Camp's Biochemistry and Cell Biology by the Numbers

Kenneth R. Camp

Buckethead Publishing

Copyright © [Year of First Publication] by [Author or Pen Name]

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No portion of this book may be reproduced in any form without written permission from the publisher or author, except as permitted by U.S. copyright law.

Contents

1.Chemistry Basics & The Periodic Table

2.Chemical Bonding, Molecules, and Compounds

3.The Chemistry of Water

4.Chemical Measurements and Reactions

5.Organic Chemistry as Applied to Biological Molecules

6.Carbohydrates

7.Lipids

8.Proteins

9.Nucleic Acids

10.Biological Energetics and Enzymes

11.Cell Biology: History, Organelles, and Functions

12.Cell Homeostasis: Membranes, Walls, and More

13.Cellular Respiration and Metabolism

14.Photosynthesis: Autotrophic Energy Production

15.Cell Division: Mitosis

16.Cell Division: Meiosis

17.GLOSSARY

Chapter one

Chemistry Basics & The Periodic Table

A) Matter and Elements

1.Matter takes up space and has mass.

2. All matter is composed of 118 basic elements (depending on who is trying to create a new element that day).

3. Only 92 elements are found in nature. Of these, only about 40 to 50 of these are of any consequence in biology.

4. Since the time of Mendeleev, who developed the first crude periodic table, all of these are symbolized with one or two letters. Many symbols don’t match the name, because they are derived from Latin.

5. For instance, C is the symbol for carbon, but Na is the symbol for sodium, since the Romans called it natrium.

6. Elements cannot be broken down by CHEMICAL reactions to substances with different chemical or physical properties. However, nuclear reactions CAN do this for unstable isotopes.

7. Chemical reactions involve the transfer of electrons from the outer shells of atoms, whereas nuclear reactions involve movement of the much heavier neutrons and protons in the nucleus.

8. Six elements make up the majority of the molecules found in life and are ALWAYS found in all living cells. These elements make an easy acryonym (C,H,O,N,P,S). The CHONPS chart is shown below.

9. The next list below shows most of the the rest of the MOST FREQUENTLY OCCURING common elements in living organisms and what some of their major purposes in living organisms are.

B) Basic Principles of Atomic Structure

1.The chemical and physical properties of atoms (e.g. mass) depend on their subatomic particles

2. Different atoms contain specific numbers of protons, neutrons, and electrons.

3. Protons have a positive charge, occupy the nucleus of an atom, and consequently define the identity of each atom, since they ultimately affect the chemical properties of that element.

4. The number of protons in the nucleus of an atom is known as the atomic number.

5. For instance, carbon atoms have 6 protons in their nucleus, yielding an atomic number of 6.

6. Neutrons also occupy space in the nucleus of an atom, but their number varies, depending on what creates stable conformations for the nucleus. They have an overall charge of zero.

7. The sum of the protons and neutrons provides the atomic mass of an atom.

8. Atomic mass can vary within the same element, based on the number of neutrons. Consequently, the published atomic mass, such as 12.011 for carbon, is an average of all its isotopes (more on isotopes shortly).

9. The strong nuclear force packs protons and neutrons tightly into the nucleus at great distance from the electrons that orbit the atom. Most of an atom is empty space.

10. The mass of protons and neutrons are about 3600 times greater than electrons.

11. While it is beyond the scope of a biology course to go into great detail, it is worth discussing a couple of atomic models so that atomic bonding is more easily understood.

12. Two useful representations of the atom for understanding bonding are those of Niels Bohr and Erwin Schrodinger, which are still commonly seen. We will discuss both in detail in the next section.

13. The Bohr model is the older of the two models and represents atoms with electrons orbiting the nucleus in fixed concentric paths, while the Schrodinger model accounts for the dual nature of electrons as particles AND as waves.

14. Instead of fixed paths, electrons are said to travel in clouds, based on the probability of their location at a given moment. This model of the atom was developed with increasing understandings of nuclear physics. Both models are shown below.

15. For now, let’s briefly summarize the very basics. First, the number of protons in the nucleus of the atom give it a specified number of positive charges.

16. Consequently, an uncharged and unbonded atom needs to pick up the correct number of electrons to balance these positive charges with negative charges and attain a net charge of zero.

17. These electrons then distribute themselves in concentric shells around the nucleus. The arrangement of these electrons determines chemical behavior and the types of bonds it can form.

18. Now, let’s go on to the specifics.

C) Electron Behavior and Chemical Bonding

1. While it is beyond the scope of a biology course to cover a great breadth of chemistry concepts, it is useful to know how electrons behave to give a greater understanding of how the periodic table is laid out.

2. Understanding electron configurations also allows one to understand bonding rules.

3. Let’s take a second deeper look at two models of the atom that are most useful for biological chemistry.

4. The Bohr model holds that electrons orbit the nucleus of the atom in specific energy levels. Inner shells are more tightly bonded to the nucleus and take less energy to hold.

5. As atoms grow in size, the electrons get further away from the nucleus and harder to retain.

6. Atomic stability increases with a full outer shell or a half-filled outer shell. The closer a shell gets to being empty or full, the more reactive that element will be.

7. Now let’s direct our attention to the periodic table and cover some basic concepts.

8. Note the atomic diagram and the periodic table representation of a carbon atom below.

9. Carbon has 6 protons, therefore it has an atomic number of 6.

10. Two Bohr model of the atom are shown below. Isotopes have the same atomic number (and are the same type of element), but have varied numbers of neutrons.

11. The isotope on the left is Carbon-12 with an atomic mass of 12, since it has 6 protons and 6 neutrons.

12. The isotope on the right is Carbon-14 with an atomic mass of 14, since it has 6 protons and 8 neutrons.

13. The published atomic mass in the table is 12.011, because it is the mass of an ‘average’ carbon atom, since 11 out of every 1000 carbon atoms have extra neutrons and are either Carbon 14 or Carbon 13 (not shown).

14. Isotopes can be useful in biological and medical studies, since the less stable forms of the isotope lose neutrons and decay at predictable rates.

15. For instance, Carbon 14 decays to Carbon 12 with a half-life of around 5,300 years. Suppose that a biological sample, such as a frozen wooly mammoth carcass contains 25% of the carbon 14 found in a living elephant. This means that it has undergone around 2 half-lives and the fossil would date to 10,600 years ago.

16. Other isotopes, such as the iodine tracers used for thyroid exams, give off radioactivity that shows up on X-rays and provides medical imaging.

17. The table at the below covers several biologically useful isotopes and examples of their use.

18. Now let’s turn our attention to the arrangement of the periodic table, along with the quantum theory about how electrons distribute themselves in shells, and then we will return to explaining the oxidation state above.

19. Use the periodic table below to follow along.

20. Note that carbon is on the second row of the periodic table at position 4. Carbon has 6 protons, therefore the atom needs 6 electrons to balance out these positive charges.

21. However, to understand why it is positioned on the second row, we need to cover the basics of quantum mechanics described by Erwin Schrodinger.

22. While the Bohr model of the atom shows electrons orbiting as spheres on distinct shells, the Schrodinger model says that since electrons travel at the speed of light, their positions cannot be known.

23. Instead, the Schrodinger model represents electrons as clouds, with the highest probability of electrons being present in prescribed 3-D shapes that stack inside one another like Russian nested dolls.

24. The periodic table is a direct evolution of Dmitri Mendeleev's original attempt to organize elements into families by common properties. The full and correct periodic table explains and organizes elements according to the derivation of Schrodinger's equations into organized schemes. A copy is shown on the next page.

25. Now enjoy a periodic study break before the periodic table.

25. Combining both models gives a clearer understanding of how the periodic table is laid out and why carbon is in this position.

26. The Schrodinger Model is actually a series of calculus equations that solve to show that there are 7 energy levels that stack within one another. These correspond to the 7 periods on the table (horizontal rows).

27. Schrodinger also solved for 4 principal quantum numbers, represented as the letters s, p, d, and f.

28. These describe energy sub-levels that are kind of like drawers on a refrigerator shelf. These are the grids referenced within the prior periodic table.

29. For instance, the 3rd shelf of most refrigerators has a small crisper drawer for your lettuce and a larger main shelf where casseroles go to die, only to be discovered months later dripping with slime and blue fur.

30. The chart below shows how many electrons can be held by each sub-shell.

31. If you look back at the periodic table, you will notice that Groups 1 and 2 put their valence electrons into the s-subshell (2 electrons), so elements like H, Li, and Na have a single s-electron on their outermost shell, while Be, Mg, and Ca have two.

32. The p-shell holds 6 electrons, so that each of the three periods have a total of 8 electrons on their outer shell, such as the elements H, Li, and Na on the left, and He, Ne, Ar on the far right.

33. The former have one valence electron, while the latter noble gases have a full shell with 8 electrons and fulfil the octet rule.

34. The d-orbital holds 10 electrons and is responsible for the 'valley' of transition metals. There is a special set of rules for how p and d orbitals fill preferentially, and these are not particularly relevant to most of the small elements seen in biochemistry.

35. The f-orbital holds 14 electrons and is used by radioactive lanthanide and actinide elements. These are, again, not relevant to the small elements that make up the cells and tissues of living organisms.

36. Now let’s look back at the carbon atom and how it distributes its electrons.

37. The first shell on the table contains only the 1s orbital and is limited to two electrons. This is why hydrogen and helium are the only occupants of that row.

38. Since a carbon atom can only stick of its 6 electrons intgo this shell, it must move on and use the second energy level to distribute the remaining four. Two go into the 2s orbital, while the final two must go into the 2p orbital.

39. This means that the 2s orbital is now filled, but only 2 of 6 slots are occupied in the p sub-shell, so there are 4 empty spaces that can be used to form chemical bonds with other atoms.

40. The same rules apply to all other atoms. Let's do one more example with a calcium atom.

41. Calcium has an atomic number of 20 and an atomic mass of 40, meaning that the typical calcium atom has 20 protons and 20 neutrons comprising its nucleus.

42. Notic ein the Bohr diagram below that the 1st level contains 2 electrons, leaving 18 more that need to be distrubuted.

43. The 2nd and 3rd rows contain spaces for 8 electrons (2s and 6p slots), and both fill and reach the octet configuration.

44. The final 2 electrons then occupy the 4s shell, finishing the Bohr diagram. Only 2 of the 8 possible slots on the 4th energy level are filled, meaning that calcium will behave like a metal, since it is easier to dump 2 electrons than to gain 6.

45. Consequently, calcium will have an oxidation number of +2, because losing 2 electrons will cause an immediate charge imbalance between the nucleus and the electron orbits.

46. There will still be 20 protons, but only 18 electrons for a net charge of +2. Therefore, it will seek out some combination of two negative charges from non-metallic ions to form a stable ionic bond and a salt compound.

47. As it turns out, the only thing that matters, with regard to chemical properties, is the number of valence electrons possessed by an atom. They give the value of the oxidation state when this atom loses these electrons to create a bond.

48. Let's give another example. Oxygen has 8 overall electrons, but only 6 are in the valence shell on the 2nd energy level. Oxygen is said to have an electron configuration of 1s2 2s2 2p4 and wants to get to the stable nobel gas configuration of 1s2 2s2 2p6 like neon, so that it is completely stable. It will need to pick up 2 more electrons to do this.

49. Once this happens, oxygen will have a total of 8 protons, but 10 electrons, and a resultant excess -2 charge. Therefore, it will try to find an atom (or atoms) with 2 positive charges to offset this and create an ionic bond with a metal ion.

50. In fact, ALL members of its group will do this. Oxygen, silicon, selenium, and tellurium will ALL try to gain 2 electrons to fill their p-shells and they will ALL end up with an oxidation state of -2 when they lose these electrons.

51. Because of their similar chemical behaviors, these aforementioned elements are grouped together into a chemical family. All of these elements in the oxygen family are known as chalcogens.

52. Members of a chemical family have the same number of valence electrons and similar chemical behaviors in chemical reactions. Let's take a look at another chemical family known as the halogens.

53. Halogens are even more reactive than chalcogens because they are closer to filling their valence shell and will seek out electrons with even greater affinity. Halogens, as a result, are more reactive than chalcogens.

54. Since atoms get larger as you add valence shells, electrons get further-and-further from the nucleus in the larger members of a chemical family. Since the nucleus has less control of the electrons, non-metals decrease in reactivity as they get larger, since the nucleus can't pull as hard to grab free electrons as the smaller atoms can.

55. Regardless of the degree of reactivity, all halogen ions have a -1 charge and gain a single electron from whatever metal cation they decide to steal from. All halogens are also brightly colored and toxic in their elemental form. Fluorine is a green gas, bromine an orange liquid, and iodine an unstable purplish-gray solid. Each element is shown below.

56. Let's take a look at the alkali metal family (also pictured above) in the very first column of the table.

57. All of these elements have one outers shell elecron, ins pite of being different diameters. Lithium, sodium, and potassium will ALL try to rid themselves of the single electron that is preventing them from dropping back to the configuration of the noble gas on the prior row of the table.

58. Metals show the opposite trend of non-metals in that they become MORE reactive as they get larger, because as the distance between the nucleus and the electron orbits increases, the weaker pull of the protons in the nucleus make it easier for electrons to be pulled off by anions needing them.

59. Therefore, in the diagram below, potassium would be predicted to have the greatest reactivity. Notice the similarity in the valence shells of their respective Bohr diagrams.

60. Now let's take a moment to recap and review the basic information contained in the periodic table, along with the meaning behind each measurement. The table below provides interpretations of each component.

61. Frankly, a biology student only really needs to know the basic uses of the periodic table for smaller elements. It's not really necessary to be an expert on the more complex rules that arise once you get into transition metals and larger elements.

62. Biological molecules are made almost entirely of small elements with simpler rules (such as C,H,O,N,P,S). However, keep in mind when we discuss chemical bonding in the next section that the rules get far more complex outside of biochemistry.

63. Let's quickly review the properties of metals and non-metals again before closing this section of basics, because there are concepts that will be needed to fully understand how ionic bonding works.

64. Metals are elements to the left of the metalloid staircase. They lose valence electrons during chemical reactions, because their outer shells have fewer electrons than empty spaces and this is energetically more favorable. They drop back to the noble gas configuration of the last element on the prior row.

65. Non-metals are elements ot the right of the metalloid staircase. They gain valence electrons during chemical reactions, because their outer shells are nearly full, and it is more energetically favorable to acquire the extra electrons to fill the shells and attain the noble gas configuration.

66. Consequently, metals form positively charged cations and non-metals form negatively charged anions due to their particular imbalance of protons in the nucleus and electrons surrounding the atom.

67. With certain excepts for transition metals that make use of space combinations from both their outer s and d orbitals, metals commonly have oxidation states of +1, +2, or +3. A few transition metals can form +4, +5, +6, or even +7 cations.

68. Non-metals commonly have oxidation states of -1, -2, and -3, and (rarely) -4.

69. Metalloids include elements such as boron, silicon, germanium, and antimony. They occupy the 'staircase' on the right side of the periodic table. Depending upon the situation, they can behave like metals or non-metals.

70. Metalloids are used for semi-conductors in computer chips because they conduct both heat and electricity slowly, allowing necessary electrons to flow through circuit boards without frying the hard drive.

71. The final chart below gives some summative examples of periodic table interpretation for 3 elements.

72. Bohr diagrams are included. Also included is a Lewis dot structure for each element, which is a very useful depiction of the atomic symbol surrounded by ONLY the valence electrons.

73. It is MUCH quicker to figure out how two elements will form chemical bonds by using Lewis dot structures to figure out how electrons will redistribute.

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ANSWERS TO FUN WITH CHEMISTRY: 1. F 2. D 3. C 4. A 5. E 6. B

Chapter two

Chemical Bonding, Molecules, and Compounds

A) Chemical Bonding Basics

1.As previously mentioned in the last section, uncharged atoms have an equal number of protons and electrons.

2. Since electrons can only be distributed into shells in distinct numbers, this usually leaves the outer shell only partially occupied in uncharged atoms.

3. Chemical bonding occurs because the most energetically stable configuration is for atoms to have a full or empty valence shell. Metals try to empty the shell and drop to the prior outer shell, while non-metals try to completely fill the partially occupied shell.

4. Partially occupied valence shells require more energy for the atom to maintain, because electrons spin and create unstable and wobbly orbits if the shell is only partially occupied.

5. A good analogy for this is the wheels of a car. Anyone who has changed a tire knows that they had better get ALL the lugnuts back on the wheel if they want to stay on the road.

6. Missing lugnuts can cause the wheel to spin right off of the axle. It also isn’t going to get you very good gas mileage, seeing as how you’ll be laying a giant skidmark down the road until you pull over or flip your car. See example below with a van load full of filthy hippies. If they had jobs and weren't following Phish to every city, this never would've happened.

7. Likewise, partially empty outer shells are unstable. Nature always seeks out the lowest energy state, which means the atom has to find a way to correct the imbalance.

8. Chemical bonding is the way this imbalance is corrected.

9. Atoms bond because they want to dump extra outer shell electrons (metals do this) or they want to fill up their outer electron shell (non-metals do this).

10. Remember that inner shell electrons aren’t involved because those shells are stable.

11. The outer shell is an all or nothing proposition…kind of like you can’t be a little bit pregnant or a little bit dead. The shell is either empty or its full. While some transition metals will settle for a half-filled d-subshell, these are the exceptions and not the rule.

12. When two or more atoms bond together, the result is a molecule.

13. A compound is any two or more different atoms bonded together. All compounds are, by definition, also molecules. However, all molecules are not compounds.

14. All compounds are, by definition, also molecules. However, all molecules are not compounds.

15. For instance, CO2 and H2O are BOTH compounds and molecules, but O2 gas is just a molecule.

16. There are 3 basic types of true chemical bonds and a couple of situations that create false chemical bonds that still aid in stability.

17. Ionic bonds occur when metal atom(s) donate(s) its valence electron(s) to a non-metal atom or atoms. The metal ends up with as a positively charged cation while the non-metal becomes a negative anion.

18. Ionic bonds are held together by oppositely charged ions.

19. Covalent bonds occur when two or more non-metal atoms overlap valence shells and share electrons.

20. Like two neighbors sharing one fence, the extra electrons from the other atom(s) orbit both or all of the atoms and fulfill the requirement for a full valence shell for everyone.

21. Metallic bonds occur in pure metal samples or alloys when metal atoms play ‘hot potato’ with their valence electrons, creating a continually circling ‘sea of electrons’ that surround the metal ions. Metallic bonds are pretty much irrelevant in biology, since there are no cyborgs walking among us.

22. Hydrogen bonds, though false chemical bonds, occur when hydrogen atoms assume a very slight negative charge and stick to the negatively charged poles of other molecules, which are usually oxygen atoms.

23. Van Der Waal Forces are similar to hydrogen bonds, but involve interactions between other types of atoms.

B) Ionic Bonds

1.Ionic bonds are a bond between a metal (left of stair step) and a non-metal (right of stair step)

2. Metals try to get rid of electrons in chemical reactions, because it is easier for them to empty their valence shells of a few electrons than to try to attain a much larger number to fill the shell.

3. Metals attain a positive charge in cation form (electrons removed) based on how many valence electrons they need to give up. For instance, sodium will be a +1 cation, while magnesium will create a +2 cation.

4. Non-metals are insulators that try to take up electrons to fill up their shells. They gain negative charges in anion form. For instance, oxygen will gain 2 electrons to become a -2 anion.

5. It is useful to touch on the concept of electronegativity with regard to ionic bonds. Electronegativity, measured on the Pauling scale, has no units, but is a relative measure of how strongly atoms attract electrons.

6. Electronegativity is a property that describes how strongly electrons are attracted to an atom and measures from 0 to 4 on the Pauling scale.

7. Metals have low or very LOW values of electronegativity. As metals grow larger in size, this value drops, since the electrons get further and further from the nucleus and the atom doesn’t want them anyway.

8. Think of metals as people trying to carry bags of groceries.

9. It isn’t hard to hold onto 2 or 3 bags of Cheetohs and hot dogs, but if you attempt to haul 16 bags of food into the house after you shopped hungry, catastrophe awaits. Busted eggs and spilled milk are sure to follow.

10. Non-metals have moderate to very HIGH values of electronegativity, because they jump on electrons with the ferocity of a hipster finding a mustache comb hidden underneath an espresso maker.

11. Let’s look at a couple of examples from the electronegativity chart below that would explain a couple of ionic compounds that would be predicted to form from the chart.

12. Let’s consider the case of Magnesium Oxide (MgO). Oxygen has a Pauling value of 3.5, while magnesium has an electronegativity of only 1.2. Oxygen would be predicted to rip the electrons off of magnesium’s outer shell.

13. This is, in fact, what happens. Magnesium oxide has a strong ionic bond.

14. Suppose a calcium atom had the opportunity to bond with either chlorine or with sulfur. In this case, it would be predicted that chlorine would win the battle for the electrons, since 3.0 is greater than 2.5.

15. The least electronegative metals are at the bottom left of the table, making them the most reactive, since they lose electrons most easily.

16. The most reactive non-metals are at the top right of the periodic table, since they can grab electrons with the most force, since the protons at the nucleus are the closest to the shells.

17. After the cation and anion form, there is an attraction between opposite charges that creates the bond.

18. Ionically bonded compounds are called salts. When they dissociate in water (most but not all do) they form an electrolyte, as (+) and (-) charges distribute around water molecules.

19. Ionic bonds are very strong. Ionic compounds have very high melting, vaporization points.

20. The average ionic bond is about 100 times stronger than the average covalent bond and 10,000 times stronger than the average hydrogen bond.

21. Below are examples of how ionic bonds are formed in two salts.

22. In the first example, there is a straightforward exchange. Sodium rids itself of one electron to have an empty 3rd electron shell, while chlorine needs one electron to fill the shell with 8 electrons to get to a stable Noble Gas configuration like neon. Sodium (+1) attracts chlorine (-1) and sodium chloride forms.

23. In the second example, the Iron atoms have 3 electrons to give up, while oxygen can only take 2. Therefore, it takes 3 oxygen atoms to bond to 2 iron atoms, since the least common multiple of electrons is 6 (see diagram).

24. Iron oxide would have a formula of Fe­2O3 since 6 positive charges attract 6 negative charges.

25. So what importance do ionic bond have in biology? Any compound that is a salt is ionically bonded. Therefore, the salts that dissolve to create electrolytes are ionically bonded, as are the salts in bones and shells.

26. Numerous body systems and cells require salt ions to run properly.

27. For instance, nerves need sodium, potassium, and calcium ions to function properly, which are all attained from salts in the diet. The heart, kidneys, brain, digestive tract, and numerous other organs also depend on salt ions to function properly.

28. With that said, ionic compounds are not as important as the covalent compounds we will cover next.

29. The chart below summarizes just a few of the places ionic bonds are found in living systems.

C) Covalent Bonds

1.Covalent bonds occur when valence shells are joined and electron pairs are shared between two or more non-metal atoms. These electrons then orbit both atoms and keep both shells filled.

2. This is much like sharing a common fence with your neighbor. Both of you benefit. He doesn’t want to see the sun beam off your pasty white spare tire belly in your pool and you don’t want to know that every day after work, he and his wife jump on matching pogo sticks while wearing gas masks, chef’s hats and cowboy boots.

3. Nobody takes anything from anybody, so there are no charges on the atoms in covalent bonds. No cations or anions come about as a result of the bond.

4. Most combinations of atoms in covalent bonds follow the octet rule. This is particularly true of the small atoms that make up biological molecules.

5. The octet rule: Non-metal atoms, with the exception of hydrogen, will try to attain a surrounding of 8 electrons in their full outer shell in order to get to the next Noble Gas configuration.

6. Noble gases are non-reactive and very stable, because they have 8 electrons in a full outer shell.

7. This is why neon, krypton, and argon can be heated to extremely high temperatures in neon signs without exploding the glass gas tubes. It is also why you can breathe helium with no ill effects.

8. The octet rule especially applies to atoms on the 2nd and 3rd row and is important to determine covalent bonding in organic chemistry and biochemistry. The octet rule pretty much applies to all biochemical compounds you’ll encounter in this course.

9. Hydrogen is the one major exception. Hydrogen only needs to get one extra electron to fill its shell, since it only has the small 1s sub-shell that holds a single pair of electrons. It wants to get to the helium configuration.

10. Covalent compounds usually have relatively low melting and vaporization points, since the forces that hold them together are shared pairs of electrons, rather than strong differences in electronegativity.

11. Covalent bonds, on average, are about 100 times weaker than ionic bonds, but much stronger than the false hydrogen bonds we will cover shortly.

12. Now let’s get to the visual representation of covalent bonds. It is much easier to use Lewis dot structures to figure out how two atoms will bond covalently, rather than drawing the entire Bohr diagram.

13. Inner shell electrons don’t bond, nor does the nucleus matter. These are redundant and waste time to draw.

14. Lewis dot structures simply show the chemical symbol, along with the outer shell electrons. From there, it is like working a Sudoku puzzle to figure out how to bond the atoms together to fulfil the octet rule.

15. Lewis dot diagrams use dots for unbonded electrons. When a bond forms, a pair of shared electrons is represented as a straight line. For example C-C shows that two carbon atoms are bonded together.

16. Two straight lines (O=O) is a double bond, three is a triple.

17. In drawing Lewis dot structures, you need to count up the dots and dashes and make sure that you get a sum of 2 around hydrogen atoms and a sum of 8 around everything else. The structure is wrong if you don’t get this.

18. Let’s look at two simple examples of water and ammonia first.

19. As you know, hydrogen atoms need only one electron to get to the stable helium configuration and fill the s-sublevel with 2 electrons.

20. In the case of water, oxygen needs two electrons to get to the neon configuration with 8 electrons in the 2nd shell. Oxygen can share an electron pair with two hydrogen atoms and everyone is happy.

21. Count around the oxygen and you will see 8 total electrons, while each hydrogen has two.

22. A similar situation occurs in ammonia, except nitrogen needs to go from 5 electrons to 8, so it must pick up a 3rd hydrogen atom to fulfil the octet rule.

23. Now let’s look at some more complex cases and see how to figure out their bonding pattern.

24. When you are given an assortment of atoms, the best strategy is to start with the atom that can bond to the most other atoms and go backwards, bonding hydrogen last.

25. Ethanol’s molecular formula is C2H6O, but this isn’t helpful in figuring out how to put the molecule together, since several other molecules can also have this formula.

26. Instead, it is written as CH3CH2OH, which provides clues as how to assemble the molecule.

27. Since carbon has 4 empty spots in the shell of 8 that can form bonds, you would start by chaining the two carbons together. Each carbon will still need to form 3 more bonds.

28. From there, oxygen has the next most open spots to bond, with 2. One of the two carbons can bond the oxygen, pushing a carbon up to 6 occupied slots.

29. Oxygen will still need to pick up another bond, as it now has 7 electrons surrounding it, with one shared pair and 5 electrons of its own remaining.

30. The entire problem can be fixed by surrounding the open slots with hydrogen atoms until the carbon and oxygen get to the magic number of 8. Each hydrogen will be satisfied sharing one pair.

31. Similar rules apply to the urea molecule below. Start with the carbon atom and attach the two nitrogen atoms first. Oxygen needs to come next.

32. You will notice when you add the oxygen that carbon and oxygen will both come up short with only 6 of 8 spots in the shell occupied. The solution is to share a second pair. Both will then get to 8.

33. From there, fill in with hydrogen atoms on the edges and everything works out.

34. As in the example of urea, double bonds are relatively common, since many oxygen compounds contain them.

35. Triple bonds are less common, as usually only nitrogen can make them. An example of a triple-bonded molecule is the N2 gas that makes up 78% of the air you breathe.

36. Double and triple bonds are higher in energy & harder to break. Carbon-oxygen double bonds yield more energy when they are metabolized by your cells than carbon-hydrogen bonds.

37.