* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Review for Chapter 6: Thermochemistry
X-ray photoelectron spectroscopy wikipedia , lookup
X-ray fluorescence wikipedia , lookup
Rutherford backscattering spectrometry wikipedia , lookup
Rotational–vibrational spectroscopy wikipedia , lookup
State of matter wikipedia , lookup
Auger electron spectroscopy wikipedia , lookup
Rotational spectroscopy wikipedia , lookup
Molecular Hamiltonian wikipedia , lookup
Heat transfer physics wikipedia , lookup
Hartree–Fock method wikipedia , lookup
Homoaromaticity wikipedia , lookup
Metastable inner-shell molecular state wikipedia , lookup
Woodward–Hoffmann rules wikipedia , lookup
Aromaticity wikipedia , lookup
Physical organic chemistry wikipedia , lookup
Atomic theory wikipedia , lookup
Atomic orbital wikipedia , lookup
Electron configuration wikipedia , lookup
Review for Chapter 10: Chemical Bonding II 1. The valence electrons, the electrons in the outermost occupied shell of an atom, are the electrons involved in chemical bonding. Their arrangement determines the geometry of a molecule. 2. The Valence-Shell Electron-Pair Repulsion (VSEPR) model is a simple model for predicting molecular geometry. It is based on the assumption that valence-shell electron pairs repel one another and tend to stay as far apart as possible. 3. To use the VSEPR model, the Lewis structure for the molecule is drawn and the number of pairs of bonding electrons and lone pairs of electrons are determined. For molecules where there are only bonding pairs of electrons (B) around the central atom (A), Table 10.1 can be used to determine the arrangement of the electron pairs and the molecular geometry. For molecules containing both bonding pairs (B) and lone pairs of electrons (E) around the central atom (A), Table 10.2 is used to predict the arrangement of the electron pairs and the molecular geometry. 4. Lone pairs of electrons repel other pairs of electrons more forcefully than bonding pairs do and thus distort the bond angles from the ideal geometry. 5. A dipole moment is a measure of the charge separation in molecules containing atoms of different electronegativities. For example, HF has a dipole moment because fluorine is more electronegative than H and the electrons spend more time in the vicinity of the F nucleus than near the H nucleus. This leads to a partial negative charge on the F end of the molecule and a partial positive charge on the H end of the molecule. HF is an example of a polar molecule. 6. For molecules containing three or more atoms, the polarity of the bonds and the molecular geometry determine whether there is a dipole moment. Even if polar bonds are present, the molecule will not have a dipole moment if the bond moments cancel out due to a symmetrical arrangement of the bonds around the central atom. If the electron pairs are not symmetrically arranged around the central atom, however, the bond moments may add together to give a dipole moment. For example, H2O has a bent shape and possesses a dipole moment because the bond moments add together instead of canceling each other out as would be the case for a linear molecule. 7. Polar molecules possessing dipole moments can be made to align using an electric field. Nonpolar molecules are not affected by an electric field. 8. Valence bond theory is a quantum mechanical explanation for covalent bond formation. Hybridized atomic orbitals are formed by the combination and rearrangement of orbitals from the same atom. The number of hybridized orbitals is equal to the number of pure atomic orbitals that are combined. 9. The hybrid orbitals and their shapes are summarized in Table 10.4. For linear molecules in which there are two pairs of bonding electrons arranged around a central atom, sp hybrid orbitals are used. For trigonal planar molecules in which three pairs of electrons are arranged around a central atom, sp2 hybrid orbitals are used. For tetrahedral molecules in which four pairs of electrons are arranged around a central atom, sp3 hybrid orbitals are used. For trigonal bipyramidal molecules in which five pairs of electrons are arranged around a central atom, sp3d hybrid orbitals are employed. The inclusion of the d orbital enables “expanded octets” to occur. For octahedral molecules in which six pairs of electrons are arranged around a central atom, sp3d2 hybrid orbitals are used. 10. Sigma bonds are covalent bonds formed by orbitals overlapping end-to-end, with the electron density concentrated between the nuclei of the bonding atoms. 11. Pi bonds are covalent bonds formed by orbitals overlapping sideways, with the electron density concentrated above and below the plane of the nuclei of the bonding atoms. Pi bonds can be made using unhybridized p orbitals. 12. A carbon-carbon double bond consists of one sigma bond and one pi bond. A carboncarbon triple bond consists of one sigma bond and two pi bonds. 13. Molecular orbital theory is another quantum mechanical explanation for covalent bond formation. Bonding is described in terms of the formation of new orbitals that are associated with the molecule as a whole. 14. Two types of molecular orbitals are formed: bonding molecular orbitals and antibonding molecular orbitals. A bonding molecular orbital is lower in energy and greater in stability than the atomic orbitals from which it was formed. An antibonding molecular orbital is higher in energy and lower in stability than the atomic orbitals from which it was formed. 15. Know how to draw molecular orbital diagrams for simple molecules as in Figure 10.26 and 10.27. The electrons fill the molecular orbitals from the lowest energy orbital up. 16. The bond order of a compound is calculated as follows: bond order = ½ (number of electrons in bonding molecular orbitals – number of electrons in antibonding molecular orbitals) Molecules with a bond order of 0 are not stable and do not exist. Molecules with bond orders greater than zero are stable and do exist. 17. The molecular orbital diagrams for Li2, B2, C2, N2, O2, and F2 are shown in Table 10.5. Note that 2p orbitals are lower in energy than 2p orbitals for elements with an atomic number of 7 or less. 18. Molecular orbital theory helps explain why O2 shows paramagnetic properties. This occurs because there are two unpaired electrons in the antibonding 2p* orbitals. Valence bond theory was unable to explain this observation. 19. Delocalized molecular orbitals, in which electrons are shared among more than two atoms, can be formed by the overlap of p orbitals on adjacent atoms. For example, in benzene (C6H6), some electrons are shared among all six carbon atoms. Benzene is described as forming one set of sigma bonds between the carbon atoms and another set of “partial” pi bonds in which electrons are shared among all six carbons. Delocalized molecular orbitals are an alternative approach to resonance structure in explaining observed molecular characteristics.