1 Chapter 1 Summary –Summer 2014 Homework: 25, 26 (assume C-H bonds are non-polar), 27-29, 31, 33 a, b, c, and d, 34 a and b, 35, 36, 37 (Don’t worry about drawing actual orbitals. Just name the orbitals used to make each bond. Then look at solution set to be sure picture of orbitals makes sense), 38 (you should be able to tell what the geometry is without Vollhardt telling you!), 39-44, 46, 49(You should be able to do this challenging problem without referring back to #38 as suggested by Vollhardt), 51(often more than one answer is acceptable, see solution). Note: This is all General Chemistry review material. Some of this material will NOT be directly discussed in lecture. However, you are responsible for understanding/knowing all material listed below regardless of whether I discuss it in lecture. Beginning: This is a broad introduction to some of the ways organic chemistry touches your life! 1.1: The Scope of Organic Chemistry: Read over this section for orientation purposes. Know that functional groups are particular arrangements of atoms that impart a particular and predictable type of chemistry to that portion of the molecule. See the front of the very last page of your textbook for a functional group list. You should be able to recognize all these functional groups and name them. o Note that alkanes are not ordinarily considered functional groups as they are generally not reactive under “normal” circumstances (except combustion and some chemistry we will see in chapter 3). o This general lack of reactivity allows alkanes to be the “backbone” of organic molecule structures. You should recognize the structural difference between a simple alkane and a cycloalkane. Don’t worry about the term “stereoisomer” at this point. You will have in depth analysis of this term in chapters 4 and 5. Know that a “chlorination” reaction replaces a hydrogen atom with a chlorine atom. Know that in a “substitution” reaction an atom (or group of atoms) is replaced by another atom (or group of atoms). Know that an “elimination” reaction produces a double bond by removing two groups/atoms on adjacent atoms. Know that organic “synthesis” is a general term used to describe the making of a particular compound from other compounds. Be comfortable with the terms reactant (substrate), products, reaction intermediates, and reaction mechanism. 1.2: Coulomb Forces: You should read this entire section carefully. It discusses important fundamentals. Understand that oppositely charged species attract and same charged species repel. Understand the components of Coulomb’s law in that the attractive force between oppositely charged species is related to the magnitude of their charges and the distance between them. The bigger the charges the more the attraction. The closer the species the more the attraction…to a point. Know that bonds are only formed when it is energetically favorable. The underlying driving force for bonding is that opposite charges attract (electrons, which have a negative charge, are simultaneously attracted to the nucleus, which have a positive charge, of two atoms). Know that bond strength is simply the attractive force that exists between two particular atoms. Know that bond length is the optimal distance between the two atoms so that the system reaches an energy minima (see Figure 1.1). o Pay particular attention to the simultaneous attractive and repulsive forces present between two atoms depicted in Figure 1.2. 1.3: Ionic and Covalent Bonds (Do in-text exercises 1, 3-5) Read this entire section carefully! Understand the differences and similarities between covalent and ionic bonding especially as depicted in Figures 1.2 and 1.3. Understand that you can relate the position of an element in the periodic table to whether the atom will tend to make covalent bonds or ionic bonds or both. While I will not ask you to do this on an exam, you should be able to reproduce Table 1.1 from memory. Knowing the position of these elements in the periodic table will help you figure things out throughout this course without having to refer to a hard-copy of the periodic table. ------------------ 2 Ionic bonds become possible when the atoms involved acquire a filled outer shell (octet) by transferring valence electron(s). The actual ionic bond is the attractive force that exists between the resulting oppositely charged ions. Know that the IP refers to the energy required to remove an electron from an atom and that EA refers to the energy involved in adding an electron to an atom. Then realize that the resulting ions will have a Coulombic attractive force. (See page 8). Be completely comfortable with the Lewis dot structures on the top of page 9. Be comfortable with the differences between the hydrogen atom, the proton (hydrogen cation), and the hydride ion (hydrogen anion). ----------------- In covalent bonding the atoms share electrons to achieve their respective noble gas configurations. Understand how periodic table position determines whether the elemental form of an atom will be monoatomic (like the noble gases), diatomic molecule with single bonds (like hydrogen or halogen molecules), double bonds (like some Lewis structure misrepresentations of molecular oxygen), or triple bonds (like nitrogen molecules). Understand how periodic table position determines how many bonds an atom will have in a molecule (Like stable form of carbon generally have 4 bonds, nitrogen generally has 3 bonds, and oxygen generally has 2 bonds). ------------------- Understand the origin of polar covalent bonds. Understand the trends of electronegativity in the periodic table. Understand that if atoms have the same electronegativity then they will form perfectly covalent bonds. However, the larger the electronegativity between two atoms the more polar the bond will be, and thus the bigger dipole that bond will have. Note that even though carbon and hydrogen don’t have the same electronegativity value the bonds between them are generally considered non-polar. Be very comfortable with partial positive (δ+) and partial negative charges (δ-). Be able to accurately place these symbols on the Lewis Structures of molecules (See homework question #26). ------------------- If you remember VSEPR theory from general chemistry then you already understand how electron repulsion determines the shape of molecules. You should be able to predict if a molecule is going to be linear, trigonal planar, trigonal pyramidal, bent, or tetrahedral. You should know that the “repulsing ability” of a lone pair of electrons is stronger than the “repulsing ability” of a bonding pair. Why? Think about how highly localized bonding pairs are relative to nonbonding electron pairs. 1.4: Lewis Structures (Do exercises 6-8) Depending on how comfortable you are generating Lewis structures you will be able to skim this section or you will have to spend an hour or two going over it. This section goes over the rules for creating Lewis structures. o Note: I think it is perfectly o.k. to use dashed lines to represent non-bonding pairs (lone pairs) of electrons as long as you make sure they won’t be misinterpreted as a negative charge symbol. You should make sure you understand the exceptions to the octet rule. You should be an expert at determining formal charges. 1.5: Resonance Forms (Do exercises 9a, 10, 11) READ THIS ENTIRE SECTION CAREFULLY! Typically the general chemistry treatment of resonance is severely lacking. Consequently prepare to work hard at this section. Resonance is a crucial part of this course that you need to take very seriously. Resonance is very useful when our method for representing molecules with a single Lewis structure is incorrect or misleading. Know that the “real structure” of a molecule is a composite (or hybrid) of all its resonance structures. Also realize that the “accepted structure” we ordinarily choose to use as a representation of a particular molecule is just another resonance structure. The real structure is a composite of all the resonance structures. 3 The degree to which a particular resonance structure contributes to the “real structure” is proportional to its stability relative to the other resonance structures. If all resonance structures are equivalent then all resonance structures equally contribute to the real structure. If you are showing that two Lewis structures are resonance forms use a double tipped arrow (↔) Remember that the position of atoms in resonance structures NEVER move (relative to the other atoms in the molecule). The only difference between resonance structures is how the electrons are arranged. Carefully read and commit to memory all the guidelines listed in your text book (page 21) for picking out the “best” resonance forms for a particular molecule. Remember the “best” is the most stable form and consequently contributes most significantly to the “real structure”. 1.6: Atomic Orbitals (Do exercise 11) This section should be an “easy” review based upon your general chemistry training. Don’t get bogged down in the subsection “the electron is described as a wave”. Just appreciate that electrons are not point charges. You should know that atomic orbitals have a characteristic shape and energy. You should know the shape of s and p orbitals. You should be very comfortable with the Pauli Exclusion Principle, Aufbau Principle (building-up order), Hund’s Rule, electron spin, and pairing. You should be very comfortable with Figures 1.8 and 1.9 on page 27 and Figure 1.10 on page 28. 1.7: Molecular Orbitals and Covalent Bonding. (Do exercise 14). Read this section carefully. This section shows you how atomic orbitals overlap to form molecular orbitals. You should be comfortable with the following: bonding and antibonding orbitals, molecular orbital diagram construction for simple diatomic molecules, and energy splitting. You should know the difference between sigma (σ) and “sigma star” (σ*) molecular orbitals. You should know that the attraction between two atoms is determined by the number of electrons in bonding molecular orbitals vs. anti-bonding molecular orbitals. 1.8: Hybrid Orbitals (Do exercise 16) Based upon your general chemistry training this section should be review. However, I have found that students still need to spend a lot of time studying this section. Pay particular attention to how hybridization determines the molecular shape (and bond-angles) around a particular atom. Also see how hybridization is connected to bond multiplicity. You should be comfortable with sp, sp2, and sp3 hybridization. Atom hybridization involves the mixing of atomic orbitals to make an equal number of hybrid orbitals. You should be able to tell what atomic orbitals were mixed to form a particular type of hybrid (like sp 3 hybridization means 1 s atomic orbital was mixed with 3 p orbitals). The hybrid orbitals on the atom adopt a relative orientation that maximizes the distances between the hybrid orbitals. Lone pair electrons will nearly always reside in hybrid orbitals. The un-hybridized orbitals (in this class they will always be pure p orbitals) are generally used to make multiple bonds. o Note: In rare cases lone pairs will occupy a p orbital if that allows the molecule to become aromatic However, you don’t have to worry too much about this exception until chapter 14 and 15 in Organic Chemistry II. Also lone pairs will occupy pure p orbitals if they can become part of a conjugated system which you don’t have to worry about until chapter 14 of Organic II. I want you to learn hybridization like it is universal law (because it will help you immensely) but I also want you to appreciate that hybridization is just a model…not a hard-fact of nature. Read over the solution to exercise (#17) to appreciate how concepts such as bonding are often over-simplified and that there are often exceptions and there are always grey areas! However, learning these simplified principles can be very helpful for learning chemistry. Thus, assume in this course that the oxygen of water is sp3 hybridized. 4 1.9: Structures and Formulas (Do exercise 19, 20, 21, and 22). This section is more “busy work” than hard mental work but it is very important work! You should know the term structural/constitutional isomer. You should have a methodical procedure for determining all the constitutional isomers for a particular molecular formula. However, I think it is best to wait to develop your procedure until after section 5 of chapter 2. I will outline a procedure in more detail after 2.5. At that time I will also assign homework question 45 from chapter 1. You should be able to convert any Kekule/Lewis or condensed formula to a bond-line representation quickly and visa-versa. You need to get very comfortable with the “bond-line” formula method for representing organic molecules since this is how almost all organic chemists communicate. You also need to be very comfortable with dash-wedge representations. 5 Lewis Dot Formulas 1. Calculate the number of valence electrons. a) For neutral molecules: # Valence electrons = all _ atom _ group _ numbers b) For polyatomic ions (anions or cations) # Valence electrons = all _ atom _ group _ numbers - total _ ch arg e _ on _ ion 2. Construct Skeleton a) Always keep in mind the number of covalent bonds an atom tends to make. b) The most electropositive atom is placed in the molecule center. An important exception is that H is always placed on the perimeter. c) The relatively electronegative atoms are placed around the central atom(s). Alternatively, they are placed between the central atom and a perimeter H). 3. Atom connection and electron distribution. a) Connect all atoms with a single pair of electrons. b) Add electron pairs to perimeter atom as needed to complete octets. The left over electron pairs are added to the central atom(s). c) Check the central atom's octet situation. If it does not have an octet then start using lone pairs on perimeter atoms to make multiple bonds until octet is complete. 4. Calculate Formal Charges a) Atom F.C = Valence electrons on free atom - (# bonds to atom) - (TOTAL #lone pair electrons around atom). Thus, FC = Group # - # bonds – number of lone electrons 5. Evaluate Resonance Structures when applicable.
