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8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule Chemical bond - attractive force between atoms or ions Ionic bond - electrostatic force between oppositely-

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Presentation on theme: "8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule Chemical bond - attractive force between atoms or ions Ionic bond - electrostatic force between oppositely-"— Presentation transcript:

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.

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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 s 2 p 6 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 1.Add up the valence electrons. (# valence e¯ of atoms = # e¯ available for molecule) Ionic charges: + charge - fewer e¯; − charge - more e¯ 2.Write symbols for the atoms and connect with single bonds. Geometry doesn’t matter at this point. 3.Complete the octets of the outer atoms. 4.Place leftover electrons (in pairs) on the central atom. 5.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 ClO 2, NO, and NO 2 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 More than an Octet Very common for central atom, rare for outer atoms Atoms from the 3 rd 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 d orbitals are low enough in energy to participate in bonding and accept the extra electron density. 8.7 Exceptions to the Octet Rule

14 BH 3 CH 4 PCl 5

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. CCl 4 : experimentally find all Cl-C-Cl bond angles are . Therefore, the molecule cannot be planar. All Cl atoms are located at the vertices of a tetrahedron with the C at its center.

16 Valence Shell Electron Pair 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) 9.2 The VSEPR Model

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18 Octahehron

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21 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. 9.2 The VSEPR Model

22 To determine the geometry: 1.Draw the Lewis structure. 2.Count the total number of electron domains around the central atom which gives electronic geometry. 3.Arrange the electron domains in a geometry which minimizes e  -e  repulsion counting multiple bonds as one bonding pair. 4.Assign molecular geometry. 5.Include multiple bonds in VSEPR structure, but lone pairs not necessary. 9.2 The VSEPR Model

23 BH 3 CH 4 PCl 5

24 H2OH2O CO 3 2-

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.

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29 To minimize e   e  repulsion, lone pairs are always placed in equatorial positions. Trigonal Bipyramidal Geometry

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