Campell Biology, 9th ed. Boston: Pearson Education. . Chap 2

2. Introduction: Who Tends This Garden?
“Devil’s gardens”

3. Living organisms are composed of about 25 chemical elements
Living organisms are composed of matter, which is anything that occupies space and has mass (weight)
Matter is composed of chemical elements
Element—a substance that cannot be broken down to other substances by chemical reactions
There are 92 elements in nature—only a few exist in a pure state
An element: the simplest forms of matter that can exist under conditions normally encountered in a chemistry lab. Elements cannot be decomposed into simpler substances by chemical reactions.
Molecule: an electrically neutral group of atoms bound tightly enough together that they behave as and can be recognized as a single particle.
Life results from the ordering of atoms into molecules and the interactions of these molecules within cells.
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
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 a species. A zoo might have trouble keeping a particular animal because zoologists have not identified all of the trace elements required in the animal’s diet.
Copyright © 2009 Pearson Education, Inc.

4. Elements can combine to form compounds
Compound—a substance consisting of two or more different elements combined in a fixed ratio by mass (and by atoms)
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
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Teaching Tips
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).
Copyright © 2009 Pearson Education, Inc.

5. + Sodium Chlorine Sodium Chloride
Figure 2.3 The emergent properties of the edible compound sodium chloride.
Sodium
Chlorine
Sodium Chloride

6. Life requires 25 essential elements; some are called trace elements (

7. CONNECTION: Trace elements are common additives to food and water
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
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
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 a species. A zoo might have trouble keeping a particular animal because zoologists have not identified all of the trace elements required in the animal’s diet.
2. Many breakfast cereals are fortified with iron. As noted in Module 2.2, you can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imaging 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.
Copyright © 2009 Pearson Education, Inc.

8. CONNECTION: Trace elements are common additives to food and water
Some trace elements are required to prevent disease
Without iron, your body cannot transport oxygen
An iodine deficiency prevents production of thyroid hormones, resulting in goiter
Although nitrogen is not a trace element, a common deficiency in plants, especially those we grow for food, is nitrogen deficiency. Without fertilization in a field that is used year after year to produce food crops, the plants are less productive.
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
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 a species. A zoo might have trouble keeping a particular animal because zoologists have not identified all of the trace elements required in the animal’s diet.
2. Many breakfast cereals are fortified with iron. As noted in Module 2.2, you can crush the cereal and extract distinct iron particles with a magnet. An overhead projector or video imaging 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.
Copyright © 2009 Pearson Education, Inc.

9. Elements can combine to form compounds
Many of the compounds in living organisms contain carbon, hydrogen, oxygen, and nitrogen
DNA, for example, contains all four of these elements
Interestingly, different arrangements of elements provide unique properties for each compound
Students will soon learn about organic compounds, especially those that are based on the properties of carbon.
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Teaching Tips
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).
Copyright © 2009 Pearson Education, Inc.

10. Atoms consist of protons, neutrons, and electrons
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
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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 eight days means that it will decay quickly.
Teaching Tips
1. Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton were as massive as a bowling ball, an electron would be the mass of a Lifesaver. (This is calculated by considering a 15-pound bowling ball, a Lifesaver with a mass of 0.12 ounces, and the mention in Module 2.4 that an electron is about 1/2,000 the mass of a proton.)
2. The text in Module 2.4 makes an analogy regarding the size of a helium atom. The text notes that 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. Such concrete examples help to relate abstract concepts.
3. After sharing teaching tips 1 and 2 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?
4. The text notes the use of radioactive isotopes in dating fossils but references Module 15.5 for further discussion. If your course does not include Chapter 15, consider explaining this process at this point in your course.
Copyright © 2009 Pearson Education, Inc.

11. Atoms consist of protons, neutrons, and electrons
Elements differ in their number of protons, neutrons, and electrons
Helium has two protons, two neutrons, and two electrons
Carbon has six protons, six neutrons, and six electrons
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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 eight days means that it will decay quickly.
Teaching Tips
1. Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton were as massive as a bowling ball, an electron would be the mass of a Lifesaver. (This is calculated by considering a 15-pound bowling ball, a Lifesaver with a mass of 0.12 ounces, and the mention in Module 2.4 that an electron is about 1/2,000 the mass of a proton.)
2. The text in Module 2.4 makes an analogy regarding the size of a helium atom. The text notes that 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. Such concrete examples help to relate abstract concepts.
3. After sharing teaching tips 1 and 2 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?
4. The text notes the use of radioactive isotopes in dating fossils but references Module 15.5 for further discussion. If your course does not include Chapter 15, consider explaining this process at this point in your course.
Copyright © 2009 Pearson Education, Inc.

12. Two models of a helium atom
Electron
cloud
Protons
2e–
Nucleus
Electrons
Mass
number = 4
Neutrons 2
Figure 2.4A Two models of a helium atom

13. Atoms consist of protons, neutrons, and electrons
Neutrons and protons are packed in the atom’s nucleus
The negative charge of electrons and the positive charge of protons keep electrons near the nucleus
The number of protons is the atom’s atomic number
Carbon with 6 protons has an atomic number of 6
The mass number is the sum of the protons and neutrons in the nucleus (carbon-12 is written 12C)
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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 eight days means that it will decay quickly.
Teaching Tips
1. Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton were as massive as a bowling ball, an electron would be the mass of a Lifesaver. (This is calculated by considering a 15-pound bowling ball, a Lifesaver with a mass of 0.12 ounces, and the mention in Module 2.4 that an electron is about 1/2,000 the mass of a proton.)
2. The text in Module 2.4 makes an analogy regarding the size of a helium atom. The text notes that 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. Such concrete examples help to relate abstract concepts.
3. After sharing teaching tips 1 and 2 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?
4. The text notes the use of radioactive isotopes in dating fossils but references Module 15.5 for further discussion. If your course does not include Chapter 15, consider explaining this process at this point in your course.
Copyright © 2009 Pearson Education, Inc.

14. Atoms consist of protons, neutrons, and electrons
Although all atoms of an element have the same atomic number, some differ in mass number
The variations are isotopes, which have the same numbers of protons and electrons but different numbers of neutrons
One isotope of carbon has 8 neutrons instead of 6 (written 14C)
Unlike 12C, 14C is an unstable (radioactive) isotope that gives off energy
Radioactive isotopes have many important applications in biology and medicine?
The isotope 12C is stable—its nucleus does not have a tendency to lose particles. The isotope 14C is radioactive—its nucleus decays spontaneously giving off particles and energy.
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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 eight days means that it will decay quickly.
Teaching Tips
1. Here is a comparison that helps make the point about the differences in mass of protons and electrons. If a proton were as massive as a bowling ball, an electron would be the mass of a Lifesaver. (This is calculated by considering a 15-pound bowling ball, a Lifesaver with a mass of 0.12 ounces, and the mention in Module 2.4 that an electron is about 1/2,000 the mass of a proton.)
2. The text in Module 2.4 makes an analogy regarding the size of a helium atom. The text notes that 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. Such concrete examples help to relate abstract concepts.
3. After sharing teaching tips 1 and 2 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?
4. The text notes the use of radioactive isotopes in dating fossils but references Module 15.5 for further discussion. If your course does not include Chapter 15, consider explaining this process at this point in your course.
Copyright © 2009 Pearson Education, Inc.

15. Table 2.4 Isotopes of Carbon.

16. CONNECTION: Radioactive isotopes can help or harm us
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
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Copyright © 2009 Pearson Education, Inc.

17. CONNECTION: Radioactive isotopes can help or harm us
Biologists use radioactive tracers in research
Radioactive 14C was used to show the route of 14CO2 in formation of sugar during plant photosynthesis
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Copyright © 2009 Pearson Education, Inc.

18. CONNECTION: Radioactive isotopes can help or harm us
Radioactive tracers are frequently used in medical diagnosis
Sophisticated imaging instruments are used to detect them
An imaging instrument that uses positron- emission tomography (PET) detects the location of injected radioactive materials
PET is useful for diagnosing heart disorders and cancer and in brain research
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Copyright © 2009 Pearson Education, Inc.

19. Chemical bonds
A covalent bond results when atoms share outer-shell electrons
A molecule is formed when atoms are held together by covalent bonds
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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 will benefit from a discussion of Figure 2.8 and the differences and limitations of representations of atomic structure. The contrast in Figure 2.8 is a good beginning for such a discussion. In addition to comparing how the positions of electrons are depicted, note the problems with the sense of scale as discussed in Module 2.4.
Teaching Tips
1. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (ionic, covalent, 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.)
2. Have your students try to calculate the number of covalent bonds possible for a variety of atoms. (Carbon, for example, can form up to four covalent bonds.) 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. (For example, carbon could form covalent bonds with four hydrogen atoms.)
Copyright © 2009 Pearson Education, Inc.

20. Table 2.8 Alternative Ways to Represent Four Common Molecules.

21. Unequal electron sharing creates polar molecules
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
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Teaching Tips
1. Modules 2.9 and 2.10 discuss the special bonding in and between water molecules. Many students do not appreciate the importance of weak chemical bonds in water and cellular chemistry. Extra time and attention may be required to address this special aspect of chemistry.
Copyright © 2009 Pearson Education, Inc.

22. Unequal electron sharing creates polar molecules
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
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Teaching Tips
1. Modules 2.9 and 2.10 discuss the special bonding in and between water molecules. Many students do not appreciate the importance of weak chemical bonds in water and cellular chemistry. Extra time and attention may be required to address this special aspect of chemistry.
Copyright © 2009 Pearson Education, Inc.

23. Unequal electron sharing creates polar molecules
Water has atoms with different electronegativities
Oxygen attracts the shared electrons more strongly than hydrogen
So, the shared electrons spend more time near oxygen
The result is a polar covalent bond
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Teaching Tips
1. Modules 2.9 and 2.10 discuss the special bonding in and between water molecules. Many students do not appreciate the importance of weak chemical bonds in water and cellular chemistry. Extra time and attention may be required to address this special aspect of chemistry.
Copyright © 2009 Pearson Education, Inc.

24. Unequal electron sharing creates polar molecules
In H2O the oxygen atom has a slight negative charge and the hydrogens have a slight positive charge
Molecules with this unequal distribution of charges are called polar molecules
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Teaching Tips
1. Modules 2.9 and 2.10 discuss the special bonding in and between water molecules. Many students do not appreciate the importance of weak chemical bonds in water and cellular chemistry. Extra time and attention may be required to address this special aspect of chemistry.
Copyright © 2009 Pearson Education, Inc.

25. (–)
(–) O H H
Figure 2.9 A water molecule.
(+)
(+)

26. Hydrogen bonds are weak bonds important in the chemistry of life
Some chemical bonds are weaker than covalent bonds
Hydrogen, as part of a polar covalent bond, will share attractions with other electronegative atoms
Examples are oxygen and nitrogen
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
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Teaching Tips
1. Modules 2.9 and 2.10 discuss the special bonding in and between water molecules. Many students do not appreciate the importance of weak chemical bonds in water and cellular chemistry. Extra time and attention may be required to address this special aspect of chemistry.
Copyright © 2009 Pearson Education, Inc.

27. Hydrogen bond
Figure 2.10 Hydrogen bonds between water molecules.

28. Ionic bonds are attractions between ions of opposite charge
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
Student Misconceptions
1. The dangers posed by certain chemicals in our food and broader environment often misled 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.
Teaching Tips
1. Consider challenging your students to suggest relationships in human lives that are analogous to each of the three types of chemical bonds (ionic, covalent, 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.)
Copyright © 2009 Pearson Education, Inc.

29. Sodium chloride (NaCl)
Transfer of
electron + – Na
Sodium atom Cl
Chlorine atom
Na+
Sodium ion
Cl–
Chloride ion
Figure 2.7A Formation of an ionic bond, producing sodium chloride.
Sodium chloride (NaCl)

30. Crystal of sodium chloride
Na+
Cl–
Figure 2.7B A crystal of sodium chloride.

31. Van der Waals interactions
Are weak interactions that occur only when atoms and molecules are very close together
These bonds are individually weak, but their cumulative is significant
Example: van de Waals is responsible for the ability of a gecko lizard to walk up a wall !!

32. Chemical reactions make and break chemical bonds
Reactants & products
Equilibrium