8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule

1. 8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule • Chemical bond - attractive force between atoms or ions • Ionic bond - electrostatic force between oppositely-charges ions; results from the transfer of electrons from a metal to a nonmetal. • Covalent bond –results from sharing electrons between the atoms; usually found between nonmetals. • Polar covalent – unequal sharing of electrons • Metallic bond – attractive force holding pure metals together.

3. Lewis Symbols • A pictorial representation of the valence electrons • Electrons are represent as dots around the symbol for the element.

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5. The Octet Rule • All noble gases except He have an s2p6 configuration. • Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). • C, N, O, and F “always” obey the octet rule. • Outer atoms obey the octet rule. • Caution: there are many exceptions to the octet rule (section 8.7).

6. 8.3 Covalent Bonding • Covalent bonds can be represented by the Lewis symbols of the elements: • In Lewis structures, each pair of electrons in a bond is represented by a single line:

7. Multiple Bonds • It is possible for more than one pair of electrons to be shared between two atoms, i.e. multiple bonds. • One shared pair of electrons = single bond (e.g. H−H) • Two shared pairs of electrons = double bond (e.g. O=O) • Three shared pairs of electrons = triple bond (e.g. N≡N) • Bond order – number of bonds between two atoms • Generally, bond strength increases and bond distance decreases as bond order increases.

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9. 8.5 Drawing Lewis Structures • Add up the valence electrons. • (# valence e¯ of atoms = # e¯ available for molecule) • Ionic charges: + charge - fewer e¯; − charge - more e¯ • Write symbols for the atoms and connect with single bonds. Geometry doesn’t matter at this point. • Complete the octets of the outer atoms. • Place leftover electrons (in pairs) on the central atom. • If there are not enough electrons to give the central atom an octet, move electrons from outer atoms to form multiple bonds.

10. 8.5 Drawing Lewis Structures • Formal Charges • There may be more than one valid Lewis Structure for a given molecule. • Formal charges are used to determine the most reasonable structure. • Calculate a formal charge (FC) for each atom: • FC = (# valence e¯) − (# e¯ belonging to atom) • Best structure? The one with lowest formal charges and one with the most negative charges on the most electronegative atoms.

11. 8.7 Exceptions to the Octet Rule • There are three classes of exceptions to the octet rule. • Molecules with: • an odd number of electrons • one or more atoms with less than an octet • one or more atoms with more than an octet • Odd Number of Electrons (radicals) • Molecules such as ClO2, NO, and NO2 have an odd number of electrons. (will not see these on worksheet)

12. 8.7 Exceptions to the Octet Rule • Less than an Octet • Relatively rare. • Typical for elements of Groups 1A, 2A, and 3A. • H – 2 electrons (duet rule) • Be – 4 electrons • B – 6 electrons • Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

13. 8.7 Exceptions to the Octet Rule • More than an Octet • Very common for central atom, rare for outer atoms • Atoms from the 3rd period onwards can accommodate more than an octet, e.g. P (10), S (12), Cl (14), Xe (16) • Only exceed the octet rule in Lewis structures when necessary. • How? Beyond the third period, the dorbitals are low enough in energy to participate in bonding and accept the extra electron density.

14. BH3 CH4 PCl5

15. 9.1 Molecular Shapes • Lewis structures show which atoms are physically connected; electron domains • The shape of a molecule is determined by its bond angles. • CCl4: experimentally find all Cl-C-Cl bond angles are 109.5. • Therefore, the molecule cannot be planar. • All Cl atoms are located at the vertices of a tetrahedron with the C at its center.

16. 9.2 The VSEPR Model • ValenceShellElectronPair Repulsion (VSEPR) theory. • Works by positioning electron domains as far apart as possible to minimize electron repulsion • Each region of electrons about central atom is an electron domain. • Single bond – one domain • Lone pair – one domain • Double or triple bond – one domain • Total number of electron domains predicts electronic geometry (or electron-domain geometry) • The arrangement of atoms in space is the molecular geometry (3D shape)

18. Octahehron

21. 9.2 The VSEPR Model • When determining the electronic geometry, all electrons (lone pairs and bonding pairs) are considered. • When naming the molecular geometry, focus only on the positions of the atoms.

22. 9.2 The VSEPR Model • To determine the geometry: • Draw the Lewis structure. • Count the total number of electron domains around the central atom which gives electronic geometry. • Arrange the electron domains in a geometry which minimizes e--e- repulsion counting multiple bonds as one bonding pair. • Assign molecular geometry. • Include multiple bonds in VSEPR structure, but lone pairs not necessary.

23. BH3 CH4 PCl5

24. H2O CO32-

25. The Effect of Lone Pairs and Multiple Bonds on Bond Angles • Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs. • Therefore, the bond angle decreases as the number of lone pairs increase. • Multiple bonds repel more than single bonds, and the same affect is seen.

29. TrigonalBipyramidal Geometry • To minimize e-- e- repulsion, lone pairs are always placed in equatorial positions.