Category: 1st Semester

Chapter 3 – Water and the Fitness of the Environment – Lecture Outline

Chapter 3    Water and the Fitness of the Environment    Lecture Outline

Overview: The Molecule That Supports All of Life

  • Because water is the substance that makes life possible on Earth, astronomers hope to find evidence of water on newly discovered planets orbiting distant stars.
  • Life on Earth began in water and evolved there for 3 billion years before colonizing the land.
  • Even terrestrial organisms are tied to water.
  • Most cells are surrounded by water.
  • Cells are about 70–95% water.
  • Water is a reactant in many of the chemical reactions of life.
  • Water is the only common substance that exists in the natural world in all three physical states of matter: solid ice, liquid water, and water vapor.

Concept 3.1 The polarity of water molecules results in hydrogen bonding

  • In a water molecule, two hydrogen atoms form single polar covalent bonds with an oxygen atom.
  • Because oxygen is more electronegative than hydrogen, the region around the oxygen atom has a partial negative charge.
  • The regions near the two hydrogen atoms have a partial positive charge.
  • A water molecule is a polar molecule in which opposite ends of the molecule have opposite charges.
  • Water has a variety of unusual properties because of the attraction between polar water molecules.
  • The slightly negative regions of one water molecule are attracted to the slightly positive regions of nearby water molecules, forming hydrogen bonds.
  • Each water molecule can form hydrogen bonds with up to four neighbors.

Concept 3.2 Four emergent properties of water contribute to Earth’s fitness for life

  • Organisms depend on the cohesion of water molecules.
  • The hydrogen bonds joining water molecules are weak, about 1/20 as strong as covalent bonds.
  • They form, break, and reform with great frequency. Each hydrogen bond lasts only a few trillionths of a second.
  • At any instant, a substantial percentage of all water molecules are bonded to their neighbors, creating a high level of structure.
  • Collectively, hydrogen bonds hold water together, a phenomenon called cohesion.
  • Cohesion among water molecules plays a key role in the transport of water and dissolved nutrients against gravity in plants.
  • Water molecules move from the roots to the leaves of a plant through water-conducting vessels.
  • As water molecules evaporate from a leaf, other water molecules from vessels in the leaf replace them.
  • Hydrogen bonds cause water molecules leaving the vessels to tug on molecules farther down.
  • This upward pull is transmitted down to the roots.
  • Adhesion, clinging of one substance to another, contributes too, as water adheres to the wall of the vessels.
  • Surface tension, a measure of the force necessary to stretch or break the surface of a liquid, is related to cohesion.
  • Water has a greater surface tension than most other liquids because hydrogen bonds among surface water molecules resist stretching or breaking the surface.
  • Water behaves as if covered by an invisible film.
  • Some animals can stand, walk, or run on water without breaking the surface.
  • Water moderates temperatures on Earth.
  • Water stabilizes air temperatures by absorbing heat from warmer air and releasing heat to cooler air.
  • Water can absorb or release relatively large amounts of heat with only a slight change in its own temperature.
  • Atoms and molecules have kinetic energy, the energy of motion, because they are always moving.
  • The faster a molecule moves, the more kinetic energy it has.
  • Heat is a measure of the total quantity of kinetic energy due to molecular motion in a body of matter.
  • Temperature measures the intensity of heat in a body of matter due to the average kinetic energy of molecules.
  • As the average speed of molecules increases, a thermometer will record an increase in temperature.
  • Heat and temperature are related, but not identical.
  • When two objects of different temperatures come together, heat passes from the warmer object to the cooler object until the two are the same temperature.
  • Molecules in the cooler object speed up at the expense of kinetic energy of the warmer object.
  • Ice cubes cool a glass of pop by absorbing heat from the pop as the ice melts.
  • In most biological settings, temperature is measured on the Celsius scale (°C).
  • At sea level, water freezes at 0°C and boils at 100°C.
  • Human body temperature is typically 37°C.
  • While there are several ways to measure heat energy, one convenient unit is the calorie (cal).
  • One calorie is the amount of heat energy necessary to raise the temperature of one g of water by 1°C.
  • A calorie is released when 1 g of water cools by 1°C.
  • In many biological processes, the kilocalorie (kcal) is more convenient.
  •  A kilocalorie is the amount of heat energy necessary to raise the temperature of 1000 g of water by 1°C.
  • Another common energy unit, the joule (J), is equivalent to 0.239 cal.
  • Water stabilizes temperature because it has a high specific heat.
  • The specific heat of a substance is the amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1°C.
  • By definition, the specific heat of water is 1 cal per gram per degree Celsius or 1 cal/g/°C.
  • Water has a high specific heat compared to other substances.
  • For example, ethyl alcohol has a specific heat of 0.6 cal/g/°C.
  • The specific heat of iron is 1/10 that of water.
  • Water resists changes in temperature because of its high specific heat.
  • In other words, water absorbs or releases a relatively large quantity of heat for each degree of temperature change.
  • Water’s high specific heat is due to hydrogen bonding.
  • Heat must be absorbed to break hydrogen bonds, and heat is released when hydrogen bonds form.
  • Investment of one calorie of heat causes relatively little change to the temperature of water because much of the energy is used to disrupt hydrogen bonds, not speed up the movement of water molecules.
  • Water’s high specific heat has effects that range from the level of the whole Earth to the level of individual organisms.
  • A large body of water can absorb a large amount of heat from the sun in daytime during the summer and yet warm only a few degrees.
  • At night and during the winter, the warm water will warm cooler air.
  • Therefore, ocean temperatures and coastal land areas have more stable temperatures than inland areas.
  • Living things are made primarily of water. Consequently, they resist changes in temperature better than they would if composed of a liquid with a lower specific heat.
  • The transformation of a molecule from a liquid to a gas is called vaporization or evaporation.
  • This occurs when the molecule moves fast enough to overcome the attraction of other molecules in the liquid.
  • Even in a low-temperature liquid (with low average kinetic energy), some molecules are moving fast enough to evaporate.
  • Heating a liquid increases the average kinetic energy and increases the rate of evaporation.
  • Heat of vaporization is the quantity of heat that a liquid must absorb for 1 g of it to be converted from liquid to gas.
  • Water has a relatively high heat of vaporization, requiring about 580 cal of heat to evaporate 1 g of water at room temperature.
  • This is double the heat required to vaporize the same quantity of alcohol or ammonia.
  • This is because hydrogen bonds must be broken before a water molecule can evaporate from the liquid.
  • Water’s high heat of vaporization moderates climate.
  • Much of the sun’s heat absorbed by tropical oceans is used for evaporation of surface water.
  • As moist tropical air moves to the poles, water vapor condenses to form rain, releasing heat.
  • As a liquid evaporates, the surface of the liquid that remains behind cools, a phenomenon called evaporative cooling.
  • This occurs because the most energetic molecules are the most likely to evaporate, leaving the lower–kinetic energy molecules behind.
  • Evaporative cooling moderates temperature in lakes and ponds.
  • Evaporation of sweat in mammals or evaporation of water from the leaves of plants prevents terrestrial organisms from overheating.
  • Evaporation of water from the leaves of plants or the skin of humans removes excess heat.
  • Oceans and lakes don’t freeze solid because ice floats.
  • Water is unusual because it is less dense as a solid than as a cold liquid.
  • Most materials contract as they solidify, but water expands.
  • At temperatures above 4°C, water behaves like other liquids, expanding as it warms and contracting as it cools.
  • Water begins to freeze when its molecules are no longer moving vigorously enough to break their hydrogen bonds.
  • When water reaches 0°C, water becomes locked into a crystalline lattice, with each water molecule bonded to a maximum of four partners.
  • As ice starts to melt, some of the hydrogen bonds break, and water molecules can slip closer together than they can while in the ice state.
  • Ice is about 10% less dense than water at 4°C.
  • Therefore, ice floats on the cool water below.
  • This oddity has important consequences for life.
  • If ice sank, eventually all ponds, lakes, and even the ocean would freeze solid.
  • During the summer, only the upper few centimeters of the ocean would thaw.
  •  Instead, the surface layer of ice insulates liquid water below, preventing it from freezing and allowing life to exist under the frozen surface.
  • Water is the solvent of life.
  • A liquid that is a completely homogeneous mixture of two or more substances is called a solution.
  • A sugar cube in a glass of water will eventually dissolve to form a uniform solution of sugar and water.
  • The dissolving agent is the solvent, and the substance that is dissolved is the solute.
  •  In our example, water is the solvent and sugar is the solute.
  • In an aqueous solution, water is the solvent.
  • Water is not a universal solvent, but it is very versatile because of the polarity of water molecules.
  • °         Water is an effective solvent because it readily forms hydrogen bonds with charged and polar covalent molecules.
  • °         For example, when a crystal of salt (NaCl) is placed in water, the Na+ cations interact with the partial negative charges of the oxygen regions of water molecules.
  • °         The Cl− anions interact with the partial positive charges of the hydrogen regions of water molecules.
  • ·         Each dissolved ion is surrounded by a sphere of water molecules, a hydration shell.
  • ·         Eventually, water dissolves all the ions, resulting in a solution with two solutes: sodium and chloride ions.
  • ·         Polar molecules are also soluble in water because they form hydrogen bonds with water.
  • ·         Even large molecules, like proteins, can dissolve in water if they have ionic and polar regions.
  • ·         Any substance that has an affinity for water is hydrophilic (water-loving).
  • °         These substances are dominated by ionic or polar bonds.
  • ·         Some hydrophilic substances do not dissolve because their molecules are too large.
  • °         For example, cotton is hydrophilic because cellulose, its major constituent, has numerous polar covalent bonds. However, its giant cellulose molecules are too large to dissolve in water.
  • °         Water molecules form hydrogen bonds with the cellulose fibers of cotton, allowing you to dry yourself with your cotton towel as the water is pulled into the towel.
  • ·         Substances that have no affinity for water are hydrophobic (water-fearing).
  • °         These substances are nonionic and have nonpolar covalent bonds.
  • °         Because there are no consistent regions with partial or full charges, water molecules cannot form hydrogen bonds with hydrophobic molecules.
  • °         Oils such as vegetable oil are hydrophobic because the dominant bonds, carbon-carbon and carbon-hydrogen, share electrons equally.
  • °         Hydrophobic molecules are major ingredients of cell membranes.
  • ·         Biological chemistry is “wet” chemistry with most reactions involving solutes dissolved in water.
  • ·         Chemical reactions depend on collisions of molecules and therefore on the concentrations of solutes in aqueous solution.
  • ·         We measure the number of molecules in units called moles.
  • ·         The actual number of molecules in a mole is called Avogadro’s number, 6.02 × 1023.
  • ·         A mole is equal to the molecular weight of a substance but scaled up from daltons to grams.
  • ·         To illustrate, how could we measure out a mole of table sugar—sucrose (C12H22O11)?
  • °         A carbon atom weighs 12 daltons, hydrogen 1 dalton, and oxygen 16 daltons.
  • °         One molecule of sucrose would weigh 342 daltons, the sum of weights of all the atoms in sucrose, or the molecular weight of sucrose.
  • °         To get one mole of sucrose, we would weigh out 342 g.
  • ·         The advantage of using moles as a measurement is that a mole of one substance has the same number of molecules as a mole of any other substance.
  • °         If substance A has a molecular weight of 10 daltons and substance B has a molecular weight of 100 daltons, then we know that 10 g of substance A has the same number of molecules as 100 g of substance B.
  • °         A mole of sucrose contains 6.02 × 1023 molecules and weighs 342 g, while a mole of ethyl alcohol (C2H6O) also contains 6.02 × 1023 molecules but weighs only 46 g because the molecules are smaller.
  • °         Measuring in moles allows scientists to combine substances in fixed ratios of molecules.
  • ·         In “wet” chemistry, we are typically combining solutions or measuring the quantities of materials in aqueous solutions.
  • °         The concentration of a material in solution is called its molarity.
  • °         A one molar solution has one mole of a substance dissolved in one liter of solvent, typically water.
  • °         To make a 1 molar (1M) solution of sucrose, we would slowly add water to 342 g of sucrose until the total volume was 1 liter and all the sugar was dissolved.

Concept 3.3 Dissociation of water molecules leads to acidic and basic conditions that affect living organisms

  • ·         Occasionally, a hydrogen atom participating in a hydrogen bond between two water molecules shifts from one molecule to the other.
  • °         The hydrogen atom leaves its electron behind and is transferred as a single proton—a hydrogen ion (H+).
  • °         The water molecule that lost the proton is now a hydroxide ion (OH−).
  • °         The water molecule with the extra proton is now a hydronium ion (H3O+).
  • ·         A simplified way to view this process is to say that a water molecule dissociates into a hydrogen ion and a hydroxide ion:
  • °         H2O <=> H+ + OH−
  • ·         This reaction is reversible.
  • ·         At equilibrium, the concentration of water molecules greatly exceeds that of H+ and OH−.
  • ·         In pure water, only one water molecule in every 554 million is dissociated.
  • °         At equilibrium, the concentration of H+ or OH− is 10−7M (at 25°C).
  • ·         Although the dissociation of water is reversible and statistically rare, it is very important in the chemistry of life.
  • ·         Because hydrogen and hydroxide ions are very reactive, changes in their concentrations can drastically affect the chemistry of a cell.
  • ·         Adding certain solutes, called acids and bases, disrupts the equilibrium and modifies the concentrations of hydrogen and hydroxide ions.
  • ·         The pH scale is used to describe how acidic or basic a solution is.
  •  Organisms are sensitive to changes in pH.
  • ·         An acid is a substance that increases the hydrogen ion concentration in a solution.
  • °         When hydrochloric acid is added to water, hydrogen ions dissociate from chloride ions: HCl -> H+ + Cl−
  • °         Addition of an acid makes a solution more acidic.
  • ·         Any substance that reduces the hydrogen ion concentration in a solution is a base.
  • ·         Some bases reduce the H+ concentration directly by accepting hydrogen ions.
  • °         Ammonia (NH3) acts as a base when the nitrogen’s unshared electron pair attracts a hydrogen ion from the solution, creating an ammonium ion (NH4+).
  • °         NH3 + H+ <=> NH4+
  • ·         Other bases reduce H+ indirectly by dissociating to OH−, which then combines with H+ to form water.
  • °         NaOH -> Na+ + OH−                OH− + H+ -> H2O
  • ·         Solutions with more OH− than H+ are basic solutions.
  • ·         Solutions with more H+ than OH− are acidic solutions.
  • ·         Solutions in which concentrations of OH− and H+ are equal are neutral solutions.
  • ·         Some acids and bases (HCl and NaOH) are strong acids or bases.
  • °         These molecules dissociate completely in water.
  • ·         Other acids and bases (NH3) are weak acids or bases.
  • °         For these molecules, the binding and release of hydrogen ions are reversible.
  • °         At equilibrium, there will be a fixed ratio of products to reactants.
  • °         Carbonic acid (H2CO3) is a weak acid:
  • §         H2CO3 <=> HCO3− + H+
  • §         At equilibrium, 1% of the H2CO3 molecules will be dissociated.
  • ·         In any solution, the product of the H+ and OH− concentrations is constant at 10−14.
  • ·         Brackets ([H+] and [OH−]) indicate the molar concentration of the enclosed substance.
  • °         [H+] [OH−] = 10−14
  • °         In a neutral solution, [H+] = 10−7 M and [OH−] = 10−7 M
  • ·         Adding acid to a solution shifts the balance between H+ and OH− toward H+ and leads to a decline in OH−.
  • °         If [H+] = 10−5 M, then [OH−] = 10−9 M
  • °         Hydroxide concentrations decline because some of the additional acid combines with hydroxide to form water.
  • ·         Adding a base does the opposite, increasing OH− concentration and lowering H+ concentration.
  • ·         The H+ and OH− concentrations of solutions can vary by a factor of 100 trillion or more.
  • ·         To express this variation more conveniently, the H+ and OH− concentrations are typically expressed via the pH scale.
  • °         The pH scale, ranging from 1 to 14, compresses the range of concentrations by employing logarithms.
  • °         pH = − log [H+] or [H+] = 10−pH
  • °         In a neutral solution, [H+] = 10−7 M, and the pH = 7.
  • ·         Values for pH decline as [H+] increase.
  • ·         While the pH scale is based on [H+], values for [OH−] can be easily calculated from the product relationship.
  • ·         The pH of a neutral solution is 7.
  • ·         Acidic solutions have pH values less than 7, and basic solutions have pH values greater than 7.
  • ·         Most biological fluids have pH values in the range of 6 to 8.
  • °         However, the human stomach has strongly acidic digestive juice with a pH of about 2.
  • ·         Each pH unit represents a tenfold difference in H+ and OH− concentrations.
  • °         A small change in pH actually indicates a substantial change in H+ and OH− concentrations.
  • ·         The chemical processes in the cell can be disrupted by changes to the H+ and OH− concentrations away from their normal values, usually near pH 7.
  • ·         To maintain cellular pH values at a constant level, biological fluids have buffers.
  • ·         Buffers resist changes to the pH of a solution when H+ or OH− is added to the solution.
  • °         Buffers accept hydrogen ions from the solution when they are in excess and donate hydrogen ions when they have been depleted.
  • °         Buffers typically consist of a weak acid and its corresponding base.
  • °         One important buffer in human blood and other biological solutions is carbonic acid, which dissociates to yield a bicarbonate ion and a hydrogen ion.
  • °         The chemical equilibrium between carbonic acid and bicarbonate acts as a pH regulator. The equilibrium shifts left or right as other metabolic processes add or remove H+ from the solution.
  • Acid precipitation threatens the fitness of the environment.
  • ·         Acid precipitation is a serious assault on water quality in some industrialized areas.
  • °         Uncontaminated rain has a slightly acidic pH of 5.6.
  • °         The acid is a product of the formation of carbonic acid from carbon dioxide and water.
  • ·         Acid precipitation occurs when rain, snow, or fog has a pH that is more acidic than 5.6.
  • ·         Acid precipitation is caused primarily by sulfur oxides and nitrogen oxides in the atmosphere.
  • °         These molecules react with water to form strong acids that fall to the surface with rain or snow.
  • ·         The major source of these oxides is the burning of fossil fuels (coal, oil, and gas) in factories and automobiles.
  • ·         The presence of tall smokestacks allows this pollution to spread from its site of origin to contaminate relatively pristine areas thousands of kilometers away.
  • °         In 2001, rain in the Adirondack Mountains of upstate New York had an average pH of 4.3.
  • ·         The effects of acids in lakes and streams are more pronounced in the spring during snowmelt.
  • °         As the surface snows melt and drain down through the snowfield, the meltwater accumulates acid and brings it into lakes and streams all at once.
  • °         The pH of early meltwater may be as low as 3.
  • ·         Acid precipitation has a great impact on the eggs and the early developmental stages of aquatic organisms that are abundant in the spring.
  • ·         Thus, strong acidity can alter the structure of molecules and impact ecological communities.
  • ·         Direct impacts of acid precipitation on forests and terrestrial life are more controversial.
  • ·         However, acid precipitation can impact soils by affecting the solubility of soil minerals.
  • °         Acid precipitation can wash away key soil buffers and plant nutrients such as calcium and magnesium ions.
  • °         It can also increase the concentrations of compounds such as aluminum to toxic levels.
  • °         This has done major damage to forests in Europe and substantial damage of forests in North America.
  • °         Progress has been made in reducing acid precipitation.

 

 

Chapter 3 – Water Objectives

 

 

Chapter 3   Water & the Fitness of the Environment
Objectives
The Properties of Water

1.        With the use of a diagram or diagrams, explain why water molecules are:

a.         polar

b.         capable of hydrogen bonding with four neighboring water molecules

2.        List four characteristics of water that are emergent properties resulting from hydrogen bonding.

3.        Define cohesion and adhesion. Explain how water’s cohesion and adhesion contribute to the movement of water from the roots to the leaves of a tree.

4.        Distinguish between heat and temperature, using examples to clarify your definitions.

5.        Explain the following observations by referring to the properties of water:

n                      Coastal areas have milder climates than adjacent inland areas.

n                      Ocean temperatures fluctuate much less than air temperatures on land.

n                      Insects like water striders can walk on the surface of a pond without breaking the surface.

n                      If you slightly overfill a water glass, the water will form a convex surface above the top of the glass.

n                      If you place a paper towel so that it touches spilled water, the towel will draw in the water.

n                      Ice floats on water.

n                      Humans sweat and dogs pant to cool themselves on hot days.

6.        Distinguish among a solute, a solvent, and a solution.

7.        Distinguish between hydrophobic and hydrophilic substances.

8.        Explain how you would make up a one molar (1M) solution of ethyl alcohol.

The Dissociation of Water Molecules

9.        Name the products of the dissociation of water and give their concentration in pure water.

10.       Define acid, base, and pH.

11.       Explain how acids and bases may directly or indirectly alter the hydrogen ion concentration of a solution.

12.       Using the bicarbonate buffer system as an example, explain how buffers work.

13.       Briefly explain the causes and effects of acid precipitation.

 
BACK

 

Chapter 1 Worksheet BI

 

Biology – Science of Life

 

Section 1-1            Themes of Biology 

1. How many species of organisms are estimated to inhabit the Earth? 

2. About how many species have actually been identified? 

3. When did the first life forms probably arise on Earth? 

4. What was the first organism like? 

5.  What is an organism? 

6. Most unicellular organisms can only be seen with a __________________________.

7. Where did these first cells live? 

8. Over time, organisms _____________ and new kinds of _______________ arose from ___________ of organisms and came to inhabit every _______________ of the Earth. 

9. Define biology. 

10. Name several things that the study of biology would include. 

11. The study of biology is unified by ___________________________. 

12. Name 6 unifying themes of biology. 

13. What is a cell? Where are they found? 

14. What is the difference between a unicellular & a multicellular organism? 

15. Cells are ___________________ but highly _______________________.

16. Are all cells alike? Explain. 

17. All cells are surrounded by a ____________ & contain _____________  instructions. 

18. Genetic instructions are used by cells to make new __________ and new cell ________________. 

19. How do new cells produced by unicellular organisms compare to the parent unicellular organism? 

20. How do mature multicellular organisms begin their life? 

21. If multicellular organisms begin their life as one cell, how do they have so many cells? Explain. 

22. Give an example of an organism maintaining a stable internal environment. 

23. Define homeostasis and tell whether it occurs in unicellular&/or multicellular organisms. 

24. Genetic information is passed to offspring during _______________________.

25. What molecule contains the cell’s hereditary information?

26.  How does DNA exist in multicellular organisms? In unicellular organisms? 

27. What is a gene?

28. In multicellular organisms, each body cell has an _____________ copy of its DNA.

29. Does each cell in a multicellular organism use all the genes on its DNA? Explain. 

30. Explain sexual reproduction. 

31. When a sperm joins with an egg to make a fertilized cell, what happens next to this cell? 

32. New organisms from sexual reproduction have _______________ material from both parents.

33. Explain asexual reproduction. 

34. Name a unicellular organism that reproduces by asexual reproduction.

35. New cells or organisms from asexual reproduction have ____________ genetic information. 

36. What is evolution? 

37. Do individuals or populations evolve? 

38. What is the driving force for evolution? 

39. Explain natural selection and give an example. 

40. Competition for what types of resources drives natural selection? 

41. Why is it so important to a species for members to survive? 

42. Organisms that survive and reproduce are ones with ________________ traits.

43. Define ecology. 

44. What are ecosystems and give an example? 

45. Name 3 things organisms must get from the environment to survive. 

46. What has been the effect of man’s activities on many ecosystems? 

47. Living things are _______________ and need a constant supply of _________________.

48. What process supplies energy for organisms on Earth? 

49. Define autotroph. 

50. Autotrophs trap _________ and use this energy to combine __________ and  _________ into__________ and starches. 

51. Define heterotrophs. 

52. Give an example of an autotroph.

53. Give several examples of heterotrophs. 

54. How do heterotrophs get their food? 

Section 1-2            World of Biology 

55. List 6 characteristics shared by all living things. 

56. All living things composed of  _________________.

57. Cells may be specialized in _________________ organisms. What does this mean?

58. What is always true about cell size?

59. Living things are organized at what 2 levels?

60. How are cells organized in multicellular organisms?

61. Define metabolism. 

62. The energy from metabolism is used for ___________, ___________, and _______________ of organisms. 

63. What is homeostasis and give an example? 

64. Is growth a characteristic of living and nonliving things? Explain.

65. What 2 things must occur for living things to grow?

 

 

66. Define cell division.

 

 

 

67. Define development.

 

 

 

68. Explain why development is necessary for multicellular organisms?

 

 

 

69. Is reproduction essential to the survival of a species? Explain.

 

 

 

70. Sexual reproduction produces offspring ______________________ to the parents.

 

71. Have all organisms been identified? Explain.

 

 

Section 1-3            Scientific Method

 

72. Scientists solve problems using the ___________________________.

 

73. The first step of the scientific method is when scientists make ___________________ of the natural world.

74. Define data.

 

 

 

75. What does a scientist usually employ in making their observations?

 

 

76. What is quantitative data?

 

77. What is sampling & why is it used by scientists?

 

 

 

78. What 2 things must be true for samples to be useful?

 

 

79. To be useful, data must be _____________________ into ____________________,

 

       ______________________, and _____________________, or maps.

 

80. Once an observation is made, the second thing a scientist must do is to develop a(n) ___________________________.

 

81. Define hypothesis.

 

 

 

82. All hypotheses must be __________________ to give supporting evidence.

 

83. What is a prediction & how are they usually written?

 

 

 

84. What is an experiment?

 

 

 

85. Name the 2 groups in a controlled experiment.

 

 

86. Both groups in an experiment are identical except for ___________ factor called the

 

       ___________________.

87. Name the 2 types of variables in a controlled experiment.

 

 

 

88. After data is collected and organized, it must be __________________ to tell if it is reliable.

 

89. If experimental data does not support the hypothesis, what should be done?

 

 

90. What is a scientific model?

 

 

91. What is an inference?

 

 

 

92. How is a theory formed?

 

 

93. Define theory.

 

 

94. What is the difference between a field biologist and a laboratory biologist? Do they both use the scientific method?

 

 

 

 

95. What do scientists do with the results of their scientific studies?

 

 

 

Section 1-4            Microscopes & Measurement

 

96. What is a microscope?

 

 

97. What is the difference between resolution & magnification?

 

 

 

 

98. When would a microscope be used?

 

 

99. Do all microscopes have the same magnification & resolution?

 

100.  Draw and label the parts of a light (LM) microscope.

 

 

 

 

 

 

 

 

 

101. Tell the function of each of these parts of an LM — stage, light source, objective lens, ocular lens, & nosepiece.

 

 

 

 

 

 

102. To view specimens with a light microscope, they must be placed on a _______________

 

         and be ________________ so light will pass through to the lenses & your eyes.

 

103. What is the power of magnification & explain how it is determined?

 

 

 

104. Light microscopes can only magnify up to ______________ before the image becomes blurry.

105. What type of scope is used to view viruses & cell parts?

 

106. What produces an image with the electron microscope?

 

 

107. Name the 2 main types of electron microscopes.

 

108. What is the highest magnification for the TEM? For the SEM?

 

 

109. Can electron microscopes be used to view living cells?

 

110. What type of scope gives a magnified view of an object’s surface?

 

111. What is the standard unit of measurement used by scientists?

 

112. Name the SI base units, what they measure, & give their abbreviations (table 1-1, page 23)

 

 

 

 

 

 

113. The SI system is based on units of __________ with designated _________________.

 

114. Give the SI prefix for these base units — 1000, .01, .001, .000001, .000000001, & .000000000001.

 

 

 

 

 

115. Give the SI unit for area, volume, and time.

 

 

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 20 AP Objectives

 

Chapter 20    DNA Technology
Objectives
DNA Cloning
1. Explain how advances in recombinant DNA technology have helped scientists study the eukaryotic genome.
2. Describe the natural function of restriction enzymes and explain how they are used in recombinant DNA technology.
3. Explain how the creation of sticky ends by restriction enzymes is useful in producing a recombinant DNA molecule.
4. Outline the procedures for cloning a eukaryotic gene in a bacterial plasmid.
5. Describe techniques that allow identification of recombinant cells that have taken up a gene of interest.
6. Define and distinguish between genomic libraries using plasmids, phages, and cDNA.
7. Describe the role of an expression vector.
8. Describe two advantages of using yeast cells instead of bacteria as hosts for cloning or expressing eukaryotic genes.
9. Describe two techniques to introduce recombinant DNA into eukaryotic cells.
10. Describe the polymerase chain reaction (PCR) and explain the advantages and limitations of this procedure.
11. Explain how gel electrophoresis is used to analyze nucleic acids and to distinguish between two alleles of a gene.
12. Describe the process of nucleic acid hybridization.
13. Describe the Southern blotting procedure and explain how it can be used to detect and analyze instances of restriction fragment length polymorphism (RFLP).
14. Explain how RFLP analysis facilitated the process of genomic mapping.
DNA Analysis and Genomics
15. Explain the goals of the Human Genome Project.
16. Explain how linkage mapping, physical mapping, and DNA sequencing each contributed to the genome mapping project.
17. Describe the alternate approach to whole-genome sequencing pursued by J. Craig Venter and the Celera Genomics company.
18. Explain how researchers recognize protein-coding genes within DNA sequences.
19. Describe the surprising results of the Human Genome Project.
20. Explain how the vertebrate genome, including that of humans, generates greater diversity than the genomes of invertebrate organisms.
21. Explain how in vitro mutagenesis and RNA interference help researchers to discover the functions of some genes.
22. Explain the purposes of gene expression studies. Describe the use of DNA microarray assays and explain how they facilitate such studies.
23. Define and compare the fields of proteomics and genomics.
24. Explain the significance of single nucleotide polymorphisms in the study of the human evolution.
Practical Applications of DNA Technology
25. Describe how DNA technology can have medical applications in such areas as the diagnosis of genetic disease, the development of gene therapy, vaccine production, and the development of pharmaceutical products.
26. Explain how DNA technology is used in the forensic sciences.
27. Describe how gene manipulation has practical applications for environmental and agricultural work.
28. Describe how plant genes can be manipulated using the Ti plasmid carried by Agrobacterium as a vector.
29. Explain how DNA technology can be used to improve the nutritional value of crops and to develop plants that can produce pharmaceutical products.
30. Discuss the safety and ethical questions related to recombinant DNA studies and the biotechnology industry.
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