1559T_ch01_01-1710/22/051:48Page3 ⊕ EQA Keys to the Chapter·3 chemis ALWAYS TRUE (these are your reminders): 1.Individual contributing resonance formsdoexist.Only the resonance ybrid.which isa"weighted 2.Altrage"or nd of s have the same total number of valence electron and the same total charge their valence shell in any and oipty to rce forms geometry and set of atomic positions for the actual chemical substance being represented. USUALLY TRUE(these are mostly implied by the guidelines in the textbook-we're just spelling them out): differ ony in the positions ofand/or nonbonding elctos Thebond stay put. ance form into another h rs from places where there is an 3. ms witl are u ally more im t contributors than resonance 4.Resonance forms with fewer charged atomsare usually more important contributors. )and below(r. 1)are nor limited to octets 1-6,1-7,and 1-8.Orbitals Atomic orbitals are a convenient way to represent the distribution of electrons in atoms.Note that the and ced with the the ital an alterative to the Lewis electron-dot method for umber of m nbing a bond is a onding atomic o ls and will give rise to st Hybrid orbitals are d ved by mixing atomic wave functions.They are used to explain the geo several ac e ny by contributing to bonding.The participation of different numbers ofs and porbitals in the hybridization al. ange of bor angles.thereby permitting electron pairs to get as far away from each other as pos- Keep in mind some points for bookkeeping with hybridization.If an atom starts with ones orbital and three p orbitals,it will always end up with a total of four orbitals,no matter how they have hybridized for bonding contains 50%s and 50 character.whereas an sp2 hybrid iss andp in n ature fore hybric 1.sp hybridized atom (linear geometry):contains two sp orbitals (each one iss andp in character)and two ordinary p orbitals 2(s+D)+2D=1s+3 chemistry. It is somewhat abbreviated. Some additional considerations regarding resonance forms (and some reminders of the basic rules) follow. ALWAYS TRUE (these are your reminders): 1. Individual contributing resonance forms do not exist. Only the resonance hybrid, which is a “weighted average” or blend of the contributing forms, is real. 2. All resonance forms of a single chemical species must have the same total number of valence electrons and the same total charge. 3. Second-row atoms (i.e., up through neon) can never exceed an octet in their valence shell in any resonance form. In other words, the rules for drawing Lewis structures apply to drawing resonance forms. 4. Atom positions and geometries do not change from one resonance form to another—there is only one geometry and set of atomic positions for the actual chemical substance being represented. USUALLY TRUE (these are mostly implied by the guidelines in the textbook—we’re just spelling them out): 1. Resonance forms differ only in the positions of and/or nonbonding electrons. The  bond electrons normally stay put. 2. We convert one resonance form into another by moving electron pairs from places where there is an excess of electron density to places where there is an electron deficiency. 3. Resonance forms with the most covalent bonds are usually more important contributors than resonance forms with fewer covalent bonds (but don’t forget that octets take priority). 4. Resonance forms with fewer charged atoms are usually more important contributors. 5. Atoms in the third row (P, S, Cl) and below (Br, I) are not limited to octets of electrons in their outer shells. In fact, Lewis structures with 10 or 12 valence electrons are frequently written for these elements. 1-6, 1-7, and 1-8. Orbitals Atomic orbitals are a convenient way to represent the distribution of electrons in atoms. Note that the
and signs associated with parts of these orbitals do not refer to electrical charges. They refer to mathematical signs of functions (wave functions) associated with the distribution of the electrons. Molecular orbitals are similar but are spread out over more than one atom. They provide an alternative to the Lewis electron-dot method for picturing bonds. The number of molecular orbitals involved in describing a bond is always exactly equal to the number of atomic orbitals contributed by the individual atoms. Overlap of atomic orbitals results in bonding, antibonding, and sometimes also nonbonding molecular orbitals. Bonding orbitals are always lower in energy (more stable) than the original constituent atomic levels, and antibonding orbitals are always higher in energy. Thus bonding electrons will be more stable than electrons in nonbonding atomic orbitals and will give rise to strong bonds. Electrons in antibonding orbitals will reduce bonding. Hybrid orbitals are derived by mixing atomic wave functions. They are used to explain the geometrical shapes of molecules. Hybridization provides several advantages for bonding. With the larger lobe of the hy￾brid orbital located in between the bonded atoms, more electron density is located where it can “do some good” by contributing to bonding. The participation of different numbers of s and p orbitals in the hybridization al￾lows a wide range of bond angles, thereby permitting electron pairs to get as far away from each other as pos￾sible and minimizing unfavorable electrostatic repulsion. Keep in mind some points for bookkeeping with hybridization. If an atom starts with one s orbital and three p orbitals, it will always end up with a total of four orbitals, no matter how they have hybridized for bonding purposes. Depending on the ratio of atomic orbitals used in the hybridization, we can describe the resulting hy￾brid orbitals as consisting of certain percentages of “s character” and “p character.” For example, an sp orbital contains 50% s and 50% p character, whereas an sp2 hybrid is  1 3  s and  2 3  p in nature. The total s and p charac￾ter around an atom after hybridization always equals that which was present in the orbitals before hybridization: 1. sp hybridized atom (linear geometry): contains two sp orbitals (each one is  1 2  s and  1 2  p in character) and two ordinary p orbitals 2( 1 2  s
 1 2  p)
2p  1s
3p Keys to the Chapter • 3 1559T_ch01_01-17 10/22/05 1:48 Page 3