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AP Biology - Unit 1 Chemistry Fundamentals and Organic Chemistry

AP Biology - Unit 1 Chemistry Fundamentals and Organic Chemistry. Campbell Biology 9 th Ed. Chapters 2, 3, 4 & 5. Biochem Unit Objectives (Chap. 2).

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AP Biology - Unit 1 Chemistry Fundamentals and Organic Chemistry

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  1. AP Biology - Unit 1Chemistry Fundamentalsand Organic Chemistry Campbell Biology 9th Ed. Chapters 2, 3, 4 & 5

  2. Biochem Unit Objectives (Chap. 2) 1. Describe the structure of an atom. State which atomic feature determines the identity of the element. Compare & contrast atoms vs. ions. 2. Explain covalent bonding.

  3. Biochem Unit Objectives (Chap. 3)  3. Distinguish between polar and nonpolar molecules. Give examples. Relate the polarity of a molecule to its behavior in various solutions. 4. Identify the characteristic behavior of water molecules and discuss their significance to living organisms (and the suitability of earth to sustain life) 5. Define acid and base. Explain the relationship between the pH scale and the concentration of hydrogen ions.

  4. Biochem Unit Objectives (Chap. 4) • Explain how carbon’s bonding arrangement allows it to form a variety of structures Note: • Number and orientation of Covalent bonds • Single, double and triple bonds • Ring structures • Identify the characteristic functional groups of organic molecules. Relate each functional group to the type of organic molecules possessing it.

  5. Biochem Unit Objectives (Chap. 5) 8. Distinguish between carbohydrates, lipids, proteins, and nucleic acids. 9. Match monomers with corresponding polymers. Explain the chemical processes which form and break apart polymers. 10. Distinguish between polypeptides and proteins. Explain the four structural levels of proteins. 11. Describe the structure and explain the function of nucleic acids. Distinguish between DNA and RNA.

  6. Atomic Structure • Atoms are the smallest particles of elements, but they consist of even smaller “subatomic” particles; Protons Neutrons and Electrons • The number of protons determines the element

  7. Subatomic Particles Note the significance of charge. Charged particles will exert forces of attraction and repulsion upon each other. Opposite charges attract. Similar charges repel.

  8. Isotopes and the Significance of Neutrons • Protons in the nucleus are positively charged, so they will repel each other • The only way to keep the nucleus from exploding is to keep protons separated, which is accomplished by Neutrons https://www.youtube.com/watch?v=SeDaOigLBTU

  9. Isotopes and the Significance of Neutrons • The number of neutrons may vary for atoms of a particular element. • This does not change the identity of the element, but it will affect the stability of the nucleus. Some combinations work better than others • The different combinations are called Isotopes https://www.youtube.com/watch?v=SeDaOigLBTU

  10. Atoms vs. Ions The number of electrons may vary. We generally think of the number of electrons as always being equal to the number of protons, but that isn’t always true. When the number of protons is equal to the number of electrons, the particle is neutral and will be called an Atom. When they aren’t equal, the particle is charged and is called an Ion https://www.youtube.com/watch?v=WWc3k2723IM

  11. Electron Configurations Because electrons repel each other, they tend to spread out in the space surrounding the nucleus. At greater distances from the nucleus there is more space available. This allows for both more movement and a greater number of electrons. Electrons closer to the nucleus are said to be at “lower energy levels”. Electrons further from the nucleus are at “higher energy levels” http://www.middleschoolchemistry.com/multimedia/chapter4/lesson6

  12. Orbitals • Regions in space where pairs of electrons are likely are called Orbitals. • Higher energy levels, being larger, contain greater numbers of orbitals • The pattern of orbitals filling is very distinct, and described mathematically In the diagram to the left we see the shapes of the different orbital types. These figures are derived from algebraic equations

  13. More Orbital Stuff • We describe orbitals by type (“s, p, d, and f”) • Each higher energy level contains one additional type of orbital • Each additional type has space for more electron pairs (following a pattern of the next odd number; 1, 3, 5, 7 . . .)

  14. Electron Configurations, Revisited • So the first energy level has only 1 orbital (an “s” orbital). That means the first energy level only has room for 1 pair of electrons • All the other energy levels also have an “s” orbital, but also 3 “p” orbitals • Those 4 orbitals have room for 4 pairs of electrons (an “octet”) Note that each higher energy level has one more type of orbital than the level before it. Each new type also has a greater number of orbitals: s = 1; p = 3, d = 5, f = 7

  15. Orbitals and the Periodic Table • Note that the periodic table is organized to demonstrate electron configuration • Electron configuration can be used to predict the chemical behavior of elements

  16. Valence Electrons • Because electrons are attracted to the nucleus, they tend to be close to it, at lower energy levels • Atoms with larger numbers of electrons will need to have some electrons at higher levels as well • The highest energy level used is called the Valence level • It is the valence electrons that will form chemical bonds In this Mg atom, the first 2 levels are full. There are 2 electrons in the 3rd level. Those 2 electrons are the valence electrons. When the atom reacts, those 2 valence electrons are lost, leaving behind only the electrons in the first 2 levels (which are full to capacity).

  17. Valence Electrons & Lewis Structure • Lewis Structures are electron dot diagrams that show the valence electrons only • Note that the diagrams show the electrons in pairs, because orbitals are electron pairs • There are spaces for 4 pairs because there is 1 s orbital and 3 p orbitals It is the unpaired electrons that will tend to be involved in chemical reactions. They may be lost, or the space may be filled by gaining or sharing electrons

  18. Valence Electrons & Lewis Structure • Note that Ne and Ar have all 4 orbitals completed. The valence level is full to capacity so they are chemically stable, no reaction will occur • Li and Na each have only 1 valence electron. It will usually be lost, forming a positively charged ion Examine each element in this chart and predict the number of chemical bonds that will be formed. Also predict whether electrons will be gained, lost or shared.

  19. Ionic Bonding • Ionic bonds happen between ions (not atoms) • In this example, Na has lost an electron (+ ion) which Cl has gained (- ion) • Oppositely charged ions attract each other Look at the Lewis structures shown for these 2 elements. Na had 1 valence electron. When that electron was lost, it left only the full energy levels behind. The Cl atom had 3 of its 4 orbitals completed. Only 1 orbital contained an unpaired electron. Gaining an electron filled that orbital and created a pair.

  20. Covalent Bonding • Covalent bonds are formed when an electron pair is attracted by the nuclei of 2 different atoms simultaneously • They form when 2 atoms each have an unpaired electron • The bond creates a pair, filling an orbital for both atoms

  21. Electronegativity • Electronegativity is a number value that describes the relative strength of attraction that an atom will have on a valence electron • This helps us to predict what will happen when 2 atoms react

  22. Electronegativity • For example, if 2 atoms of H collide, each of them will attract a valence electron from the other atom with a relative force of 2.1 • Since the attraction force is equal for both atoms, the electron pair will be shared equally

  23. Electronegativity • If an atom of Na collides with an atom of Cl, the electron pair will be attracted by Na with a force of .9, but by Cl with a force of 3.0. • Cl wins. It will take the electron, ionize both, and form an ionic bond As a rule, if the difference between the 2 atoms is 1.7 or greater the bond will be Ionic. If the difference is less than 1.7 the bond will be Covalent

  24. Electonegativity Difference • In this example, both atoms have an unpaired valence electron. When they collide a chemical bond will form • H has an electronegativity of 2.1 and C has an electronegativity of 2.5. • Neither can exert a force strong enough to ionize the other. The difference is .4, far below the 1.7 needed to ionize, but not zero. • The bond will be covalent, but will have a slight polarity in favor of the C atom, represented by the “dipole” shown The H has a partial + charge, the C has a partial - charge

  25. Nonpolar vs. Polar Molecules • Shape matters. • Both of the molecules shown have 3 atoms forming 2 covalent bonds • Both of the molecules have a “dipole moment”, meaning the bonds have a polarity • The bent shape of the first molecule concentrates the negative charge at one end, but the linear shape of the second molecule puts the polarities directly opposite each other Dipole moment is a vector quantity, it has a magnitude and a direction https://www.youtube.com/watch?v=PVL24HAesnc

  26. Try these 2 examples: • Use the following electonegativities: • H = 2.1, O = 3.5, C = 2.5 • Determine the bond polarity for each bond • Label each atom as d+ or d – • Draw the dipole moment for each bond • Predict the polarity of the whole molecule based on shape and bond dipoles

  27. Water Molecules • Water is special for a number of reasons • The bond between H and O is extremely polar. It’s right on the fringe between covalent and ionic • The oxygen atom has 2 pairs of unbonded electrons, resulting in both a bent shape and a concentration of negative charge on the oxygen atom • The hydrogen has only 1 electron, and it is in the bond between the H and O, leaving the H nucleus exposed at the end of the molecule https://www.youtube.com/watch?v=HVT3Y3_gHGg

  28. Water Molecules • In the diagrams given, recognize the bent shape, unbonded pairs of electrons, and dipole moment • Construct several water molecules with your modeling kit and predict how they will behave towards each other

  29. Hydrophilic vs Hydrophobic Hydrophilic = Water Loving Hydrophobic = Water Hating Substances that are hydrophilic are attracted to water, and usually dissolved by water Hydrophobic substances have no attraction to water and will not dissolve https://www.youtube.com/watch?v=i3jA40arq9Y

  30. Characteristics & Behavior of Water • Water molecules have strong attractions to each other and to other polar substances • Water has a high “specific heat”, and resists temperature change • Water absorbs significant amounts of heat energy when it evaporates • Water expands when it freezes • Water is an excellent solvent of polar and ionic solutes • Acid/Base chemistry is dependent upon water https://www.youtube.com/watch?v=iOOvX0jmhJ4

  31. Adhesion and Cohesion Adhesion and Cohesion both refer to things sticking together. Adhesion Water sticks to something other than water Cohesion Water sticks to itself

  32. Surface Tension Strong cohesion between water molecules at the surface can create unusual effects. Note how the water beads up on the surface of the penny. The cohesion forces are stronger than the force of gravity, so the water doesn’t fall off and drop to the table. The insect shown should theoretically sink, but the surface tension is strong enough to support its weight. Watch the videos linked below. You’ll see some interesting effects. https://www.youtube.com/watch?v=ntQ7qGilqZE https://www.youtube.com/watch?v=ynk4vJa-VaQ

  33. Heat Capacity • Compared to other substances, water requires a huge amount of energy to change its temperature. • That means that water tends to hold a steady temperature. Consider a rain forest and a desert. Both are hot, but one holds its heat and the other varies temperature drastically from day to night • It also means that water can store and move energy very well. The oceans are a huge heat sink https://www.youtube.com/watch?v=dQk5yi05PIk

  34. Heat Capacity and World Climate • Consider Los Angeles and Phoenix. Almost identical in latitude, but LA has a more moderate climate because the ocean stabilizes the temperature • How about Ireland. Further north than Saskatoon, but with a moderate climate. The gulf stream carries warm water from the equator to the European coast and brings heat energy to the British Isles

  35. Evaporation and Heat Transfer • Evaporation is an endothermic reaction. It requires the input of a significant amount of heat energy • Sweating doesn’t cool you off. It’s the evaporation of the sweat. As it evaporates it draws heat energy from your body • MJ’s sweat isn’t really orange https://www.youtube.com/watch?v=vLfWnX0ahtc

  36. Freezing, Expansion, and Floating The shape and polarity of water molecules cause them to form very distinct arrangements when they crystallize. Note the difference between the molecular arrangement of liquid water and ice. In ice the molecules are closer together, but the crystals contain empty spaces in the hollow part of the molecular bend. Water is the only substance that expands when it freezes. That expansion makes ice less dense than liquid water, causing it to float https://www.youtube.com/watch?v=bzTZx1RDV3w

  37. Solvent Properties • The polarity of water creates a distinct + and – end of the molecule. Those charged regions will pull apart crystals of polar and ionic substances, causing them to dissolve. • Ionic substances will not only dissolve, the ions will “dissociate” from each other As the salt dissolves, the Na+ dissociates from the Cl-. Note the orientation of the water molecules relative to the ions in the hydration shells https://www.youtube.com/watch?v=0cPFx0wFuVs

  38. Dissociation, Acids and Bases • The covalent bonds between H and O are so polar they are almost ionic. • When water molecules collide, especially when the H from one collides with the O from the other, the H will ionize forming H+ and OH- https://www.youtube.com/watch?v=6gjZ88JbJas

  39. Dissociation, Acids and Bases • The H+ ion is really a free proton • H+ ions form Acids • OH- ions form Bases • In pure water there is constant activity of the molecules. Some dissociate while others form back in to H2O On the average, every liter of water has 10-7 moles dissociated into H+ and OH- ions https://www.youtube.com/watch?v=6gjZ88JbJas

  40. pH Scale and H+ ion concentration • Remember that H+ ions make acids, so the greater the concentration of H+ ions, the stronger the acid • In pure water, the concentration of H+ ions is 10-7mol/L • The p in pH is a mathematical function (-log). If I look at 10-7 and lose the base 10 and the – sign, I end up with pH = 7 https://www.youtube.com/watch?v=LS67vS10O5Y

  41. Organic Compounds • The word Organic comes from the word Organism • Organic compounds are mostly compounds that are formed by living organisms • Carbohydrates, Lipids, Proteins and DNA are all organic compounds, but there are others • All organic compounds are structured around Carbon

  42. Organic Compounds • The diagram to the right shows an apparatus used for a famous experiment to determine whether organic compounds could be produced from simple inorganic substances without living organisms to direct the reaction • Watch the video and read the text • Campbell Biology p58-59 https://www.youtube.com/watch?v=gWqJfBEzU98

  43. Carbon is a unique element • Carbon has 4 valence electrons, all unpaired, which allows it to form 4 covalent bonds all equally spaced around the carbon in 3 dimensions • It bonds readily with H, N, O, S, P and with other C • It can form single, double and triple bonds, and ring structures • It can form long, stable chains when bonded with other C’s https://www.youtube.com/watch?v=QnQe0xW_JY4

  44. Carbon’s Bonding Arrangements • Using your molecular modeling kit, construct 2 carbon hydrocarbon molecules with • Only single bonds • One double bond • One triple bond • Note: • The number of hydrogens • The orientation of the bonds • Can the C-C bond rotate freely?

  45. Hydrocarbons • The simplest organic compounds are saturated hydrocarbons. Only H and C, with only single bonds • There is no limit to the number of carbons in the chain • The chart to the left shows names, formulas and structures of some simple “alkanes” https://www.youtube.com/watch?v=UloIw7dhnlQ

  46. Substituted Hydrocarbons • Because Carbon can form covalent bonds with all the other nonmetals, it can easily bond to something other than H at a given position • In the Ethanol shown, an oxygen atom has been substituted on one of the carbon atoms in Ethane Any nonmetal atom or group of nonmetal atoms can be substituted an any position on the chain. The substitution may change the shape, polarity, and behavior of the resulting molecule

  47. Isomers • Using your molecular modeling kit, construct • A 2 carbon saturated hydrocarbon with a single substitution • A 2 carbon unsaturated hydrocarbon with a single substitution • A 2 carbon unsaturated hydrocarbon with 2 different substitutions • A 3 carbon saturated hydrocarbon with a single substitution • A 4 carbon saturated hydrocarbon with no substitutions https://www.youtube.com/watch?v=22PkbCu3vYI

  48. More complex substitutions • Using your molecular modeling kit, construct: • A 2 carbon saturated hydrocarbon (“ethane”) • The simplest molecule you can construct using only • Oxygen and hydrogen • Nitrogen and hydrogen • Carbon and oxygen • Identify the resulting simple molecules • Substitute each onto your ethane molecule (one at a time, please)

  49. Functional Groups • Each of these functional groups is a characteristic structure that is commonly found in organic molecules • Each is fundamentally a simple, inorganic molecule substituted onto the hydrocarbon chain • Each will change the shape and polarity of the molecule, causing characteristic chemical behaviors https://www.youtube.com/watch?v=OGD3q1eQ1TE

  50. Monomers and Polymers • Mono = 1, Poly = many • Monomers are simple organic molecules • Polymers are larger, more complex molecules that are formed by attaching many monomers together in a chain https://www.youtube.com/watch?v=VigpwmH7E3M

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