Chemical Bonding

Chemical Bonds, Lewis Symbols, and the Octet Rule Chemical bond: attractive force holding two or more atoms together. Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals. Ionic bond results from the transfer of electrons from a metal to a nonmetal. Metallic bond: attractive force holding pure metals together.

Ionic Bonding

Electron Configurations of Ions and their Lewis Structures

Covalent Bonding

Energy Potential Diagram for Covalent Bonds!

Chemical Bonds Bond Type Single Double Triple # of e’s 2 4 6 Notation — =  Bond order 1 3 Bond strength Increases from Single to Triple Bond length Decreases from Single to Triple

Strengths of Covalent Bonds

Chemical Bonds, Lewis Symbols, and the Octet Rule

Chemical Bonds, Lewis Symbols, and the Octet Rule All noble gases except He has an s2p6 configuration. Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). Caution: there are many exceptions to the octet rule.

Bond Polarity and Electronegativity Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Electronegativity increases across a period and down a group.

Figure 8.6: Electronegativities of Elements Electronegativity

Difference in Electronegativities 0 to 0.5 Nonpolar covalent 0.5 to 1.7 Polar covalent >1.7 Ionic

Bond Polarity and Electronegativity Figure 8.7: Electronegativity and Bond Polarity There is no sharp distinction between bonding types. The positive end (or pole) in a polar bond is represented + and the negative pole -. HyperChem

Drawing Lewis Structures Follow Step by Step Method Total all valence electrons. [Consider Charge] Use Gr # on PT. ex Gr1=1v.e, Gr2=2ve etc Write symbols for the atoms and guess skeleton structure [ define a central atom ]. Place a pair of electrons in each bond. Complete octets of surrounding atoms. [ H = 2 only ] Place leftover electrons in pairs on the central atom. If there are not enough electrons to give the central atom an octet, look for multiple bonds by transferring electrons until each atom has eight electrons around it.

Lewis Dot Diagrams for Individual Elements

Exceptions to the Octet Rule Central Atoms Having Less than an Octet Relatively rare. Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. Most typical example is BF3. Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

Exceptions to the Octet Rule Central Atoms Having More than an Octet This is the largest class of exceptions. Atoms from the 3rd period onwards can accommodate more than an octet. Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density. HyperChem

Drawing Lewis Structures Formal Charge Consider: For C: There are 4 valence electrons (from periodic table). In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge: 4 - 5 = -1.

Drawing Lewis Structures Formal Charge Consider: For N: There are 5 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge = 5 - 5 = 0. We write:

Formal Charge facts to keep in mind: Formal charges need to cancel out on a neutral compound. Example SO3 Formal charges equal (when added) the total charge of the compound. Example NO3-1

Molecular Shapes: VSEPR There are five fundamental geometries for molecular shape:

Molecular Shapes – 3D Notations VSEPR (Ballons)-Movie Clip

Figure 9.3 HyperChem

Summary of VSEPR Molecular Shapes e-pairs Notation Name of VSEPR shape Examples 2 AX2 Linear HgCl2 , ZnI2 , CS2 , CO2 3 AX3 Trigonal planar BF3 , GaI3 AX2E Non-linear (Bent) SO2 , SnCl2 4 AX4 Tetrahedral CCl4 , CH4 , BF4- AX3E (Trigonal) Pyramidal NH3 , OH3- AX2E2 Non-Linear (Bent) H2O , SeCl2 5 AX5 Trigonal bipyramidal PCl5 , PF5 AX4E Distorted tetrahedral (see-sawed) TeCl4 , SF4 AX3E2 T-Shaped ClF3 , BrF3 AX2E3 I3- , ICl2- 6 AX6 Octahedral SF6 , PF6- AX5E Square Pyramidal IF5 , BrF5 AX4E2 Square Planar ICl4- , BrF4-

Examples: VSEPR Molecular Shapes - I # electron pairs on Central Atom A Notation Example Lewis VSEPR & Name of Shape 2 AX2 2 bp on A 3 AX3 3 bp on A AX2E 2 bp and 1 lp on A

Examples: VSEPR Molecular Shapes – I – F08

Examples: VSEPR Molecular Shapes - II # electron pairs on Central Atom A Notation Example Lewis VSEPR & Name of Shape 4 AX4 4 bp on A AX3E 3 bp and 1 lp on A AX2E2 2 bp and 2 lp on A

Examples: VSEPR Molecular Shapes – II – F08

Examples: VSEPR Molecular Shapes - III # electron pairs on Central Atom A Notation Example Lewis VSEPR & Name of Shape 5 AX5 5 bp on A AX4E 4 bp and 1 lp on A AX3E2 3 bp and 2 lp on A AX2E3 2 bp and 3 lp on A

Examples: VSEPR Molecular Shapes – III – F08 HyperChem

Examples: VSEPR Molecular Shapes - IV # electron pairs on Central Atom A Notation Example Lewis VSEPR & Name of Shape 6 AX6 6 bp on A AX5E 5 bp and 1 lp on A AX4E2 4 bp and 2 lp on A

VSEPR Model The Effect of Nonbonding Electrons By experiment, the H-X-H bond angle decreases on moving from C to N to O: 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 increases HyperChem

VSEPR Model Figure 9.10: Shapes of Larger Molecules HyperChem Figure 9.10: Shapes of Larger Molecules In acetic acid, CH3COOH, there are three central atoms.

Figure 8.10: Drawing Lewis Structures Resonance Structures

HyperChem Figure 9.12

Figure 9.11: Molecular Shape and Molecular Polarity HyperChem

Figure 9.13: Molecular Shape and Molecular Polarity HyperChem

Covalent Bonding and Orbital Overlap Lewis structures and VSEPR do not explain why a bond forms. How do we account for shape in terms of quantum mechanics? What are the orbitals that are involved in bonding? We use Valence Bond Theory: Bonds form when orbitals on atoms overlap. There are two electrons of opposite spin in the orbital overlap.

Figure 9.14: Covalent Bonding and Orbital Overlap Optional Topic

VSEPR Model (Figure 9.6) To determine the electron pair geometry: draw the Lewis structure, count the total number of electron pairs around the central atom, arrange the electron pairs in one of the above geometries to minimize e--e- repulsion, and count multiple bonds as one bonding pair.

VSEPR Model

Chemical Bonding Lewis VSEPR shapes AXE notation Polarity