Valence Bond Theory. Valence Bond Theory. Example: H 2. How do bonds form? Other Points about Valence Bond Theory. Example HF. Hybridization and VSEPR

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1 ow do bonds form? Valence Bond Theory ybridization and VSEPR The valence bond model or atomic orbital model was developed by Linus Pauling in order to explain how atoms come together and form molecules. The model theorizes that a covalent bond forms when two orbitals overlap to produce a new combined orbital containing two electrons of opposite in. This overlapping results in a decrease in the energy of the atoms forming the bond. The shared electron pair is most likely to be found in the ace between the two nuclei of the atoms forming the bonds. Valence Bond Theory Example: 2 Valence bond theory describes covalent bond in terms of the overlap of atomic orbitals. 2 : F: F 2 : This type of end-to end overlap of orbitals produces a sigma,σ bond Covalent Bond - Overlapping of the orbitals The newly combined orbital will contain an electron pair with opposite in just like a filled atomic orbital. Example F In hydrogen fluoride the orbital of the will overlap with the half-filled orbital of the F forming a covalent bond. F Overlapping of the and orbitals Other Points about Valence Bond Theory This theory can also be applied to molecules with more than two atoms such as water. Each covalent bond results in a new combined orbital with two oppositely inning electrons. In order for atoms to bond according to the valence bond model, the orbitals must have an unpaired electron. Covalent Bond -F

2 Bonding and Molecular Shape When two atoms like hydrogen come together there is a precise distance between the two orbitals that ensures maximum overlap of the two orbitals The need for maximum overlap is reonsible for the different shapes of molecules found in nature - VSEPR Bond length The distance between the two nuclei in a bond is referred to as bond length The two shared electrons of opposite ins end most of their time between the two nuclei Overlap of orbitals can be between like orbitals (s and s) or unlike orbitals (s and p) Valence Bond Theory and Molecular Geometry Consider the 2 O molecule: 2s Central atom O: 2s Orbital overlap suggests that the bond angle is 90, but we know that the angle is 104.5, therefore there must be a different orbital configuration. ybrid Orbitals To make the connection between Quantum theory and VSEPR shapes, we theorize that the central atom hybridizes the available orbitals to achieve the required number of bonds and the correct molecular shape. What s a ybrid? ybrid Orbitals ybrid orbitals are mixtures of s, p, and d atomic orbitals with intermediate energies. The number of s, p, and d orbitals that have combined, equals the number of hybrid orbitals. s p

3 ybrid Orbitals in BeF 2 Linear geometry is achieved using two hybrid orbitals. 2 ybrid Orbitals in BF 3 Trigonal planar geometry is achieved using three 2 hybrid orbitals. Be atom: B atom: p p p ybrid Orbitals in C 4 Tetrahedral geometry is achieved using four 3 hybrid orbitals. C atom: 3 d and 3 d 2 ybridization Trigonal bipyramidal geometry is achieved using five 3 d hybrid orbitals Octahedral geometry is achieved using six 3 d 2 hybrid orbitals. Types of Bonds There are 2 main types of covalent bonds Single Bonds Areas of electron density are concentrated between the nuclei of the bonding atoms (along the bond) Bond is extremely strong Double and triple bonds Areas of electron density are above and below the plane of the molecule Bonds are highly reactive and easy to break Types of Bonds - σ End to end overlap of orbitals (s, p, d, f, or hybrid) forms sigma, σ, bonds The first bond between the central atom and a ligand is a sigma bond

4 Covalent Bond Formation A (sigma) bond results from end-to-end overlap of orbitals. The maximum electron density lies along the bond. Types of Bonds - π Double and triple bonds are pi, π, bonds A single bond is composed of 2 areas of electron density above and below the sigma bond extremely reactive because they are so far away from the influence of the nucleus and due to their location π bonds are formed from regular p orbitals π bond Covalent Bond Formation A (pi) bond results from side-to-side overlap of p orbitals. The electron density is zero along the bond. Overlap of p y orbitals in O 2 : ybridization in Carbon Carbon is able to form several different hybrid orbitals depending on how many other atoms it is bound to 2 C 3 2 hybridization in Carbon, C 2 4 Recall you only need hybrid orbitals for the first bond between 2 atoms! The rest of the orbitals (the ones that will form the double bond) remain as normal p orbitals 2 hybridization in Carbon, C 2 4 This hybridization forms a planar molecule with 120 angles between atoms. C

5 hybridization in Carbon, C 2 2 Each carbon needs 2 hybrid orbitals for the first bond with the C and atoms. The double and triple bond is formed from 2 normal p orbitals which will form 2 π bonds above and below the central plane of the molecule hybridization in Carbon, C 2 2 The unpaired electrons in the two p orbitals of the two adjacent carbon atoms share electrons by forming two π bonds. hybridization in Carbon, C 2 2 This hybridization forms a linear molecule with 180 angles between atoms. UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2 Determining the ybridization of the Central Atom of a Molecule or Ion