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Orbitals And Bonding

Orbitals And Bonding
Chapter 1: Orbitals And Bonding Chapter 1 Topics: Bonding Concepts Look back at your General Chemistry Textbook Ionization Potential Quantum Numbers Ionic Bonds Electronic Configuration Hund’s Rule Pauli Exclusion Principle Aufbau Principle Atomic Orbitals Lewis Structures Dipole Moment Electronegativity Valence Electrons Octet Rule Resonance Structures Bond Dissociation Energy Formal Charge Covalent Bonds Nodes Wave Functions Molecular Orbitals Orbital Nomenclature Organic Chemists usually don’t use quantum numbers – but we have to remember the correlations: principal quantum azimuthal quantum number (n) number (l) One S orbital (1s) Shell 1: (2s, 2p) One S, three P Shell 2: (3s, 3p, 3d) One S, three P, five D Shell 3: (4s, 4p, 4d, 4f) One S, three P, five D, seven F Shell 4: The three P orbitals are designated P , P , P (different spacial orientations). x y z magnetic quantum number (m) Higher EnergyIonic Bonding Ionic bonds: One atom transfers electrons to another. Molecules are held together by electostatic (coulombic) forces. Ionic bonds are formed between two atoms of very different electonegativities (2.0 electronegativity difference) or Li F LiF F Li Loss of one electron will Addition of one electron lead to a completely will lead to a completely empty valence shell filled valence shell (full octet) Atoms are especially stable when all of the valence orbitals are either completely filled or completely empty (the "noble gas" configuration). This has been adapted to the octet rule: (most) atoms are stable when there are 8 electrons in their outermost (valence) shell. For this course: 1st 2nd row atoms can never have more than 8 valence electrons and/or 4 valence orbitalsElectronegativity Covalent Bonding Covalent bonds: Two atoms share electrons. Both atoms can count the shared electrons toward their octet. This type of bond is formed between two atoms of similar electonegativities (2.0 electronegativity difference) H + or or H H H H H H 2 Sharing one additional Sharing one additional electron will lead to a A line (bond) signifies electron will lead to a completely filled valence 2 shared electrons completely filled valence shell shell Both hydrogens have a filled valence shell (shared electrons count for both atoms)Electronegativity and Percent Covalency Polar Covalent Bonding or H F or HF H F – + δ δ H F Electronegativity 2.1 Electronegativity 4.0 Fluorine has a higher electronegativity, and will "pull" electrons toward itself causing bond polarization. This creates a dipole along the bond axis. Identical atoms will share electrons equally: a nonpolar covalent bond. Nonidentical atoms will not share electrons equally: a polar covalent bond.Writing Lewis Structures Lewis Structures: Represent connectivity of a chemical species. Dots reach represent one electron; lines represent a shared electron pair; atomic symbols represent the nucleus and all nonvalence electrons. Nonbonding electron pairs are frequently omitted Writing Lewis Structures Lewis Structures: Represent connectivity of a chemical species. Dots reach represent one electron; lines represent a shared electron pair; atomic symbols represent the nucleus and all nonvalence electrons. Nonbonding electron pairs are frequently omitted H H H NH 3 or N H H H N N ammonia H H HWriting Lewis Structures Lewis Structures: Represent connectivity of a chemical species. Dots reach represent one electron; lines represent a shared electron pair; atomic symbols represent the nucleus and all nonvalence electrons. Nonbonding electron pairs are frequently omitted H CCH 2 2 H C H H C C C H ethylene H H H H H H or H H C C C C H H H HWriting Lewis Structures Lewis Structures: Represent connectivity of a chemical species. Dots reach represent one electron; lines represent a shared electron pair; atomic symbols represent the nucleus and all nonvalence electrons. Nonbonding electron pairs are frequently omitted O O H H H H H CS(O)CH C H H S C 3 3 H S H C C dimethylsulfoxide H H H H H H O O H H H C S C H H C S C H H H H H – – – Formal Charge = Valence e – Nonbonding e – 1/2 Bonding eAtomic Orbitals: A Brief Review node (nodal plane) A 1s orbital A 2p orbitalAtomic Orbitals: A Brief Review In General, electrons are lower in energy if: 1. They are closer to a positive charge (nucleus or multiple nucleii) 2. They are in an orbital with fewer nodes Molecular Orbitals: A Brief Review Mixing orbitals of opposite phase: leads H H to an antibonding interaction that is destabilizing (higher in energy than the atomic orbitals) H or H H H H H Mixing orbitals with same phase: leads to a bonding interaction that is stabilizing (lower in energy than the atomic orbitals)Molecular Orbitals: A Brief Review node (nodal plane) a σorbital (antibonding) H σ H–H H H H a σorbital (bonding) H H σ H–HMolecular Orbitals: A Brief Review a σorbital (antibonding) H σ H–H H H H 52 kcal/mol a σorbital (bonding) H H σ H–H Electrons are 52 kcal/mol (per electron) more stable in a σ H–H orbital (larger orbital space) than in a hydrogen 1s orbital. EnergyMolecular Orbitals: A Brief Review bond formation + H H H H exothermic by 104 Kcal/mol (ΔH = –104 Kcal/mol) bond cleavage (BDE) H H + endothermic by 104 Kcal/mol (ΔH = +104 Kcal/mol) H H The bond dissociation energy (BDE) is the energy required to break a bond homolytically (into a diradical). BDE (H ) = 104 Kcal/mol 2Molecular Orbitals: A Brief Review Mixing orbitals of opposite phase: leads to an antibonding interaction that is C C destabilizing (higher in energy than the atomic orbitals) or C C C C C C Mixing orbitals with same phase: leads to a bonding interaction that is stabilizing (lower in energy than the atomic orbitals)Molecular Orbitals: A Brief Review C C a πorbital (antibonding) π C–C C C a πorbital (bonding) C C π C–C All mutiple bonds that we will encounter will be πtypeX Curved Arrow Notation Curved arrows are used to designate the movement or flow of electrons. the electrons start here (a filled orbital) correct + H O H O H H this atom has the empty orbital to receive the electrons incorrect + H O H O H HResonance Structures Resonance Structures allow Lewis Structures to describe multicenter bonding (more than 2 atoms sharing electrons). There are also situations where resonance structures are used to show bond polarization. O O O O O O O O O O 3 ozone O O O O O O O O O +1 O A more accurate single representation: –1/2 –1/2 O O
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