ex. Line Bond Structure like Lewis Dot but using lines not dots (and no lone pairs shown)

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1 hemical Formulas: 1. Empirical smallest whole number ratio (ex. 2 ) 2. Molecular actual number of atoms in the specified ratio (ex ) 3. Structural shows order of attachment of all atoms ex. Lewis Dot Structure N ex. Line Bond Structure like Lewis Dot but using lines not dots (and no lone pairs shown) N 4. ondensed Structures (line bond structure w/o lines) an use a line when connecting a functional group. onvert the following into a condensed structure: O Or you can use parenthesis: on the left ( 3 ) on the right 3 2 ( 3 ) 2 in the middle - 3 (O) 2 3 onsider the following and draw the line bond structure: 3 (O) 3

2 5. Skeletal (Zig-Zag) Structures (line bond w/o letters) Rules for Skeletal Structures a. Do not draw in (as in - or -) b. At the end of a line, assume there is a c. At the intersection of two or more lines, assume there is a d. s attached to are not drawn in (recall that can make 4 bonds ) e. Everything besides - and - must be drawn. Must be able to convert from skeletal to condensed and vice versa O onvert to skeletal structure: ( 3 ) 2 (Br) 2 ( 3 ) 3 Now: convert to condensed formula: O Last tidbit on structures: Drawing tetrahedral carbons in three-dimensions The wedge/bold line bonds coming at you The dashed line bonds going away from you In what direction are the O s facing? O O O O Bonding in Molecules atoms coming together to form molecules If atoms are too close, their positively charged nuclei repel each other If atoms are too far apart, no overlap of electrons can occur to form bonds

3 Ideal Distance between two nuclei = Bond Length Orbitals regions of space where electrons are found The shape of an orbital is described by a math equation and is determined by probability of where electrons are most likely to be found around the nucleus. Like math equations (think of a sine or cosine wave), the orbitals will also have positive and negative areas (when thought of on a graph, shown in black and/or white below). Every orbital holds 1 pair of electrons Recall : s p (p x, p y, p z ) Each atomic orbital can hold up to two electrons. Atomic orbitals (s, p, and others) combine to form molecular orbitals. Once combined, they are no longer called ATOMI orbitals (NOT s or p or any hybrid orbital any longer!) The number of atomic orbitals, before combining, must equal the number of molecular orbitals after bond formation has occurred. Each molecular orbital still only holds 1 pair of electrons. Molecular orbitals are either o BONDING molecular orbitals (those that make bonds when they overlap) or o NON-BONDING or ANTIBONDING molecular orbitals (those that do not). Bonding Orbitals overlap positively in what is called an Additive fashion when the orbitals are mathematically both positive or both negative (in the same phase) and are very stable because these orbitals will place the negatively charged electron density (shown in gray below) between the positively charged nuclei. Relatively speaking, where are the nuclei? (Look at all that gray electron density in the middle!) Very Stable! Antibonding orbitals overlap negatively in what is called a Subtractive fashion when the orbitals are out of phase (one positive and one negative ) and are unstable because

4 there isn t any negatively charged electron density between the positively charged nuclei. There s no actual overlap occurring! Where are the nuclei? See how the nuclei are exposed and facing each other with very little electron density to shield the positive charges from each other? IG ENERGY!! Types of bonds that we commonly use in Organic hemistry: (formed with bonding molecular orbitals, during in-phase, positive overlap) Sigma Bonds formed whenever two spherical orbitals come together Pi Bonds only formed only when two p orbitals overlap sideways. Most atoms do not form bonds using just s or p orbitals. Most atoms use YBRID orbitals. Why? Reasons for ybridization 1. Allows atoms to have the correct number of unpaired electrons to make necessary number of bonds. 2. Best spatial orientation achieved least amount of electron repulsions 3. Stronger, more stable bonds as a result of overlapping orbitals with greater s character To determine hybridization of most atoms count the number of SIGMA bonds and LONE PAIRS of electrons

5 Bonds: single bond = 1 sigma bond double bond = 1 sigma bond, 1 pi bond triple bond = 1 sigma bond, 2 pi bond Lone Pairs (neutral atoms, like O, N): Group 5 = one lone pair Group 6 = two lone pair Find a review of ybridization in detail in the problem set for Exam 1 material and view as either a powerpoint or pdf file.