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Chapter 11 Chemical Bonding II.docx

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Pippa Lock

1 Chapter 11: Chemical Bonding II: Additional Aspects What a Bonding Theory Should Do  Imagine bringing together two H atoms that are initially far apart  When H atoms are far apart, the two do not interact with each other and by convention, the net energy of interaction between the two H atoms is zero  As the two H atoms approach each other, three types of interactions occur o Each electron is attracted to the other nucleus o Electrons repel each other o The two nuclei repel each other  Net energy of interaction of two H atoms o Starts at zero when atoms are very far apart o At intermediate distances, attractive forces predominate and potential energy is negative o At very small internuclear distances, repulsive forces exceed attractive forces and potential energy is positive o At one particular internuclear distance, potential energy reaches its lowest value  Condition in which two H atoms combine into a H mole2ule through a covalent bond o The nuclei continuously move back and forth, as the molecule vibrates, but the average internuclear distance remains constant o This internuclear distance corresponds to bond length o Potential energy corresponds to the negative of the bond-dissociation energy  Several approaches to understanding bonding o Strength of Lewis theory is in the ease with which it can be applied o VSEPR theory makes it possible to propose molecular shapes that are generally in good agreement with experimental results o However, neither method yields quantitative information about bond energies and bond lengths o Lewis theory has problems with odd-electron species and situations in which it is not possible to represent a molecule through a single structure (ex. resonance) Introduction to the Valence-Bond Method  As two H atoms approach each other, these regions begin to interpenetrate  We say that the two orbitals overlap  Furthermore, we say that a bond is produced between the two atoms because of the high electron density probability found in the region between the atomic nuclei where the 1s orbitals overlap  The increased electron density, with its negative charge, attracts the two positively charged nuclei  In this way, a covalent bond is formed between the two H atoms in the H mo2ecule 2  A description of covalent bond formation in terms of atomic orbital overlap is called the valence-bond method  Creation of a covalent bond in the valence-bond method is normally based on overlap of half-filled orbitals, but sometimes an overlap involves a filled orbital on one atom and an empty orbital on another  Valence-bond method gives a localized electron model of bonding, core electrons and lone-pair valence electrons retain the same orbital locations as in separated atoms and the charge density of the bonding electrons is concentrated in the region of orbital overlap Hybridization of Atomic Orbitals  Number of hybrid orbitals is equal to the number of combining atomic orbitals  In most cases, our description of molecular geometry based on the simple overlap of unmodified atomic orbitals do not conform to observed measurements  Ex. Based on the ground-state electron configuration of valence shell of carbon and employing only half-filled orbitals, we expect existence of a molecule with formula CH 2 and a bond angle of 90 o  CH m2lecule is highly reactive and is observed under only special conditions  The simplest hydrocarbon is methane, a stable, unreactive molecule with molecular formula consistent with the octet rule of the Lewis theory  To obtain this molecular formula by the valence-bond method, we need an orbital diagram for carbon in which there are four unpaired electrons so that the orbital overlap leads to four C-H bonds  To get such a diagram, imagine that one of the 2s electrons in a ground-state C atom absorbs energy and is promoted to an empty 2p orbital  The resulting electron configuration is that of an excited state  Electron configuration of this excited state suggests a molecule with three mutually o perpendicular C-H bonds based on the 2p orbitals of the C atom (90 bond angles)  Fourth bond would be directed to whatever position in the molecule could accommodate the fourth H atom  However, this description does not agree with experimentally determined H-C-H bond o angles, all four of which are fond to be 109.5  Problem is with the way the situation has been defined o We have been describing bonded atoms as though they have the same kinds of orbitals as isolated, nonbonded atoms o This assumption works well for H S and 2H , but we3cannot expect these unmodified pure atomic orbitals to work equally well in all cases 3  One way to deal with this problem is to modify the atomic orbitals of the bonded atoms  These new orbitals, which are directed in a tetrahedral fashion, have energies intermediate between those of 2s and 2p orbitals  This mathematical process of replacing pure atomic orbitals with reformulated atomic orbitals for bonded atoms is called hybridization and the new orbitals are called hybrid orbitals  For methane, the hybridization of one s and three p orbitals forms a new set of four sp 3 hybrid orbitals 3  Each sp hybrid orbital has 25% s character and 75% p character  In a hybridization scheme, the number of hybrid orbitals equals the total number of atomic orbitals that are combined  The symbols identify the numbers and kinds of orbitals involved o Thus, sp signifies that one s and three p orbitals are combined  Note that the three p orbitals move down by ¼ of the energy difference between the s orbitals and p orbitals and that the s orbital moves up by ¾ of that energy difference (
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