Exam 3 Review Sheet
AP Biology
USE THIS SHEET AFTER YOU HAVE STUDIED AS A METHOD OF EVALUATING WHAT YOU
STILL NEED TO WORK ON…OR DO WHAT YOU ARE LIKELY TO DO, WHICH IS WHATEVER
YOU WANT. SERIOUSLY, STUDY FIRST AND THEN FLY THROUGH MOST OF THESE.
Exam 3 will cover:
Chapter 3 – ALL
Chapter 4 – ALL
Make sure you look at ALL the questions in chapter 3 as well as bold words,
headings and figures, the PowerPoint, and this review sheet.
CHAPTER 3:
96. Write out a chemical reaction. It can be any reaction you want. Indicate the reactants
and the products. What is a chemical reaction?
97. Compare a reversible to an irreversible reaction.
97.5. Enzymes bring chemical reactions to equilibrium…explain.
100. Define the term electronegativity.
101. Describe what makes one element more electronegative than another element.
For example why is oxygen more electronegative than carbon? Why is oxygen more
electronegative than sulfur? Make sure you can explain the trend of the periodic
table in terms of nuclear charge and distance from the nucleus.
102. When two elements with very different electronegative values form a covalent
bond this bond is more specifically called what? Explain why this name is given to
this bond (hint: Earth, magnets, personality disorder).
103. If two elements with similar electronegativity values form a covalent bond, this
is what type of covalent bond? Compare this type of bond to the type of bond in
question number 5.
104. List every polar covalent bond we spoke about in class (Ex. O-H) and indicate
partial charges. What makes these bonds polar covalent?
105. List every non-polar covalent bond we discussed in class (Ex. O=O, C-C). What
makes these bonds non-polar?
106. In class we said that the C-H covalent bond is non-polar even thought carbon
has more protons than hydrogen. Explain why in detail.
107. Is water a polar or non-polar molecule? Explain why.
108. You should be able to identify polar and non-polar covalent bonds in any
molecule I give you like we did in class.
109. Give one example of a non-polar molecule. Why is this molecule considered
non-polar?
110. Explain how multiple water molecules would interact with each other in a glass
of water.
111. Explain why a single water molecule can hydrogen bond to four other water
molecules in terms of loan pair electrons.
112. What makes a hydrogen bond different from an ionic or covalent bond? Order
these in terms of bond strengths.
113. Explain why hydrogen bonds are able to form between two water molecules in
terms of the properties of the individual atoms that make up a water molecule.
2. Describe the structure of water in as much detail as possible including its
geometry, partial charges, etc…
3. All of waters properties having to do with cohesion/adhesion come down to the
atomic number of the elements that make it up…explain.
4. When one pours glucose crystals into water, the glucose “dissolves”. Explain what
this means on the atomic level. What does it mean to “dissolve”? Include hydration
shell in your response. Did you break any covalent bonds???? I hope NOT!!
5. If I were to pour oil into water, the two substances do not mix. Explain why on the
molecular level. Why can water dissolve salt or glucose, but not oil? Make sure to
include hydration cage.
5.5 Explain why the oil molecules will eventually come together and not readily
separation…make sure to discuss hydrogen bonding, random molecular motion,
hydration cage, etc…
6. What is meant by like dissolves like.
6.5. Explain why just because a molecule has polar covalent bonds it does not mean
that is will be highly soluble (dissolvable) in water.
7. How is a water molecule held together? How are water molecules held to each
other?
8. In class we discussed many properties of water. The first was cohesion, which
leads to high surface tension, a high specific heat, moderation of temperature, a high
heat of vaporization (perfect for evaporative cooling), transpiration and more.
Describe cohesion and these resulting properties in detail.
9. The second property was adhesion. Adhesion allows water to be the “universal
solvent” and makes water extremely sticky. Describe adhesion on the molecular
level. What types of molecules would be involved in adhesion with water? Explain
how adhesion makes water the universal solvent and is critical for dissolving
substances. Draw a picture of a single sodium atom surrounded by water to
demonstrate adhesion, universal solvent and dissolving all in one picture. Why is
adhesion, like all the other properties, so important for life on Earth as we know it?
Think about all the substances in and around your cells…
10. Explain the difference between solution, aqueous solution, solvent and solute.
Give an example of an aqueous solution and identify the solvent and solute. Give an
example of a non-aqueous solution and do the same.
11. What is heat on the molecular level? Explain how you heat something up (on the
MOLECULAR LEVEL). Don’t just say I turn on the stove. What is the difference
between heat and temperature? Use an example to make this crystal clear.
12. Explain why wood burns and why in order for wood to burn, molecular oxygen
needs to be present. Why does this process generate so much “heat”. Your
explanations need to be specific on the molecular level.
13. Explain the importance of water’s high specific heat for life on Earth. Consider
the water that makes you up as well as the water of the oceans. What does it mean
that water has a high specific heat?
14. Explain why water has a high heat of vaporization. Why is this important in
terms of evaporative cooling?
15. Explain how evaporative cooling works on the molecular level. Why does water
secreted onto your skin cool your body down if the water you are secreting is as
warm as your body? Make sure to discuss the definition of temperature as being the
average of molecular motion.
16. Explain, on the molecular level, why ice is less dense than water and therefore
floats. Why is this unique property of water so important for life on Earth? What
would happen otherwise? Explain how the freezing of water can crack boulders on
the molecular level?
17. In class we discussed that the water in any aqueous solution is not all in the form
of H2O. A very tiny amount breaks apart into H+ and OH-. Write out this chemical
reaction on paper and explain why this happens. What do we call these ions?
18. Define the term concentration. Go through the powerpoint slides and make sure
you understand the concept of concentration.
19. What does Molarity mean?
20. In terms of numbers of molecules dissolved, how many glucose molecules would
be in 0.5L of a 6M (molar) glucose solution?
21. If I have 250ml of the solution in the question above, how many glucose
molecules would be in the solution I am holding in moles.
22. Be able to explain how to make solutions as shown in the PowerPoint. Here is a
link to some Molarity questions with answers:
http://www.chemteam.info/Solutions/WS-Molarity.html
You do not need to do them all! It is just extra practice if you wish. There are also
practice problems in the book at the end of chapter 3 and on page 53.
23. Logarithms. What does it mean to take the logarithm (log; base 10) of a number
(in words)?
24. What is the Molar concentration of H+ in pure water. What does this mean in
terms of actual numbers of free protons? Show how you would convert this
concentration to pH.
25. What does the pH tell us about a solution? Why don’t we just discuss the
concentration of H+ instead of using pH.
26. What do we think pH stands for?
27. What is the pH of a solution with an [H+] = 10-3? What is the [H+] of a solution
with a pH of 9? Do not forget units!!
28. Compare and contrast the terms acidic, basic and neutral. Give examples of
solutions of each.
29. Which is more acidic, a solution with an [H+] = 10-4 or a solution with an
[H+] = 10-9? Explain. What would the pH of these solutions be?
30. Discuss the relationship between the [H+] and the pH in terms of when one goes
up, what happens to the other? Explain why this relationship exists.
31. As the H+ concentration changes 10-fold (10X higher or 10X low [H+]), what
happens to the pH?
32. If the [H+] of one solution is 10,000X times higher than a second solution that
has a pH of 8, what is the pH of the first solution?
33. There are two solutions in front of you. One has a pH 9 and the other a pH 13.
First, are these solutions basic, acidic or neutral? Second, how much higher or lower
is the [H+] of the pH13 solution compared to the pH9 solution?
34. Compare and contrast an acid to a base. Make sure to define each. Give an
example of an acid and a base.
35. Sodium hydroxide (Na+OH-) is considered a very strong base. Explain how, on
the molecular level, this compound acts like a base. (How does it remove free
protons from a solution?)
36. If I had a solution with a measured pH of 3, what would I need to do in order to
change the pH to 7? What if I wanted to change it to 1?
37. If you add acid to a solution, does it become acidic or basic? You don’t know!
Explain.
38. If I give you a solution, blue litmus paper and red litmus paper, explain how you
would determine if this solution is acidic or basic.
39. Which would give you a more accurate pH value, red litmus paper, blue litmus
paper or hydrion paper? Explain.
40. Your fish looks very sick and you wonder if the pH is off. They typically prefer to
be in a pH below 8. You have with you a bottle of phenolphthalein. Explain how you
would determine if you have a pH problem.
41. Swimming pools should be kept at a pH around neutral. You are going to have a
huge pool party tomorrow night and need to make sure that the pH is correct. You
have with you a bottle of bromothymol blue. Explain how you would use the
bromothymol blue to determine the pH of the pool. What would you expect to find if
the pool has a pH of 4? of 7? of 10?
42. You measured the pH of the pool and found it to be 12!! What do you need to do?
43. You inform Dr. T of your pool problem and he suggests adding a buffer. What is a
buffer and how would it help? Explain exactly how a buffer works in terms of le
Chatlier’s Principle.
44. Are buffers permanent? Will they force the pH to never change? Explain.
44.5 Explain what the pKa value tells you and how knowing these values for various
buffers can help you make a decision on what buffer to use in a lab or swimming
pool…
45. In class we said that the normal pH of rain is around 5.5. Explain why it is acidic and
not neutral. Write the chemical formula that explains this.
46. Explain acid precipitation (why is occurs) and why it is a problem to organisms on
the molecular level.
47. Define the calorie and the Joule.
48. Approximately how fast do molecules move at room/body temperature?
49. What is a hydronium ion and how does this ion relate to pH of a solution?
50. Describe van der Waals interactions. Why do they occur in terms of electron
positions? Between what types of molecules do van der Waals occur? These interactions
are the weakest relative to the bonds we have spoken about. When will they be
significant?
51. Describe how transpiration works and be sure to discuss adhesion, cohesion, solar
energy, stomates, xylem and H-bonding.
52. Make sure you can do specific heat problems. What is the specific heat of water in
calories and Joules? There are practice problems on the notes section of the website.
53. Compare a food calorie to a calorie.
54. Compare pH to pOH. If I know the pOH, how can I calculate the pH? If I know the
[OH-], how can I determine the pOH and then the pH?
55. I will post some molarity problems under the notes section…make sure you can do
them.
56. Do not forget that I will put conversions on the test again. I will try and put a review
sheet on the notes section for you to practice.
CHAPTER 4
0.5. Make sure you can do number conversions / scientific notation
1. Explain on the molecular level what happens to a crystal composed of glucose when
you throw it in water (don’t just say that it dissolves).
2. Compare and contrast organic compounds and hydrocarbons. Why are all
hydrocarbons organic, but not all organic molecules are hydrocarbons?
3. Be able to identify/circle the carbon skeleton of a hydrocarbon.
4. Compare and contrast structural, geometric and enantiomeric isomers. Be able to draw
examples of each.
5. Explain how you can determine by looking at a molecule if it will have an
enantiomeric isomer in terms of an asymmetric carbon.
5.5 Describe the requirements for a carbon to be considered asymmetric
6. We said the all amino acids are “L” and all sugar are “D”. What does this refer to?
7. What is meant by “cis” and “trans” in terms of geometric isomers? Draw a pair of cistrans (geometric) isomers using a double bond as well as a ring. Why is the double bond /
ring necessary? What property makes it necessary?
8. Be able to convert a skeletal formula into the corresponding structural formula.
9. Be able to draw a small molecules like methane, ethane, propane, ammonia etc…
three-dimensionally using the wedge-and-dash projection learned in class. THERE IS A
REVIEW SHEET ONLINE UNDER THE NOTES TAB!!
10. Compare tetrahedral to planar molecular shapes and draw examples of each. Be able
to explain why each forms. Carbon with four attachments would form what shape?
Carbon with three attachments? Two attachments?
11. Describe the formation of early Earth in terms of the Great Bombardment,
atmosphere formation, ocean formation, etc…
12. How are elements larger than lithium formed? Where did the carbon, oxygen,
nitrogen, phosphorus, sulfur, zinc, molybdenum, etc… that make up you come from since
universe began with just hydrogen, helium and perhaps some lithium?
12.5 Describe a leading hypthosis concerning the formation of the moon. Why is
complex life like ourselves so thankful for such a large moon?
13. Describe the Miller-Urey experiment and explain the results/conclusions.
14. Define teratogen and give an example.
15. Know the functional group table in the textbook. Be able to draw the functional
groups.
16. Identify the functional groups in the molecule below (there are three):
17. I will ask you a question similar to: draw an organic molecule that contains a
sulfhydryl group and is an acid.
18. Compare and contrast aldehydes to ketones. Why are these NOT considered
functional groups, but rather a type of compound? What is the functional group found in
each of these? Which class of macromolecule contains both types of compounds?
19. Which types of compounds usually end in –ol? What functional group is involved?
20. Describe in detail why carboxylic acids are considered acids (You will need to
discuss electronegativity differences between atoms). What class of macromolecules
contain this functional group?
21. Compare and contrast the carbonyl to the carboxyl group. Which two functional
groups add to be a carboxyl? Reminder: if the group is a carboxyl, it is not also a
carbonyl and a hydroxyl (You cannot circle the C=O in question 6 and say it’s a
carbonyl. Why?)
21.5 Which functional groups act as acids? Which will act as a base?
21.75 Describe the importance of the sulfhydryl group in proteins. Reminder: cysteine
and cystine.
22. Compare hydrophobic to hydrophilic. What makes something hydrophobic or
hydrophilic? Describe how this works on the MOLECULAR LEVEL. Don’t just say
nonpolar or polar… Speak in terms of electronegativity and the electromagnetic force.