Category: 1st Semester

Chapter 21 – Genetic Basis of Development

 

Chapter 21    Genetic Basis of Development
Objectives
1. Explain how advances in recombinant DNA technology have helped scientists study the eukaryotic genome.

2. Describe the natural function of restriction enzymes.

3. Describe how restriction enzymes and gel electrophoresis are used to isolate DNA fragments.

4. Explain how the creation of sticky ends by restriction enzymes is useful in producing a recombinant DNA molecule.

5. Outline the procedures for producing plasmid and phage vectors.

6. Explain how vectors are used in recombinant DNA technology.

7. List and describe the two major sources of genes for cloning.

8. Describe the function of reverse transcriptase in retroviruses and explain how they are useful in recombinant DNA technology.

9. Describe how “genes of interest” can be identified with the use of a probe.

10. Explain the importance of DNA synthesis and sequencing to modern studies of eukaryotic genomes.

11. Describe how bacteria can be induced to produce eukaryotic gene products.

12. List some advantages for using yeast in the production of gene products.

13. List and describe four complementary approaches used to map the human genome.

14. Explain how RFLP analysis and PCR can be applied to the Human Genome Project.

15. Describe how recombinant DNA technology can have medical applications such as diagnosis of genetic disease, development of gene therapy, vaccine production, and development of pharmaceutical products.

16. Describe how gene manipulation has practical applications for agriculture.

17. Describe how plant genes can be manipulated using the Ti plasmid carried by Agrobacterium as a vector.

18. Explain how foreign DNA may be transferred into monocotyledonous plants.

19. Describe how recombinant DNA studies and the biotechnology industry are regulated with regards to safety and policy matters.

 
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Chapter 22 AP Objectives

 

Chapter 22    Darwinian View of Life
Objectives
The Historical Context for Evolutionary Theory
1. Explain the mechanism for evolutionary change proposed by Charles Darwin in On the Origin of Species.
2. Define evolution and adaptation.
3. Compare and contrast Aristotle’s scala naturae to Carolus Linnaeus’ classification scheme.
4. Describe the theories of catastrophism, gradualism, and uniformitarianism.
5. Explain the mechanism for evolutionary change proposed by Jean-Baptiste de Lamarck. Explain why modern biology has rejected Lamarck’s theories.
The Darwinian Revolution
6. Describe how Darwin’s observations on the voyage of the HMS Beagle led him to formulate and support his theory of evolution.
7. Explain how the principle of gradualism and Charles Lyell’s theory of uniformitarianism influenced Darwin’s ideas about evolution.
8. Explain what Darwin meant by “descent with modification.”
9. Explain what evidence convinced Darwin that species change over time.
10. Explain how Linnaeus’ classification scheme fit Darwin’s theory of evolution by natural selection.
11. Describe the three inferences Darwin made from his observations that led him to propose natural selection as a mechanism for evolutionary change.
12. Explain how an essay by the Rev. Thomas Malthus influenced Charles Darwin.
13. Distinguish between artificial selection and natural selection.
14. Explain why an individual organism cannot evolve.
15. Describe the experiments that supported Reznick and Endler’s hypothesis that differences in life-history traits between guppy populations are due to selective pressure based on predation.
16. Explain how the existence of homologous and vestigial structures can be explained by Darwin’s theory of natural selection.
17. Explain how evidence from biogeography supports the theory of evolution by natural selection.
18. Explain the problem with the statement that Darwinism is “just a theory.” Distinguish between the scientific and colloquial use of the word theory.

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Chapter 23 AP Objectives

 

Chapter 23    Evolution of Populations
Objectives
Population Genetics
1. Explain the statement “It is the population, not the individual, that evolves.”
2. Explain how Mendel’s particulate hypothesis of inheritance provided much-needed support for Darwin’s theory of evolution by natural selection.
3. Distinguish between discrete and quantitative traits. Explain how Mendel’s laws of inheritance apply to quantitative traits.
4. Explain what is meant by “the modern synthesis.”
5. Define the terms population, species, and gene pool.
6. Explain why meiosis and random fertilization alone will not alter the frequency of alleles or genotypes in a population.
7. List the five conditions that must be met for a population to remain in Hardy-Weinberg equilibrium.
8. Write the Hardy-Weinberg equation. Use the equation to calculate allele frequencies when the frequency of homozygous recessive individuals in a population is 25%.
Mutation and Sexual Recombination
9. Explain why the majority of point mutations are harmless.
10. Explain why mutation has little quantitative effect on allele frequencies in a large population.
11. Describe the significance of transposons in the generation of genetic variability.
12. Explain how sexual recombination generates genetic variability.
Natural Selection, Genetic Drift, and Gene Flow
13. Explain the following statement: “Only natural selection leads to the adaptation of organisms to their environment.”
14. Explain the role of population size in genetic drift.
15. Distinguish between the bottleneck effect and the founder effect.
16. Describe how gene flow can act to reduce genetic differences between adjacent populations.
Genetic Variation, the Substrate for Natural Selection
17. Explain how quantitative and discrete characters contribute to variation within a population.
18. Distinguish between average heterozygosity and nucleotide variability. Explain why average heterozygosity tends to be greater than nucleotide variability.
19. Define a cline.
20. Define relative fitness.
a. Explain why relative fitness is zero for a healthy, long-lived, sterile organism.
b. Explain why relative fitness could be high for a short-lived organism.
21. Distinguish among directional, disruptive, and stabilizing selection. Give an example of each mode of selection.
22. Explain how diploidy can protect a rare recessive allele from elimination by natural selection.
23. Describe how heterozygote advantage and frequency-dependent selection promote balanced polymorphism.
24. Define neutral variations. Explain why natural selection does not act on these alleles.
25. Distinguish between intrasexual selection and intersexual selection.
26. Explain how female preferences for showy male traits may benefit the female.
27. Describe the disadvantages of sexual reproduction.
28. Explain how the genetic variation promoted by sex may be advantageous to individuals on a generational time scale.
29. List four reasons why natural selection cannot produce perfect organisms.

 
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Chemistry of Organisms

Chemistry
All Materials © Cmassengale

Composition of Matter

Ø  Everything in the universe is made of matter

Ø  Matter takes up space & has mass

Ø  Mass is a measure of the amount of matter in the substance

Ø  Mass & weight are NOT the same

Ø  Weight is a measure of the pull of gravity on an object

Question: Is the mass of an object the same on the moon as it is on the Earth? Is its weight the same? (Hint: Gravitational pull on the moon is 1/6 of that on the Earth.)

Ø  Matter exists in 4 states – solid, liquid, gas, & plasma

Ø  Solids have both a definite volume & definite shape (rock)

Ø  Liquids have a definite volume but no definite shape; they can be    poured (water)

Ø  Gases do not have a definite volume or definite shape, but they take the  volume & shape of their container

Ø  Plasmas have no definite volume, no definite shape, and only exist at extremely high temperatures such as the sun

Ø  Chemical Changes in matter are essential to all life processes

Ø  Biologists study chemistry because all living things are made of the same kinds of matter that make up nonliving things

Elements

Ø     Elements are pure substances which cannot be chemically broken down into simpler kinds of matter

Ø     More than 100 elements have been identified, but only about 30 are important in living things

Ø     All of the Elements are arranged on a chart known as the Periodic Table

Ø     Periodic charts tell the atomic number, atomic mass, & chemical symbol for every element

Ø     Four elements, Carbon – C, Hydrogen – H, Oxygen – O, and Nitrogen – N make up almost 90% of the mass of living things

Ø     Every element has a different chemical symbol composed of one to two letters

Ø     Chemical symbols usually come from the first letter or letters of an element like C for Carbon and Cl for Chlorine

Ø     Some chemical symbols come form their Latin or Greek name such as  Na for Sodium (natrium) or K for Potassium (Kalium)

Ø      Elements in the same horizontal period on the periodic table have the same number of energy levels (e.g. H & He in period 1 have only a K energy level)

[Periodic Table]
All Period 2 elements have 2 energy levels
(K & L)

Ø      Elements in the same vertical Family on the periodic table have the same number of electrons in their outermost energy level & react similar (e.g. Family IV, the Carbon family all have 4 electrons in their outermost energy level)

Atoms

Ø     Atoms are the simplest part of an element that keeps all of the element’s properties

Ø     Atoms are too small to be seen so scientists have developed models that show their structure & properties

Ø     Atoms consist of 3 kinds of subatomic particles – protons & neutrons in the center or nucleus, and electrons spinning in energy levels around the center

Ø     The nucleus is the center of an atom where most of the mass is concentrated

Ø     Protons are positively charged ( p+ ),  have a mass of 1 amu (atomic mass unit) , are found in the nucleus, and determine the atomic number of the element

Example:  Carbon has 6 protons so its atomic number is 6

Ø     Neutrons are neutral or have no electrical charge (n), have a mass of 1 amu, are found in the nucleus, and when added to the number of protons, determine the atomic mass of the element

Example:  Sodium has 11 protons and 12 neutrons so its atomic mass is 11+12=23 amu

Ø     Electrons (e-) are negatively charged, high energy particles with little mass that spin around the nucleus in energy levels

Ø     Seven energy levels (K, L, M, N, O, P, & Q) exist around the nucleus and each holds a certain number of electrons

Ø     The K energy level is closest to the nucleus & only holds 2 electrons, while the  L – Q energy levels can hold 8 electrons  

Ø     Electrons in outer energy level are traveling faster & contain more energy than electrons in inner levels  

Ø     The number of protons (positive charges) and electrons (negative charges in an atom are equal so the net electrical charge on a atom is zero making it electrically neutral

Ø     Stable or non-reactive atoms have an outer energy level that is filled with electrons  

Compounds

Ø     Most elements do not exist by themselves; Most elements combine with other elements

Ø      Compounds are made of atoms of two or more elements chemically combined

Ø      Chemical Formulas represent a compound & show the kind & number of atoms of each element  (e.g. H2O has 2 hydrogen & 1 oxygen)

Ø      Compounds have different physical & chemical properties than the atoms that compose them  (e.g. hydrogen & oxygen are gases but H2O is a liquid)

Ø      The number & arrangement of electrons in an atom determines if it will combine to form compounds

Ø      Chemical reactions occur whenever unstable atoms (outer energy level not filled) combine to form more stable compounds

Ø      Chemical bonds form between atoms during chemical reactions

Types of Chemical Bonds

Ø     Covalent bonds form between atoms whenever they share 1 or more pairs of electrons (e.g. H2O)  

Ø     Molecules form from covalent bonding & are the simplest part of a compound (e.g. NaCl, H2O, O2)  

Ø     Ionic bonding occurs between a positively & negatively charged atom or ion  

Ø     Positively charged ions have more electrons (-) than protons (+); negatively charged ions have more protons than electrons

Ø     Table salt (NaCl) forms when the 1 outer electron of Na is transferred to the outer energy level of chlorine that has 7 electrons (e-)

Ø     Sodium (Na) with 1 less e- becomes positively charged, while Chlorine (Cl) with 1 more e- becomes negatively charged; the + and – charges attract & form the ionic bond holding NaCl together

Ø     Other types of chemical bonding include hydrogen bonding

Energy

Ø     Energy is the ability to do work

Ø     Energy occurs in several forms & may be converted from one form to another

Ø     Sunlight is the ultimate energy for all life on earth

Ø     Forms of energy include chemical, electrical, mechanical, thermal, light, & sound

Ø     Free energy is the energy available for work (e.g. cells have energy to carry out cell processes)

Ø     Cells convert the chemical energy stored in food into other types of energy such as thermal & mechanical

Ø     Energy is used to change matter form one state into another (e.g. liquid into a gas)

Chemical Reactions

Ø     Living things undergo thousands of chemical reactions

Ø     Chemical equations represent chemical reactions

Ø     CO2 + H20—–goes to—–H2CO3  (carbonic acid) is a sample Chemical Reaction in living things

Ø     Reactants are on the left side of the equation, while products are on the right side

Ø Activation energy is required to start many reactions

Ø     Chemical bonds are broken, atoms rearranged, and new bonds form in chemical reaction

Ø     Plants use sunlight to produce sugars such as C6H12O6 glucose; the chemical energy from the sun is stored in the chemical bonds of glucose

Ø      Organisms eat plants, break down the sugars, and release energy along with CO2 & H2O

Ø      Exergonic reactions involve a net release of energy; while endergonic reactions involve a net absorption of energy

Ø      Energy must be added to the reactants for most chemical reactions to occur; called activation energy

Ø      Enzymes are chemical substances in living things that act as catalysts & reduce the amount of activation energy needed

Ø      Organisms contain thousands of different enzymes

Ø      Most enzymes end with –ase (e.g. lipase is the enzyme that acts on lipids)

Reduction-Oxidation (Redox) reactions

Ø     Reactions in which e- are transferred between atoms is a redox or reduction-oxidation reaction (e.g. formation of table salt NaCl)

Ø     In oxidation reactions, a reactant loses 1 or more e- & becomes positively (+) charged (e.g. Sodium atom becomes a Na+ ion)

Ø     In a reduction reaction, a reactant gains 1 or more e- and becomes negatively (-) charged (e.g. Chlorine atom becomes a Cl- ion)

Ø     REDOX reactions always occur together; the electron(s) from the oxidation reaction are then accepted by another substance in the reduction reaction

Solutions

Ø     A large percentage of the mass of organisms is water & many of the chemical reactions of life occur in water

Ø     A solution  is a uniform mixture of one substance in anther

Ø     Solutions may be mixtures of solids, liquids, or gases

Ø     The solute is the substance uniformly dissolved in the solution & may be ions, molecules, or atoms

Ø     The solvent is the substance in which the solute is dissolved

Ø     Water is known as the universal solvent 

Ø     Dissolving one substance in another does not alter their chemical properties

Ø     The concentration of a solution is a measure of the amount of solute dissolved in a given volume of solvent

Ø     Increasing the amount of solute increases the solution’s concentration

Ø     Aqueous solutions are solutions in which water is the solvent; these are important in living things (e.g. blood, cytoplasm of cell…)

Acids and Bases

Ø     The degree of acidity or alkalinity (basic) is important in organisms

Ø     The force of attraction between molecules is so strong that the oxygen atom of one molecule can actually remove the hydrogen from other water molecules; called Dissociation

Ø      H20—–GOES TO—– H+  +  OH-

Ø     OH- called hydroxide ion; H+ called hydrogen ion

Ø     Free H+ ion can react with another water molecule to form H3O+  (hydronium ion)

Ø     Acidity or alkalinity is a measure of the relative amount of H+ and OH- ions dissolved in a solution

Ø     Neutral solutions have an equal number of H+ and OH- ions

Ø     Acids have more H3O+ ions than OH- ions; taste sour; and can be corrosive

Ø     Bases contain more OH- ions than H3O+ ions; taste bitter; & feel slippery  

 

Examples of Common Acids

  • citric acid (from certain fruits and veggies, notably citrus fruits)
  • ascorbic acid (vitamin C, as from certain fruits)
  • vinegar (5% acetic acid)
  • carbonic acid (for carbonation of soft drinks)
  • lactic acid (in buttermilk)
 Examples of Common Bases

  • detergents
  • soap
  • lye (NaOH)
  • household ammonia

PH Scale

Ø     Compares the relative concentration of H3O+ ions and OH- ions

Ø     Scale ranges from 0 to 14; 0-3 is very acidic; 7 is neutral; 11-14 is very basic or alkaline

 

Ø    Litmus paper, phenolphthalein, pH paper, & other indicators that change color can be used to measure pH

Buffers

Ø     Control of pH is important to organisms

Ø     Enzymes function only within a narrow pH range; usually neutral

Ø     Buffers neutral acids or bases in organisms to help control pH

Chemistry Study Guide Chemistry On-line

 

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Chapter 24 AP Objectives

 

Chapter 24    Origin of Species
Objectives
What Is a Species?
1. Distinguish between anagenesis and cladogenesis.
2. Define Ernst Mayr’s biological species concept.
3. Distinguish between prezygotic and postzygotic isolating mechanisms.
4. Describe five prezygotic isolating mechanisms and give an example of each.
5. Explain a possible cause for reduced hybrid viability.
6. Explain how hybrid breakdown maintains separate species even if fertilization occurs.
7. Describe some limitations of the biological species concept.
8. Define and distinguish among the following: ecological species concept, paleontological species concept, phylogenetic species concept, and morphological species concept.
Modes of Speciation
9. Distinguish between allopatric and sympatric speciation.
10. Explain the allopatric speciation model and describe the mechanisms that may lead to divergence of isolated gene pools.
11. Describe examples of adaptive radiation in the Galápagos and Hawaiian archipelagoes.
12. Explain how reproductive barriers evolve. Describe an example of the evolution of a prezygotic barrier and the evolution of a postzygotic barrier.
13. Define sympatric speciation and explain how polyploidy can cause reproductive isolation.
14. Distinguish between an autopolyploid and an allopolyploid species and describe examples of each.
15. Describe how cichlid fishes may have speciated in sympatry in Lake Victoria.
Adaptive Radiation
16. Define adaptive radiation and describe the circumstances under which adaptive radiation may occur.
17. Describe the two gene loci implicated in speciation in Mimulus.
From Speciation to Macroevolution
18. Explain in general terms how a complex structure can evolve by natural selection.
19. Define exaptation and illustrate this concept with an example.
20. Explain how slight genetic divergences may lead to major morphological differences between species.
21. Explain how the evolution of changes in temporal and spatial developmental dynamics can result in evolutionary novelties.
22. Define evo-devo, heterochrony, allometric growth, and paedomorphosis.
23. Explain why extracting a single evolutionary progression from a fossil record can be misleading.
24. Define and illustrate the concept of species selection.
25. Explain why evolutionary change is not goal-directed.
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