Homework. Chapter 9. Chapter 9. Sigma Bond Formation by Orbital Overlap. Valence Bond Theory VALENCE BOND THEORY
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1 Homework Chapter 9 Chapter 9 11, 21, 25, 27, 29, 31, 35, 39, 45, 51, 65 Bonding and Molecular Structure: Orbital Hybridization and Molecular Chapter 9 Broken into two different sections discussing two different theories on molecular bonding. Valence bond theory Molecular orbital theory VALENCE BOND THEORY VALENCE BOND THEORY Linus Pauling valence electrons are localized between atoms (or are lone pairs). half-filled atomic orbitals overlap to form bonds. Valence Bond Theory A model of bonding in which a bond arises from the overlap of atomic orbitals on two atoms to give a bonding orbital with electrons localized between the atoms. Helps explain the molecular shapes. Takes the different orbital types (s, p, d) and creates hybrids of them. Sigma Bond Formation by Orbital Overlap Two s orbitals overlap
2 Sigma Bond Formation Two s orbitals overlap Two p orbitals overlap Types of Bonds Two basic types of bonds. Sigma (σ) bonds. A bond formed by the overlap of orbitals head to head, and with bonding electron density concentrated along the axis of the bond. Can think of it as a single bond. Pi (π) bonds. The second (and third, if present) bond in a multiple bond; results from sideways overlap of p atomic orbitals. Components of our known bonds. Single bond = 1 sigma bond Double bond = 1 sigma bond + 1 pi bond Triple bond = 1 sigma bond + 2 pi bonds Using VB Theory Bonding in BF 3 F Boron configuration B F F 1s 2s 2p planar triangle angle = 120 o Bonding in BF 3 How to account for 3 bonds 120 o apart using a spherical s orbital and p orbitals that are 90 o apart? Pauling said to modify VB approach with ORBITAL HYBRIDIZATION mix available orbitals to form a new set of orbitals HYBRID ORBITALS that will give the maximum overlap in the correct geometry. Bonding in BF 3 Bonding in BF 3 2s hydridize orbs. 2p rearrange electrons The three hybrid orbitals are made from 1 s orbital and 2 p orbitals 3 sp 2 hybrids. three sp 2 hybrid orbitals unused p orbital Now we have 3, half-filled HYBRID orbitals that can be used to form B-F sigma bonds.
3 Bonding in BF 3 Bonding in CH 4 An orbital from each F overlaps one of the sp 2 hybrids to form a B-F σ bond. F F B F How do we account for 4 C H sigma bonds 109 o apart? Need to use 4 atomic orbitals s, p x, p y, and p z to form 4 new hybrid orbitals pointing in the correct direction. 109 o Bonding in a Tetrahedron Formation of Hybrid Atomic 4 C atom orbitals hybridize to form four equivalent sp 3 hybrid atomic orbitals. Bonding in a Tetrahedron Formation of Hybrid Atomic 4 C atom orbitals hybridize to form four equivalent sp 3 hybrid atomic orbitals. Bonding in CH 4 The Hybridization of the CH 4 How many bonds off of the carbon center? 4 What is the noble gas notation of carbon? [He] 2s 2, 2p 2 So we can take the s and p orbitals of the carbon and create hybrid orbitals.
4 sp 3 Hybridization for carbon in CH 4 Also Works for Other sp 3 Hybrid Compounds Energy 2s 2p orbitals Electrons available to form σ bonds Four sp 3 hybrid orbitals NH 3 Nitrogen has 3 bonds off of it and a lone pair of electrons. Four sp H 2 O 3 hybrid orbitals Oxygen has 2 bonds off of it and two lone pair of electron. Giving us the ability to form 4 σ bonds. Four sp 3 hybrid orbitals How to Determine the Hybridization of an Atom 1. Draw the Lewis Structure if not given. 2. Pick a specific atom if one it not specified. 3. Count up the number of σ bonds off of it. 4. Count up the number of lone pairs of electrons around it. 5. Add these two numbers up. 6. Use the s, p, d orbitals (IN THAT ORDER) to equal the number in part 5. Remember there is only one s orbital, three p orbitals and five d orbitals. For Example (all bonds are single bonds) Multiple Bonds An atom with 4 bonds off of it. sp 3 hybridized An atom with 3 bonds and one lone electron pair. sp 3 hybridized An atom with 1 bond off of it. s hybridized (which is really not a hybridization) An atom with 2 bonds off of it. sp hybridized An atom with 2 bonds and one lone electron pair. sp 2 hybridized An atom with 5 bonds and one lone electron pair. sp 3 d 2 hybridized An atom with 5 bond off of it. sp 3 d hybridized C 2 H 4 Draw the Lewis Structure. What is the hybridization around carbon? sp 2
5 Hybridization Around Carbon in C 2 H 4 σ and π Bonding in C 2 H 4 Un-hybridized p orbital Energy 2p orbitals Three sp 2 hybrid orbitals 2s Giving us the ability to form 3 σ bonds and 1 π bond. σ and π Bonding in CH 2 O σ and π Bonding in C 2 H 2 More Examples H 2 CO C 2 H 2 CSe 2 NH 4 + Exploring O 2 Lewis Dot Structure Determine the Hybridization around the oxygens How many σ bonds? 1 How many lone pairs? 2 Therefore we need how many orbitals? 3 What is the hybridization? sp 2
6 Hybridization Around Oxygen in O 2 The Bonding in O 2 Forms the π bond Un-hybridized p orbital Un-hybridized p orbital Un-hybridized p orbital Energy 2p orbitals Three sp 2 hybrid orbitals Three sp 2 hybrid orbitals of oxygen Three sp 2 hybrid orbitals of oxygen 2s Giving us the ability to form 1 σ bonds and 1 π bond. Forms the σ bond Remember that since the bond between the oxygens is covalent the electrons are shared equally. This causes the two ½ empty orbitals to be filled. Therefore oxygen is magnetic. Paramagnetic or Diamagnetic? Based upon the valence bond theory what would you guess be about the magnetism of O 2? Molecular orbital (MO) theory is an alternative way to view orbitals in molecules. MO theory assumes that pure atomic orbitals of the atoms in the molecule combine to produce orbitals that are spread out, or delocalized, over several atoms or even an entire molecule. The Basic Principles of 1. The total number of molecular orbitals is always equal to the total number of atomic orbitals contributed by the atom that have combined. 2. The bonding molecular orbital is lower in energy than the parent orbitals, and the anti-bonding orbital is higher in energy. 3. The electrons of the molecule are assigned to orbitals of successively higher energy according to the Pauli Exclusion Principle and Hund s Rule. 4. Atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitals are of similar energy 1 st Principle of MO Theory Example of a molecular orbital diagram of H 2. Showing that the number of molecular orbitals that form is equal to the total number of atomic orbitals that are used.
7 1 st Principle of MO Theory This leads to a increased probability that the electrons will reside in the bond region between the two nuclei. This is called a bonding molecular orbital and is the same as a chemical bond in valence bond theory. The probability of finding an electron between the nuclei in the molecular orbital is reduced and the probability of finding the electron in other regions is higher. This type of orbital is called an anti-bonding molecular orbital. There is no counterpart in valence bond theory. The Basic Principles of 1. The total number of molecular orbitals is always equal to the total number of atomic orbitals contributed by the atom that have combined. 2. The bonding molecular orbital is lower in energy than the parent orbitals, and the anti-bonding orbital is higher in energy. 3. The electrons of the molecule are assigned to orbitals of successively higher energy according to the Pauli Exclusion Principle and Hund s Rule. 4. Atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitals are of similar energy 2 nd Principle of Molecular Orbital Theory We can see that the bonding orbital is lower in energy than the parent orbitals, providing stability due to chemical bond formation. What do you think happens when we start to populate the anti-bonding orbital? The Basic Principles of 1. The total number of molecular orbitals is always equal to the total number of atomic orbitals contributed by the atom that have combined. 2. The bonding molecular orbital is lower in energy than the parent orbitals, and the anti-bonding orbital is higher in energy. 3. The electrons of the molecule are assigned to orbitals of successively higher energy according to the Pauli Exclusion Principle and Hund s Rule. 4. Atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitals are of similar energy 3 rd Principle of MO Theory Bond Order Bond Order = ½ [# bonding e - s - # antibonding e - s] 1 0 We can see here how we fill the molecular orbitals based upon the Pauli Exclusion Principle and Hund s Rule. What is the bond order of the two compounds?
8 The Basic Principles of 1. The total number of molecular orbitals is always equal to the total number of atomic orbitals contributed by the atom that have combined. 2. The bonding molecular orbital is lower in energy than the parent orbitals, and the anti-bonding orbital is higher in energy. 3. The electrons of the molecule are assigned to orbitals of successively higher energy according to the Pauli Exclusion Principle and Hund s Rule. 4. Atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitals are of similar energy. 4 th Principle of MO Theory It is theoretically possible that the 1s and 2s orbitals combine to form a molecular orbital, but since the energies are so different it does not occur. Basic MO Diagram for Periods 1 & 2 Sigma Bonding from p π Bonding from p σ & π Bonding from p Sideways overlap of atomic 2p orbitals that lie in the same direction in space give π bonding and antibonding MOs.
9 Chemical Reactivity When we talk about the chemistry of compounds we concern ourselves with what? Valence electrons We do the same thing in MO theory but we also extend it. Highest Occupied Molecular Orbital (HOMO) Lowest Unoccupied Molecular Orbital (LUMO) We can use these two two orbitals describe the chemical reactivity, the shape, and other important properties of a compound. Chapter Highlights 1. Demonstrated more uses of the Lewis Structure 2. Valence Bond Theory Hybridization of the valence orbitals. Helps describe the molecular shapes when we use VESPR Theory 3. Improves upon valence bond theory. Helps explain things where valence bond theory fails. Ex) paramagnetism of O 2
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