Campbells Biology Review Ch. 1-14 Flashcard

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CHAPTER 1 INTRODUCTION: THEMES IN THE STUDY OF LIFE OUTLINE

I. Life’s Hierarchical Order

  • A. The living world is a hierarchy, with each level of biological structure building on the level below it
  • B. Each level of biological structure has emergent properties
  • C. Cells are an organism’s basic units of structure and function
  • D. The continuity of life is based on heritable information in the form of DNA
  • E. Structure and function are correlated at all levels of biological organization
  • F. Organisms are open systems that interact continuously with their environments
  • G. Regulatory mechanisms ensure a dynamic balance in living systems Evolution, Unity, and Diversity

A. Diversity and unity are the dual faces of life on Earth

B. Evolution is the core theme of biology Science as a Process

A. Testable hypotheses are the hallmarks of the scientific process

B. Science and technology are functions of society

C. Biology is a multidisciplinary adventure

II. III. OBJECTIVES

After reading this chapter and attending lecture, the student should be able to:

  1. Briefly describe unifying themes that pervade the science of biology.
  2. Diagram the hierarchy of structural levels in biology.
  3. Explain how the properties of life emerge from complex organization.
  4. Describe seven emergent properties associated with life.
  5. Distinguish between holism and reductionism.
  6. Explain how technological breakthroughs contributed to the formulation of the cell theory and our current knowledge of the cell.
  7. Distinguish between prokaryotic and eukaryotic cells.
  8. Explain, in their own words, what is meant by “form fits function. “
  9. List the five kingdoms of life and distinguish among them.
  10. Briefly describe how Charles Darwin’s ideas contributed to the conceptual framework of biology.
  11. Outline the scientific method.
  12. Distinguish between inductive and deductive reasoning.
  13. Explain how science and technology are interdependent.

Chapter 1

Introduction: Themes in the Study of Life KEY TERMS emergent property population community ecosystem biome biogenesis holism reductionism prokaryotic eukaryotic taxonomy evolution natural selection scientific method hypothesis inductive reasoning control group variable experimental group deductive reasoning scientific theory

LECTURE NOTES

Biology, the study of life, is a human endeavor resulting from an innate attraction to life in its diverse forms (E. O. Wilson’s biophilia). The science of biology is enormous in scope.

• It reaches across size scales from submicroscopic molecules to the global distribution of biological communities.
• It encompasses life over huge spans of time from contemporary organisms to ancestral life forms stretching back nearly four billion years. As a science, biology is an ongoing process.
• As a result of new research methods developed over the past few decades, there has been an information explosion. Technological advances yield new information that may change the conceptual framework accepted by the majority of biologists. With rapid information flow and new discoveries, biology is in a continuous state of flux.

There are, however, enduring unifying themes that pervade the science of biology:
• A hierarchy of organization
• The cellular basis of life
• Heritable information
• The correlation between structure and function
• The interaction of organisms with their environment
• Unity in diversity
• Evolution: the core theme
• Scientific process: the hypothetico-deductive method I. Life’s Hierarchical Order A. The living world is a hierarchy, with each level of biological structure building on the level below it A characteristic of life is a high degree of order. Biological organization is based on a hierarchy of structural levels, with each level building on the levels below it. Chapter 1 Introduction: Themes in the Study of Life 3 Atoms Complex biological molecules Subcellular organelles are ordered into Cells In multicellular organisms similar cells are organised into Tissues Organs Organ systems Complex organism There are levels of organization beyond the individual organism: Population Community Ecosystem = = = Localized group of organisms belonging to the same species Populations of species living in the same area An energy-processing system of community interactions that include abiotic environmental factors such as soil and water Large scale communities classified by predominant vegetation type and distinctive combinations of plants and animals The sum of all the planet’s ecosystems Biomes = Biosphere = B. Each level of biological organization has emergent properties Emergent property = Property that emerges as a result of interactions between components.
• With each step upward in the biological hierarchy, new properties emerge that were not present at the simpler organizational levels.
• Life is difficult to define because it is associated with numerous emergent properties that reflect a hierarchy of structural organization. Some of the emergent properties and processes associated with life are the following: 1. Order. Organisms are highly ordered, and other characteristics of life emerge from this complex organization.

Chapter

1 Introduction: Themes in the Study of Life

2. Reproduction. Organisms reproduce; life comes only from life (biogenesis).

3. Growth and Development. Heritable programs stored in DNA direct the species-specific pattern of growth and development.

4. Energy Utilization. Organisms take in and transform energy to do work, including the maintenance of their ordered state.

5. Response to Environment. Organisms respond to stimuli from their environment.

6. Homeostasis. Organisms regulate their internal environment to maintain a steady-state, even in the face of a fluctuating external environment.

7. Evolutionary Adaptation. Life evolves in response to interactions between organisms and their environment. Because properties of life emerge from complex organization, it is impossible to fully explain a higher level of order by breaking it into its parts. Holism = The principle that a higher level of order cannot be meaningfully explained by examining component parts in isolation.
• An organism is a living whole greater than the sum of its parts.
• For example, a cell dismantled to its chemical ingredients is no longer a cell. It is also difficult to analyze a complex process without taking it apart. Reductionism = The principle that a complex system can be understood by studying its component parts.
• Has been a powerful strategy in biology
• Example: Watson and Crick deduced the role of DNA in inheritance by studying its molecular structure. The study of biology balances the reductionist strategy with the goal of understanding how the parts of cells, organisms, and populations are functionally integrated. C . Cells are an organism’s basic units of structure and function The cell is an organism’s basic unit of structure and function.
• Lowest level of structure capable of performing all activities of life. All organisms are composed of cells.
• May exist singly as unicellular organisms or as subunits of multicellular organisms. The invention of the microscope led to the discovery of the cell and the formulation of the cell theory.
• Robert Hooke (1665) reported a description of his microscopic examination of cork. Hooke described tiny boxes which he called “cells” (really cell walls). The significance of this discovery was not recognized until 150 years later.
• Antonie van Leeuwenhok (1600’s) used the microscope to observe living organisms such as microorganisms in pond water, blood cells, and animal sperm cells. Matthias Schleiden and Theodor Schwann (1839) reasoned from their own microscopic studies and those of others, that all living things are made of cells. This formed the basis for the cell theory.
• The cell theory has since been modified to include the idea that all cells come from preexisting cells. Over the past 40 years, use of the electron microscope has revealed the complex ultrastructure of cells.
• Cells are bounded by plasma membranes that regulate passage of materials between the cell and its surroundings.
• All cells, at some stage, contain DNA. Chapter 1 Introduction: Themes in the Study of Life 5 Based on structural organization, there are two major kinds of cells: prokaryotic and eukaryotic. Prokaryotic cell = Cell lacking membrane-bound organelles and a membrane-enclosed nucleus.
• Found only in the archaebacteria and bacteria
• Generally much smaller than eukaryotic cells
• Contains DNA that is not separated from the rest of the cell, as there is no membrane-bound nucleus
• Lacks membrane-bound organelles
• Almost all have tough external walls Eukaryotic cell = Cell with a membrane-enclosed nucleus and membrane-enclosed organelles. Found in protists, plants, fungi, and animals
• Subdivided by internal membranes into different functional compartments called organelles
• Contains DNA that is segregated from the rest of the cell. DNA is organized with proteins into chromosomes that are located within the nucleus, the largest organelle of most cells.
• Cytoplasm surrounds the nucleus and contains various organelles of different functions
• Some cells have a tough cell wall outside the plasma membrane (e. g. , plant cells). Animal cells lack cell walls. Though structurally different, eukaryotic and prokaryotic cells have many similarities, especially in their chemical processes. D. The continuity of life is based on heritable information in the form of DNA Biological instructions for an organism’s complex structure and function are encoded in DNA.
• Each DNA molecule is made of four types of chemical building blocks called nucleotides.
• The linear sequence of these four nucleotides encode the precise information in a gene, the unit of inheritance from parent to offspring.
• An organism’s complex structural organization is specified by an enormous amount of coded information. Inheritance is based on:
• A complex mechanism for copying DNA.
• Passing the information encoded in DNA from parent to offspring. All forms of life use essentially the same genetic code.
• A particular nucleotide sequence provides the same information to one organism as it does to another.
• Differences among organisms reflect differences in nucleotide sequence. E. Structure and function are correlated at all levels of biological organization There is a relationship between an organism’s structure and how it works. Form fits function.
• Biological structure gives clues about what it does and how it works. Knowing a structure’s function gives insights about its construction.
• This correlation is apparent at many levels of biological organization. 6 Chapter 1 Introduction: Themes in the Study of Life F. Organisms are open systems that interact continuously with their environments Organisms interact with their environment, which includes other organisms as well as abiotic factors.
• Both organism and environment are affected by the interaction between them.
• Ecosystem dynamics include two major processes: 1. Nutrient cycling 2. Energy flow (see Campbell, Figure 1. 7) G. Regulatory mechanisms ensure a dynamic balance in living systems Regulation of biological processes is critical for maintaining the ordered state of life. Many biological processes are self-regulating; that is, the product of a process regulates that process (= feedback regulation; see Campbell, Figure 1. 8).
• Positive feedback speeds a process up
• Negative feedback slows a process down Organisms and cells also use chemical mediators to help regulate processes.
• The hormone insulin, for example, signals cells in vertebrate organisms to take up glucose. As a result, blood glucose levels go down. In certain forms of diabetes mellitus, insulin is deficient and cells do not take up glucose as they should, and as a result, blood glucose levels remain high. II. Evolution, Unity, and Diversity A. Diversity and unity are the dual faces of life on Earth Biological diversity is enormous.
• Estimates of total diversity range from five million to over 30 million species.
• About 1. 5 million species have been identified and named, including approximately 260,000 plants, 50,000 vertebrates, and 750,000 insects. To make this diversity more comprehensible, biologists classify species into categories. Taxonomy = Branch of biology concerned with naming and classifying organisms.
• Taxonomic groups are ranked into a hierarchy from the most to least inclusive category: domain, kingdom, phylum, class, order, family, genus, species.
• A six-kingdom system recognizes two prokaryotic groups and divides the Monera into the Archaebacteria and Eubacteria.
• The kingdoms of life recognized in the traditional five-kingdom system are Monera, Protista, Plantae, Fungi, and Animalia (see Campbell, Figure 1. 10). There is unity in the diversity of life forms at the lower levels of organization. Unity of life forms is evident in:
• A universal genetic code.
• Similar metabolic pathways (e. g. , glycolysis).
• Similarities of cell structure (e. g. , flagella of protozoans and mammalian sperm cells). B. Evolution is the core theme of biology Evolution is the one unifying biological theme.
• Life evolves. Species change over time and their history can be described as a branching tree of life.
• Species that are very similar share a common ancestor at a recent branch point on the phylogenetic tree.
• Less closely related organisms share a more ancient common ancestor. Chapter 1 Introduction: Themes in the Study of Life 7 All life is connected and can be traced back to primeval prokaryotes that existed more than three billion years ago. In 1859, Charles Darwin published On the Origin of Species in which he made two major points: 1. Species change, and contemporary species arose from a succession of ancestors through a process of “descent with modification. ” 2. A mechanism of evolutionary change is natural selection. Darwin synthesized the concept of natural selection based upon the following observations:
• Individuals in a population of any species vary in many inheritable traits. Populations have the potential to produce more offspring than will survive or than the environment can support.
• Individuals with traits best suited to the environment leave a larger number of offspring, which increases the proportion of inheritable variations in the next generation. This differential reproductive success is what Darwin called natural selection. Organisms’ adaptations to their environments are the products of natural selection.

• Natural selection does not create adaptations; it merely increases the frequency of inherited variants that arise by chance. Adaptations are the result of the editing process of natural selection. When exposed to specific environmental pressures, certain inheritable variations favor the reproductive success of some individuals over others. Darwin proposed that cumulative changes in a population over long time spans could produce a new species from an ancestral one. Descent with modification accounts for both the unity and diversity of life.
• Similarities between two species may be a reflection of their descent from a common ancestor. Differences between species may be the result of natural selection modifying the ancestral equipment in different environmental contexts. III. Science as a Process A. Testable hypotheses are the hallmarks of the scientific process As the science of life, biology has the characteristics associated with science in general. Science is a way of knowing. It is a human endeavor that emerges from our curiosity about ourselves, the world, and the universe. Good scientists are people who:
• Ask questions about nature and believe those questions are answerable. Are curious, observant, and passionate in their quest for discovery.
• Are creative, imaginative, and intuitive.
• Are generally skeptics. Scientific method = Process which outlines a series of steps used to answer questions.
• Is not a rigid procedure.
• Based on the conviction that natural phenomena have natural causes.
• Requires evidence to logically solve problems. The key ingredient of the scientific process is the hypothetico-deductive method, which is an approach to problem-solving that involves: 1. Asking a question and formulating a tentative answer or ypothesis by inductive reasoning. 2. Using deductive reasoning to make predictions from the hypothesis and then testing the validity of those predictions.

Chapter 1 Introduction: Themes in the Study of Life Hypothesis = Educated guess proposed as a tentative answer to a specific question or problem. Inductive reasoning = Making an inference from a set of specific observations to reach a general conclusion. Deductive reasoning = Making an inference from general premises to specific consequences, which logically follow if the premises are true. Usually takes the form of If… then logic.
• In science, deductive reasoning usually involves predicting experimental results that are expected if the hypothesis is true. Some students cannot make the distinction between inductive and deductive reasoning. An effective teaching strategy is to let them actually experience both processes. To illustrate inductive reasoning, provide an every day scenario with enough pieces of information for student to hypothesize a plausible explanation for some event. Demonstrate deductive reasoning by asking students to solve a simple problem, based upon given assumptions. Useful hypotheses have the following characteristics:
• Hypotheses are possible causes. Generalizations formed by induction are not necessarily hypotheses. Hypotheses should also be tentative explanations for observations or solutions to problems.
• Hypotheses reflect past experience with similar questions. Hypotheses are not just blind propositions, but are educated guesses based upon available evidence.
• Multiple hypotheses should be proposed whenever possible. The disadvantage of operating under only one hypothesis is that it might restrict the search for evidence in support of this hypothesis; scientists might bias their search, as well as neglect to consider other possible solutions.
• Hypotheses must be testable via the hypothetico-deductive method. Predictions made from hypotheses must be testable by making observations or performing experiments. This limits the scope of questions that science can answer.
• Hypotheses can be eliminated, but not confirmed with absolute certainty. I f repeated experiments consistently disprove the predictions, then we can assume that the hypothesis is false. However, if repeated experimentation supports the deductions, we can only assume that the hypothesis may be true; accurate predictions can be made from false hypotheses. The more deductions that are tested and supported, the more confident we can be that the hypothesis is true. Another feature of the scientific process is the controlled experiment which includes control and experimental groups. Control group = In a controlled experiment, the group in which all variables are held constant.
• Controls are a necessary basis for comparison with the experimental group, which has been exposed to a single treatment variable. Allows conclusions to be made about the effect of experimental manipulation.
• Setting up the best controls is a key element of good experimental design. Variable = Condition of an experiment that is subject to change and that may influence an experiment’s outcome. Experimental group = In a controlled experiment, the group in which one factor or treatment is varied. Science is an ongoing process that is a self-correcting way of knowing. Scientists:
• Build on prior scientific knowledge.
• Try to replicate the observations and experiments of others to check on their conclusions. Chapter 1 Introduction: Themes in the Study of Life 9 Share information through publications, seminars, meetings, and personal communication. What really advances science is not just an accumulation of facts, but a new concept that collectively explains observations that previously seemed to be unrelated.
• Newton, Darwin, and Einstein stand out in the history of science because they synthesized ideas with great explanatory power.
• Scientific theories are comprehensive conceptual frameworks which are well supported by evidence and are widely accepted by the scientific community. B. Science and technology are functions of society Science and technology are interdependent.
• Technology extends our ability to observe and measure, which enables scientists to work on new questions that were previously unapproachable.
• Science, in turn, generates new information that makes technological inventions possible.
• Example: Watson and Crick’s scientific discovery of DNA structure led t o further investigation that enhanced our understanding of DNA, the genetic code, and how to transplant foreign genes into microorganisms. The biotechnology industry has capitalized on this knowledge to produce valuable pharmaceutical products such as human insulin. We have a love-hate relationship with technology.
• Technology has improved our standard of living.
• The consequence of using technology also includes the creation of new problems such as increased population growth, acid rain, deforestation, global warming, nuclear accidents, ozone holes, toxic wastes, and endangered species.
• Solutions to these problems have as much to do with politics, economics, culture and values as with science and technology. A better understanding of nature must remain the goal of science. Scientists should:
• Try to influence how technology is used.
• Help educate the public about the benefits and hazards of specific technologies. C . Biology is a multidisciplinary adventure Biology is a multidisciplinary science that integrates concepts from chemistry, physics and mathematics. Biology also embraces aspects of humanities and the social sciences.
• 10 Chapter 1 Introduction: Themes in the Study of Life REFERENCES Campbell, N. Biology. 5th ed. Menlo Park, California: Benjamin/Cummings, 1998. Moore, J. A. “Science as a Way of Knowing–Evolutionary Biology. ” American Zoologist, 24(2): 470-475, 1980.

CHAPTER 2 THE CHEMICAL CONTEXT OF LIFE OUTLINE

I. Chemical Elements and Compounds A. Matter consists of chemical elements in pure form and in combinations called compounds B. Life requires about 25 chemical elements Atoms and Molecules A. Atomic structure determines the behavior of an element B. Atoms combine by chemical bonding to form molecules C. Weak chemical bonds play important roles in the chemistry of life D. A molecule’s biological function is related to its shape E. Chemical reactions make and break chemical bonds II. OBJECTIVES After reading this chapter and attending lecture, the student should be able to: 1. Define element and compound. 2. State four elements essential to life that make up 96% of living matter. 3. Describe the structure of an atom. 4. Define and distinguish among atomic number, mass number, atomic weight, and valence. 5. Given the atomic number and mass number of an atom, determine the number of neutrons. 6. Explain why radioisotopes are important to biologists. 7. Explain how electron configuration influences the chemical behavior of an atom. . Explain the octet rule and predict how many bonds an atom might form. 9. Explain why the noble gases are so unreactive. 10. Define electronegativity and explain how it influences the formation of chemical bonds. 11. Distinguish among nonpolar covalent, polar covalent and ionic bonds. 12. Describe the formation of a hydrogen bond and explain how it differs from a covalent or ionic bond. 13. Explain why weak bonds are important to living organisms. 14. Describe how the relative concentrations of reactants and products affect a chemical reaction. 12 Unit I The Chemistry of Life KEY TERMS atter element trace element atom neutron proton electron atomic nucleus dalton atomic number mass number atomic weight isotope radioactive isotope energy potential energy energy level energy potential energy energy level electron shell orbital valence electron valence shell chemical bond covalent bond molecule structural formula molecular formula double covalent bond valence electronegativity nonpolar covalent bond polar covalent bond ion cation anion ionic bond hydrogen bond chemical reactions reactants products chemical equilibrium LECTURE NOTES I. Chemical Elements and Compounds A. Matter consists of chemical elements in pure form and in combinations called compounds Chemistry is fundamental to an understanding of life, because living organisms are made of matter. Matter = Anything that takes up space and has mass. Mass = A measure of the amount of matter an object contains. You might want to distinguish between mass and weight for your students. Mass is the measure of the amount of matter an object contains, and it stays the same regardless of changes in the object’s position. Weight is the measure of how strongly an object is pulled by earth’s gravity, and it varies with distance from the earth’s center. The key point is that the mass of a body does not vary with its position, whereas weight does. So, for all practical purposes—as long as we are earthbound—weight can be used as a measure of mass. B. Life requires about 25 chemical elements Element = A substance that cannot be broken down into other substances by chemical reactions.
• All matter is made of elements.
• There are 92 naturally occurring elements.
• They are designated by a symbol of one or two letters. About 25 of the 92 naturally occurring elements are essential to life. Biologically important elements include: C O H N = = = = carbon oxygen hydrogen nitrogen ake up 96% of living matter Chapter 2 The Chemical Context of Life 13 Ca = calcium P = phosphorus K = potassium S = sulfur Na = sodium Cl = chlorine Mg = magnesium Trace elements make up remaining 4% of an organism’s weight Trace element = Element required by an organism in extremely minute quantities.
• Though required by organisms in small quantity, they are indispensable for life
• Examples: B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V and Zn Elements can exist in combinations called compounds.
• Compound = A pure substance composed of two or more elements combined in a fixed ratio. Example: NaCl (sodium chloride)
• Has unique emergent properties beyond those of its combined elements (Na and Cl have very different properties from NaCl). See Campbell, Figure 2. 2. Since a compound is the next structural level above element or atom, this is an excellent place to emphasize the concept of emergent properties, an integral theme found throughout the text and course. II. Atoms and Molecules A. Atomic structure determines the behavior of an element Atom = Smallest possible unit of matter that retains the physical and chemical properties of its element.
• Atoms of the same element share similar chemical properties. Atoms are made up of subatomic particles. 1. Subatomic particles The three most stable subatomic particles are: 1. Neutrons [no charge (neutral)]. 2. Protons [+1 electrostatic charge]. 3. Electrons [-1 electrostatic charge]. NEUTRON No charge PROTON +1 charge ELECTRON -1 charge Found together in a dense core called the nucleus Orbits around nucleus (held (positively charged because of protons) by electrostatic attraction to positively charged nucleus) 1. 009 dalton 1. 007 dalton 1/2000 dalton Mass is so small, usually not used to calculate atomic mass Masses of both are about the same (about 1 dalton) NOTE: The dalton is a unit used to express mass at the atomic level. One dalton (d) is equal to 1. 67 x 10-24 g. If an atom is electrically neutral, the number of protons equals the number of electrons, which yields an electrostatically balanced charge. 14 Unit I The Chemistry of Life 2. Atomic number and atomic weight Atomic number = Number of protons in an atom of a particular element.
• All atoms of an element have the same atomic number.
• Written as a subscript to the left of the element’s symbol (e. g. , 11 Na)
• In a neutral atom, # protons = # electrons. Mass number = Number of protons and neutrons in an atom.
• Written as a superscript to left of an element’s symbol (e. g. , 23 Na)
• Is approximate mass of the whole atom, since the mass of a proton and the mass of a neutron are both about 1 dalton
• Can deduce the number of neutrons by subtracting atomic number from mass number
• Number of neutrons can vary in an element, but number of protons is constant
• Is not the same as an element’s atomic weight, which is the weighted mean of the masses of an element’s constituent isotopes In a large classroom with up to 300 students, it can be difficult to interact. Try putting examples on an overhead transparency and soliciting student input t o complete the information. It is a quick way to check for understanding and t o actively involve students. Examples: (Mass #) 23 (Atomic #) 11 Na # of electrons # of protons # of neutrons 12 6 C # of electrons # of protons # of neutrons 3. Isotopes Isotopes = Atoms of an element that have the same atomic number but different mass number.
• They have the same number of protons, but a different number of neutrons.
• Under natural conditions, elements occur as mixtures of isotopes. Different isotopes of the same element react chemically in same way.
• Some isotopes are radioactive. Radioactive isotope = Unstable isotope in which the nucleus spontaneously decays, emitting subatomic particles and/or energy as radioactivity. Loss of nuclear particles may transform one element to another (e. g. , 146 C ; 147 N).
• Has a fixed half life.
• Half life = Time for 50% of radioactive atoms in a sample to decay. Biological applications of radioactive isotopes include: a. Dating geological strata and fossils
• Chapter 2 The Chemical Context of Life 5 Radioactive decay is at a fixed rate. By comparing the ratio of radioactive and stable isotopes in a fossil with the ratio of isotopes in living organisms, one can estimate the age of a fossil.
• The ratio of 14 C to 12 C is frequently used to date fossils less than 50,000 years old. b. Radioactive tracers
• Chemicals labelled with radioactive isotopes are used to trace the steps of a biochemical reaction or to determine the location of a particular substance within an organism (see Campbell, p. XX, Methods: The Use of Radioactive Tracers in Biology). Radioactive isotopes are useful as biochemical tracers because they chemically react like the stable isotopes and are easily detected at low concentrations.
• Isotopes of P, N, and H were used to determine DNA structure.
• Used to diagnose disease (e. g. , PET scanner)
• Because radioactivity can damage cell molecules, radioactive isotopes can also be hazardous c. Treatment of cancer
• e. g. , radioactive cobalt 4. The energy levels of electrons Electrons = Light negatively charged particles that orbit around nucleus. Equal in mass and charge
• Are the only stable subatomic particles directly involved in chemical reactions
• Have potential energy because of their position relative to the positively charged nucleus Energy = Ability to do work Potential energy = Energy that matter stores because of its position or location.
• There is a natural tendency for matter to move to the lowest state of potential energy.
• Potential energy of electrons is not infinitely divisible, but exists only in discrete amounts called quanta.
• Different fixed potential energy states for electrons are called energy levels or electron shells (see Campbell, Figure 2. 7). Electrons with lowest potential energy are in energy levels closest to the nucleus.
• Electrons with greater energy are in energy levels further from nucleus. Electrons may move from one energy level to another. In the process, they gain or lose energy equal to the difference in potential energy between the old and new energy level.

• 16 Unit I The Chemistry of Life 5. Electron orbitals Orbital = Three-dimensional space where an electron will most likely be found 90% of the time (see Campbell, Figure 2. 8).
• Viewed as a three-dimensional probability cloud (a statistical concept)
• No more than two electrons can occupy same orbital. First energy level:
• Has one spherical s orbital (1s orbital)
• Holds a maximum of two electrons Second energy level
• Holds a maximum of eight electrons
• One spherical s orbital (2s orbital)
• Three dumbbell-shaped p orbitals each oriented at right angles to the other two (2p x , 2p y , 2p z orbitals) Higher energy levels:
• Contain s and p orbitals
• Contain additional orbitals with more complex shapes 6. Electron configuration and chemical properties An atom’s electron configuration determines its chemical behavior. Electron configuration = Distribution of electrons in an atom’s electron shells The first 18 elements of a periodic chart are arranged sequentially by atomic number into three rows (periods). In reference to these representative elements, note the following:
• Outermost shell of these atoms never have more than four orbitals (one s and three p) or eight electrons.
• Electrons must first occupy lower electron shells before the higher shells can be occupied. (This is a reflection of the natural tendency for matter to move to the lowest possible state of potential energy—the most stable state.
• Electrons are added to each of the p orbitals singly, before they can be paired.
• If an atom does not have enough electrons to fill all shells, the outer shell will be the only one partially filled. Example: O2 with a total of eight electrons: Chapter 2 The Chemical Context of Life 17 OXYGEN 😯 Two electrons have the 1s orbital of the first electron shell. First two electrons in the second shell are both in the 2s orbital. Next three electrons each have a p orbital (2p x, 2p y, 2p z). Eighth electron is paired in the 2p x orbital. 1s 2 2s 2px 2py 2pz 2 2 1 1 Chemical properties of an atom depend upon the number of valence electrons. Valence electrons = Electrons in the outermost energy shell (valence shell). Octet rule = Rule that a valence shell is complete when it contains eight electrons (except H and He).
• An atom with a complete valence shell is unreactive or inert.
• Noble elements (e. g. , helium, argon, and neon) have filled outer shells in their elemental state and are thus inert.
• An atom with an incomplete valence shell is chemically reactive (tends to form chemical bonds until it has eight electrons to fill the valence shell).
• Atoms with the same number of valence electrons show similar chemical behavior. NOTE: The consequence of this unifying chemical principle is that the valence electrons are responsible for the atom’s bonding capacity. This rule applies to most of the representative elements, but not all. B. Atoms combine by chemical bonding to form molecules Atoms with incomplete valence shells tend to fill those shells by interacting with other atoms. These interactions of electrons among atoms may allow atoms to form chemical bonds.
• Chemical bonds = Attractions that hold molecules together Molecules = Two or more atoms held together by chemical bonds. 1. Covalent bonds Covalent bond = Chemical bond between atoms formed by sharing a pair of valence electrons.
• Strong chemical bond
• Example: molecular hydrogen (H2 ); when two hydrogen atoms come close H2 H H H-H 18 Unit I The Chemistry of Life enough for their 1s orbitals to overlap, they share electrons, thus completing the valence shell of each atom. Structural formula = Formula which represents the atoms and bonding within a molecule (e. g. , H-H). The line represents a shared pair of electrons. Molecular formula = Formula which indicates the number and type of atoms (e. . , H2 ). Single covalent bond = Bond between atoms formed by sharing a single pair of valence electrons.
• Atoms may freely rotate around the axis of the bond. Double covalent bond = Bond formed when atoms share two pairs of valence electrons (e. g. , O2 ). O2 O O O=O Molecules = Two or more atoms held together by chemical bonds. Triple covalent bond = Bond formed when atoms share three pairs of valence electrons (e. g. , N2 or N? N). NOTE: Double and triple covalent bonds are rigid and do not allow rotation. Valence = Bonding capacity of an atom which s the number of covalent bonds that must be formed to complete the outer electron shell.
• Valences of some common elements: hydrogen = 1, oxygen = 2, nitrogen = 3, carbon = 4, phosphorus = 3 (sometimes 5 as in biologically important compounds, e. g. , ATP), sulfur = 2. Compound = A pure substance composed of two or more elements combined in a fixed ratio.
• Example: water (H2 O), methane (CH4 )
• Note that two hydrogens are necessary to complete the valence shell of oxygen in water, and four hydrogens are necessary for carbon to complete the valence shell in methane. Chapter 2 The Chemical Context of Life 19 2. Nonpolar and polar covalent bonds Electronegativity = Atom’s ability to attract and hold electrons.
• The more electronegative an atom, the more strongly it attracts shared electrons.
• Scale determined by Linus Pauling: O = 3. 5 N = 3. 0 S and C = 2. 5 P and H = 2. 1 Nonpolar covalent bond = Covalent bond formed by an equal sharing of electrons between atoms.
• Occurs when electronegativity of both atoms is about the same (e. g. , CH4 )
• Molecules made of one element usually have nonpolar covalent bonds (e. . , H2 , O 2 , Cl2 , N 2 ). Polar covalent bond = Covalent bond formed by an unequal sharing of electrons between atoms.
• Occurs when the atoms involved have different electronegativities.
• Shared electrons spend more time around the more electronegative atom.
• In H2 O, for example, the oxygen is strongly electronegative, so negatively charged electrons spend more time around the oxygen than the hydrogens. This causes the oxygen atom to have a slight negative charge and the hydrogens to have a slight positive charge (see also Campbell, Figure 2. 1). 3. Ionic bonds Ion = Charged atom or molecule. Anion = An atom that has gained one or more electrons from another atom and has become negatively charged; a negatively charged ion. Cation = An atom that has lost one or more electrons and has become positively charged; a positively charged ion. Ionic bond = Bond formed by the electrostatic attraction after the complete transfer of an electron from a donor atom to an acceptor.
• The acceptor atom attracts the electrons because it is much more electronegative than the donor atom.
• Are strong onds in crystals, but are fragile bonds in water; salt crystals will readily dissolve in water and dissociate into ions.
• Ionic compounds are called salts (e. g. , NaCl or table salt) (see Campbell, Figure 2. 13). NOTE: The difference in electronegativity between interacting atoms determines if electrons are shared equally (nonpolar covalent), shared unequally (polar covalent), gained or lost (ionic bond). Nonpolar covalent bonds and ionic bonds are two extremes of a continuum from interacting atoms with similar electronegativities to interacting atoms with very different electronegativities. 20 Unit I The Chemistry of Life C . Weak chemical bonds play important roles in the chemistry of life Biologically important weak bonds include the following:
• Hydrogen bonds, ionic bonds in aqueous solutions, and other weak forces such as Van der Waals and hydrophobic interactions
• Make chemical signaling possible in living organisms because they are only temporary associations. Signal molecules can briefly and reversibly bind t o receptor molecules on a cell, causing a short-lived response.
• Can form between molecules or between different parts of a single large molecule.
• Help stabilize the three-dimensional shape of large molecules (e. . , DNA and proteins). 1. Hydrogen bonds Hydrogen bond = Bond formed by the charge attraction when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom.
• Weak attractive force that is about 20 times easier to break than a H O covalent bond
• Is a charge attraction between H oppositely charged portions of polar Electronegative Hydrogen atoms molecules bond
• Can occur between a hydrogen that N has a slight positive charge when H covalently bonded to an atom with H high electronegativity (usually O and H N)
• Example: NH3 in H2 O (see Campbell, Figure 2. 4) 2. Van der Waals interactions Weak interactions that occur between atoms and molecules that are very close together and result from charge asymetry in electron clouds. D. A molecule’s biological function is related to its shape A molecule has a charasteric size and shape. The function of many molecules depends upon their shape Insulin causes glucose uptake into liver and muscle cells of veterbrates because the shape of the insulin molecule is recognized by specific receptors on the target cell.
• Molecules with only two atoms are linear.
• Molecules with more than two atoms have more complex shapes. When an atom forms covalent bonds, orbitals in the valence shell rearrange into the most stable configuration. To illustrate, consider atoms with valence electrons in the s and three p orbitals:
• The s and three p orbitals hybridize into four new orbitals.
• The new orbitals are teardrop shaped, extend from the nucleus and spread out as far apart as possible.
• Example: If outer tips of orbitals in methane (CH4 ) are connected by imaginary lines, the new molecule has a tetrahedral shape with C at the center (see Campbell, Figure 2. 15). Chapter 2 The Chemical Context of Life 21 E. Chemical reactions make and break chemical bonds Chemical reactions = process of making and breaking chemical bonds leading to changes in the composition of matter.
• Process where reactants undergo changes into products.
• Matter is conserved, so all reactant atoms are only rearranged to form products.
• Some reactions go to completion (all reactants converted to products), but most reactions are reversible. For example: 3H 2 + N 2
• 2NH 3 The relative concentration of reactants and products affects reaction rate (the higher the concentration, the greater probability of reaction). Chemical equilibrium = Equilibrium established when the rate of forward reaction equals the rate of the reverse reaction.
• Is a dynamic equilibrium with reactions continuing in both directions
• Relative concentrations of reactants and products stop changing. Point out to students that chemical equilibrium does NOT mean that the concentrations of reactants and products are equal. REFERENCES Atkins, P. W. Atoms, Electrons and Change. W. H. Freeman and Company, 1991. Campbell, N. , et al. Biology. 5th ed. Menlo Park, California: Benjamin/Cummings, 1998. Weinberg, S. The Discovery of Subatomic Particles. New York, San Francisco: W. H. Freeman and Company, 1983. Brown, T. L. , H. E. Le May, Jr. , and B. Bursten. Chemistry: The Central Science. 7th ed. Upper Saddle River, New Jersey: Prentice Hall, 1997. CHAPTER 3 WATER AND THE FITNESS OF THE ENVIRONMENT OUTLINE I. Water’s Polarity and Its Effects A. The polarity of water molecules results in hydrogen bonding B. Organisms depend on the cohesion of water molecules C. Water moderates temperatures on Earth D. Oceans and lakes don’t freeze solid because ice floats E. Water is the solvent of life The Dissociation of Water A. Organisms are sensitive to changes in pH Acid Precipitation Threatens the Fitness of the Environment II. III. OBJECTIVES After reading this chapter and attending lecture, the student should be able to: 1. Describe how water contributes to the fitness of the environment to support life. 2. Describe the structure and geometry of a water molecule, and explain what properties emerge as a result of this structure. 3. Explain the relationship between the polar nature of water and its ability to form hydrogen bonds. 4. List five characteristics of water that are emergent properties resulting from hydrogen bonding. 5. Describe the biological significance of the cohesiveness of water. 6. Distinguish between heat and temperature. 7. Explain how water’s high specific heat, high heat of vaporization and expansion upon freezing affect both aquatic and terrestrial ecosystems. 8. Explain how the polarity of the water molecule makes it a versatile solvent. 9. Define molarity and list some advantages of measuring substances in moles. 10. Write the equation for the dissociation of water, and explain what is actually transferred from one molecule to another. 11. Explain the basis for the pH scale. 12. Explain how acids and bases directly or indirectly affect the hydrogen ion concentration of a solution. 3. Using the bicarbonate buffer system as an example, explain how buffers work. 24 Unit I The Chemistry of Life 14. Describe the causes of acid precipitation, and explain how it adversely affects the fitness of the environment. KEY TERMS polar molecule cohesion adhesion surface tension kinetic energy heat temperature acid precipitation Celsius scale calorie kilocalorie joule specific heat evaporative cooling solution solute solvent aqueous solution hydrophilic hydrophobic mole molecular weight hydrogen ion molarity hydroxide ion acid base pH scale buffer LECTURE NOTES Water
• contributes to the fitness of the environment to support life. Life on earth probably evolved in water. Living cells are 70%-95% H2 O. Water covers about 3/4 of the earth. In nature, water naturally exists in all three physical states of matter—solid, liquid and gas. Water’s extraordinary properties are emergent properties resulting from water’s structure and molecular interactions. I. Water’s Polarity and Its Effects A. The polarity of water molecules results in hydrogen bonding Water is a polar molecule. Its polar bonds and asymmetrical shape give water molecules opposite charges on opposite sides. Four valence orbitals of O point t o Unbonded electron pairs corners of a tetrahedron.
• Two corners are orbitals with unshared pairs of electrons and weak negative charge.
• Two corners are occupied by H atoms which are in polar covalent bonds with O. Oxygen is so electronegative, that shared electrons spend more time around the O causing a weak positive charge near H’s. Hydrogen bonding orders water into a higher level of structural organization.
• The polar molecules of water are held together by hydrogen bonds.
• Positively charged H of one molecule is attracted to the negatively charged O of another water molecule. H O
• Each water molecule can form a H maximum of four hydrogen bonds with neighboring water molecules. Chapter 3 Water and the Fitness of the Environment 25 Water has extraordinary properties that emerge as a consequence of its polarity and hydrogen-bonding. Some of these properties are that water:
• has cohesive behavior
• resists changes in temperature
• has a high heat of vaporization and cools surfaces as it evaporates
• expands when it freezes
• is a versatile solvent B. Organisms depend on the cohesion of water molecules. Cohesion = Phenomenon of a substance being held together by hydrogen bonds. Though hydrogen bonds are transient, enough water molecules are hydrogen bonded at any given time to give water more structure than other liquids.
• Contributes to upward water transport in plants by holding the water column together. Adhesion of water to vessel walls counteracts the downward pull of gravity. Surface tension = Measure of how difficult it is to stretch or break the surface of a liquid.
• Water has a greater surface tension than most liquids; function of the fact that at the air/H2 O interface, surface water molecules are hydrogen bonded to each other and to the water molecules below. Causes H2 O to bead (shape with smallest area to volume ratio and allows maximum hydrogen bonding). C . Water moderates temperatures on Earth 1. Heat and temperature Kinetic energy = The energy of motion. Heat = Total kinetic energy due to molecular motion in a body of matter. Temperature = Measure of heat intensity due to the average kinetic energy of molecules in a body of matter. Calorie (cal) = Amount of heat it takes to raise the temperature of one gram of water by one degree Celsius. Conversely, one calorie is the amount of heat that one gram of water releases when it cools down by one degree Celsius. NOTE: The “calories” on food packages are actually kilocalories (kcal). Kilocalorie (kcal or Cal) = Amount of heat required to raise the temperature of one kilogram of water by one degree Celsius (1000 cal). Celsius Scale at Sea Level 100°C (212°F) 37°C (98. 6°F) 23°C (72°F) 0°C (32°F) = water boils = human body temperature = room temperature = water freezes °K = °C °F Scale Conversion = = 5(°F- 32) 9 9° C+ 32 5 °C + 273 2. Water’s high specific heat Water has a high specific heat, which means that it resists temperature changes when it absorbs or releases heat. Specific heat = Amount of heat that must be absorbed or lost for one gram of a substance to change its temperature by one degree Celsius. Specific heat of water = One calorie per gram per degree Celsius (1 cal/g/°C). 26 Unit I The Chemistry of Life As a result of hydrogen bonding among water molecules, it takes a relatively large heat loss or gain for each 1°C change in temperature.
• Hydrogen bonds must absorb heat to break, and they release heat when they form.
• Much absorbed heat energy is used to disrupt hydrogen bonds before water molecules can move faster (increase temperature). A large body of water can act as a heat sink, absorbing heat from sunlight during the day and summer (while warming only a few degrees) and releasing heat during the night and winter as the water gradually cools. As a result:
• Water, which covers three-fourths of the planet, keeps temperature fluctuations within a range suitable for life.
• Coastal areas have milder climates than inland.
• The marine environment has a relatively stable temperature. 3. Evaporative cooling Vaporization (evaporation) = transformation from liquid to a gas. Molecules with enough kinetic energy to overcome the mutual attraction of molecules in a liquid, can escape into the air. Heat of vaporization = Quantity of heat a liquid must absorb for 1 g to be converted to the gaseous state.
• For water molecules to evaporate, hydrogen bonds must be broken which requires heat energy.
• Water has a relatively high heat of vaporization at the boiling point (540 cal/g or 2260 J/g; Joule = 0. 239 cal). Evaporative cooling = Cooling of a liquid’s surface when a liquid evaporates (see Campbell, Figure 3. 4). The surface molecules with the highest kinetic energy are most likely to escape into gaseous form; the average kinetic energy of the remaining surface molecules is thus lower. Water’s high heat of vaporization:
• Moderates the Earth’s climate.
• Solar heat absorbed by tropical seas dissipates when surface water evaporates (evaporative cooling).
• As moist tropical air moves poleward, water vapor releases heat as it condenses into rain.
• Stabilizes temperature in aquatic ecosystems (evaporative cooling).
• Helps organisms from overheating by evaporative cooling. D. Oceans and lakes don’t freeze solid because ice floats Because of hydrogen bonding, water is less dense as a solid than it is as a liquid. Consequently, ice floats.
• Water is densest at 4°C.
• Water contracts as it cools to 4°C.
• As water cools from 4°C to freezing (0°C), it expands and becomes less dense than liquid water (ice floats).
• When water begins to freeze, the molecules do not have enough kinetic energy to break hydrogen bonds.
• As the crystalline lattice forms, each water molecule forms a maximum of four hydrogen bonds, which keeps water molecules further apart than they would be in the liquid state; see Campbell, Figure 3. .
• Chapter 3 Water and the Fitness of the Environment 27 Expansion of water contributes to the fitness of the environment for life:
• Prevents deep bodies of water from freezing solid from the bottom up.
• Since ice is less dense, it forms on the surface first. As water freezes it releases heat to the water below and insulates it.
• Makes the transitions between seasons less abrupt. As water freezes, hydrogen bonds form releasing heat. As ice melts, hydrogen bonds break absorbing heat. E. Water is the solvent of life Solution = A liquid that is a completely homogenous mixture f two or more substances. Solvent = Dissolving agent of a solution. Solute = Substance dissolved in a solution. Aqueous solution = Solution in which water is the solvent. Water is a versatile solvent owing to the polarity of the water molecule. Hydrophilic { Hydrophobic { Ionic compounds dissolve in water (see Campbell, Figure 3. 8).
• Charged regions of polar water molecules have an electrical attraction to charged ions.
• Water surrounds individual ions, separating and shielding them from one another. Polar compounds in general, are water-soluble. Charged regions of polar water molecules have an affinity for oppositely charged regions of other polar molecules. Nonpolar compounds (which have symmetric distribution in charge) are NOT water-soluble. 1. Hydrophilic and hydrophobic substances Ionic and polar substances are hydrophilic, but nonpolar compounds are hydrophobic. Hydrophilic = (Hydro = water; philo = loving); property of having an affinity for water.
• Some large hydrophilic molecules can absorb water without dissolving. Hydrophobic = (Hydro = water; phobos = fearing); property of not having an affinity for water, and thus, not being water-soluble. . Solute concentration in aqueous solutions Most biochemical reactions involve solutes dissolved in water. There are two important quantitative properties of aqueous solutions: solute concentration and pH. Molecular weight = Sum of the weight of all atoms in a molecule (expressed in daltons). Mole = Amount of a substance that has a mass in grams numerically equivalent to its molecular weight in daltons. 28 Unit I The Chemistry of Life For example, to determine a mole of sucrose (C12 H22 O11 ):
• Calculate molecular weight: C = 12 dal 12 dal ? 12 = 144 dal H = 1 dal 1 dal ? 2 = 22 dal O = 16 dal 16 dal ? 11 = 176 dal 342 dal
• Express it in grams (342 g). Molarity = Number of moles of solute per liter of solution
• To make a 1M sucrose solution, weigh out 342 g of sucrose and add water up to 1L. Advantage of measuring in moles:
• Rescales weighing of single molecules in daltons to grams, which is more practical for laboratory use.
• A mole of one substance has the same number of molecules as a mole of any other substance (6. 02 ? 1023 ; Avogadro’s number).
• Allows one to combine substances in fixed ratios of molecules. II. The Dissociation of Water Occasionally, the hydrogen atom that is shared in a hydrogen bond between two water molecules, shifts from the oxygen atom to which it is covalently bonded to the unshared orbitals of the oxygen atom to which it is hydrogen bonded.
• Only a hydrogen ion (proton with a +1 charge) is actually transferred.
• Transferred proton binds to an unshared orbital of the second water molecule creating a hydronium ion (H 3 O+).
• Water molecule that lost a proton has a net negative charge and is called a hydroxide ion (OH -). H2 O + H 2 O H3 O+ + OH
• By convention, ionization of H2 O is expressed as the dissociation into H+ and OH -. H2 O H+ + OH
• Reaction is reversible.
• At equilibrium, most of the H2 O is not ionized. A. Organisms are sensitive to changes in pH 1. Acids and bases At equilibrium in pure water at 25°C:
• Number of H+ ions = number of OH- ions. 1
• [H +] = [OH-] = M = 10-7 M 10,000,000
• Note that brackets indicate molar concentration. This is a good place to point out how few water molecules are actually dissociated (only 1 out of 554,000,000 molecules). Chapter 3 Water and the Fitness of the Environment 9 ACID Substance that increases the relative [H +] of a solution. Also removes OH- because it tends to combine with H+ to form H2 O. For example: (in water) HCl H+ + Cl- BASE Substance that reduces the relative [H+] of a solution. May alternately increase [OH-]. For example: A base may reduce [H+] directly: NH3 + H + NH 4 + A base may reduce [H+] indirectly: NaOH ; Na + + OH OH- + H + ; H2 O A solution in which:
• [H +] = [OH-] is a neutral solution.
• [H +] ; [OH-] is an acidic solution.
• [H +] ; [OH-] is a basic solution. Strong acids and bases dissociate completely in water. Example: HCl and NaOH
• Single arrows indicate complete dissociation. NaOH ; Na + + OH Weak acids and bases dissociate only partially and reversibly.
• Examples: NH3 (ammonia) and H2 CO3 (carbonic acid)
• Double arrows indicate a reversible reaction; at equilibrium there will be a fixed ratio of reactants and products. H2 CO3 HCO3 H+ Carbonic Bicarbonate + Hydrogen acid ion ion 2. The pH scale In any aqueous solution: [H +][OH -] = 1. 0 ? 10-14 For example:
• In a neutral solution, [H+] = 10-7 M and [OH-] = 10-7 M.
• In an acidic solution where the [H+] = 10-5 M, the [OH-] = 10-9 M. In a basic solution where the [H+] = 10-9 M, the [OH-] = 10-5 M. pH scale = Scale used to measure degree of acidity. It ranges from 0 to 14. pH = Negative log10 of the [H+] expressed in moles per liter.
• pH of 7 is a neutral solution.
• pH ; 7 is an acidic solution.
• pH ; 7 is a basic solution. 30 Unit I The Chemistry of Life

• Most biological fluids are within the pH range of 6 to 8. There are some exceptions such as stomach acid with pH = 1. 5. (See Campbell, Figure 3. 9) Each pH unit represents a tenfold difference (scale is logarithmic), so a slight change in pH represents a large change in actual [H+]. To illustrate this point, project the following questions on a transparency and cover the answer. The students will frequently give the wrong response (3? ), and they are surprised when you unveil the solution. How much greater is the [H+] in a solution with pH 2 than in a solution with pH 6? 1 ANS: pH 2 = [H+] of 1. 0 ? 10-2 = M 100 1 pH 6 = [H+] of 1. 0 ? 10-6 = M 1,000,000 10,000 times greater. 3. Buffers By minimizing wide fluctuations in pH, buffers help organisms maintain the pH of body fluids within the narrow range necessary for life (usually pH 6-8). Buffer = Substance that minimizes large sudden changes in pH. Are combinations of H+-donor and H+-acceptor forms in a solution of weak acids or bases
• Work by accepting H+ ions from solution when they are in excess and by donating H+ ions to the solution when they have been depleted Example: Bicarbonate buffer response to a rise in pH H2 CO3 H donor (weak acid) + HCO3 response to a drop in pH H acceptor (weak base) + + H+ Hydrogen ion HCl + NaHCO3 strong acid H2 CO3 + NaCl weak acid NaOH + H 2 CO3 strong base NaHCO3 + H 2 O weak base III. Acid Precipitation Threatens the Fitness of the Environment Acid precipitation = Rain, snow, or fog more strongly acidic than pH 5. .
• Has been recorded as low as pH 1. 5 in West Virginia
• Occurs when sulfur oxides and nitrogen oxides in the atmosphere react with water in the air to form acids which fall to Earth in precipitation
• Major oxide source is the combustion of fossil fuels by industry and cars
• Acid rain affects the fitness of the environment to support life.
• Lowers soil pH which affects mineral solubility. May leach out necessary mineral nutrients and increase the concentration of minerals that are potentially toxic to vegetation in higher concentration (e. g. , aluminum). This is contributing to the decline of some European and North American forests. Chapter 3 Water and the Fitness of the Environment 31 Lowers the pH of lakes and ponds, and runoff carries leached out soil minerals into aquatic ecosystems. This adversely affects aquatic life. Example: In the Western Adirondack Mountains, there are lakes with a pH ; 5 that have no fish. What can be done to reduce the problem?
• Add industrial pollution controls.
• Develop and use antipollution devices.
• Increase involvement of voters, consumers, politicians, and business leaders. The political issues surrounding acid rain can be used to enhance student awareness and make this entire topic more relevant and interesting to the students.
• REFERENCES Campbell, N. , et al. Biology. 5th ed. Menlo Park, California: Benjamin/Cummings, 1998. Gould, R. Going Sour: Science and Politics of Acid Rain. Boston: Birkhauser, 1985. Henderson, L. J. The Fitness of the Environment. Boston: Beacon Press, 1958. Mohnen, V. A. “The Challenge of Acid Rain. ” Scientific American, August 1988. CHAPTER 4 CARBON AND MOLECULAR DIVERSITY OUTLINE I. The Importance of Carbon A. Organic chemistry is the study of carbon compounds B. Carbon atoms are the most versatile building blocks of molecules C. Variation in carbon skeletons contributes to the diversity of organic molecules Functional Groups A. Functional groups also contribute to the molecular diversity of life II. OBJECTIVES After reading this chapter and attending lecture, the student should be able to: 1. Summarize the philosophies of vitalism and mechanism, and explain how they influenced the development of organic chemistry, as well as mainstream biological thought. 2. Explain how carbon’s electron configuration determines the kinds and number of bonds carbon will form. 3. Describe how carbon skeletons may vary, and explain how this variation contributes to the diversity and complexity of organic molecules. 4. Distinguish among the three types of isomers: structural, geometric and enantiomers. 5. Recognize the major functional groups, and describe the chemical properties of organic molecules in which they occur. KEY TERMS organic chemistry hydrocarbon isomer structural isomer geometric isomer enantiomer functional group hydroxyl group alcohol carbonyl group aldehyde ketone carboxyl group carboxylic acid amino group amine sulfhydryl group thiol phosphate group LECTURE NOTES Aside from water, most biologically important molecules are carbon-based (organic). The structural and functional diversity of organic molecules emerges from the ability of carbon to form large, complex and diverse molecules by bonding to itself and to other elements such as H, O, N, S, and P. 34 I. Unit I The Chemistry of Life The Importance of Carbon A. Organic chemistry is the study of carbon compounds Organic chemistry = The branch of chemistry that specializes in the study of carbon compounds. Organic molecules = Molecules that contain carbon Vitalism = Belief in a life force outside the jurisdiction of chemical/physical laws. Early 19th century organic chemistry was built on a foundation of vitalism because organic chemists could not artificially synthesize organic compounds. It was believed that only living organisms could produce organic compounds. Mechanism = Belief that all natural phenomena are governed by physical and chemical laws.
• Pioneers of organic chemistry began to synthesize organic compounds from inorganic molecules. This helped shift mainstream biological thought from vitalism to mechanism.
• For example, Friedrich Wohler synthesized urea in 1828; Hermann Kolbe synthesized acetic acid. Stanley Miller (1953) demonstrated the possibility that organic compounds could have been produced under the chemical conditions of primordial Earth. B. Carbon atoms are the most versatile building blocks of molecules The carbon atom:
• Usually has an atomic number of 6; therefore, it has 4 valence electrons.
• Usually completes its outer energy shell by sharing valence electrons in four covalent bonds. (Not likely to form ionic bonds. ) Emergent properties, such as the kinds and number of bonds carbon will form, are determined by their tetravalent electron configuration. It makes large, complex molecules possible. The carbon atom is a central point from which the molecule branches off into four directions.
• It gives carbon covalent compatibility with many different elements. The four major atomic components of organic molecules are as follows:
• It determines an organic molecule’s three-dimensional shape, which may affect molecular function. For example, when carbon forms four single covalent bonds, the four valence orbitals hybridize into teardrop-shaped orbitals that angle from the carbon atoms toward the corners of an imaginary tetrahedron. Students have problems visualizing shapes of organic molecules in three dimensions. Specific exa

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