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Chapter 1: Orbitals and Bonding Chapter 1 Topics: Bonding Concepts Chapter 1: Orbitals And Bonding Chapter 1 Topics: Bonding Concepts Look back at “Chemistry, The Central Science” by Brown, Lemay, Bursten! Ionization Potential (Ch. 8) Quantum Numbers (Ch. 8) Ionic Bonds (Ch. 8) Electronic Configuration (Ch. 8) Hund’s Rule (Ch. 8) Pauli Exclusion Principle (Ch. 8) Aufbau Principle (Ch. 8) Atomic Orbitals (Ch. 8) Lewis Structures (Ch. 8) Dipole Moment (Ch. 8) Electronegativity (Ch. 8) Valence Electrons (Ch. 8) Octet Rule (Ch. 8) Resonance Structures (Ch. 8) Bond Dissociation Energy (Ch. 8) Formal Charge (Ch. 8) Covalent Bonds (Ch. 8/9) Nodes (Ch. 8/9) Wave Functions (Ch. 9) Molecular Orbitals (Ch. 9) Orbital Nomenclature Organic Chemists usually don’t use quantum numbers – but we have to remember the correlations: principal quantum number (n) azimuthal quantum number (l) y Shell 1: One S orbital (1s) g r e n Shell 2: One S, three P (2s, 2p) E r e h Shell 3: One S, three P, five D (3s, 3p, 3d) g i H Shell 4: One S, three P, five D, seven F (4s, 4p, 4d, 4f) The three P orbitals are designated Px, Py, Pz (different spacial orientations). magnetic quantum number (m) Ionic Bonding Ionic bonds: One atom transfers electron to another. Molecule held together by electostatic (magnetic) forces. Formed between two atoms of very different electonegativities (>2.0 electronegativity difference) Li F Li F or LiF 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 orbitals!!! Electronegativity Electronegativity and Percent Covalency 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 + H H H or H H or H2 Sharing one additional Sharing one additional electron will lead to a electron will lead to a A line (bond) signifies completely filled valence completely filled valence 2 shared electrons shell shell Both hydrogens have a filled valence shell (shared electrons count for both atoms) Polar Covalent Bonding H F or H F or HF !+ !– 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 non-valence 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 non-valence electrons. Nonbonding electron pairs are frequently omitted! H H H NH 3 N H N H or N H ammonia H H H 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 non-valence electrons. Nonbonding electron pairs are frequently omitted! H CCH 2 2 H C C H H C C H ethylene H H H H H H H C C H or C C H H H H 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 non-valence electrons. Nonbonding electron pairs are frequently omitted! O O H H H H H3CS(O)CH3 H C S C H H C S C H dimethylsulfoxide H H H H O H H H H O H C S C H H C S C H H H H H Formal Charge = Valence e– – Nonbonding e– – 1/2 Bonding e– Atomic Orbitals: A Brief Review node (nodal plane) A 1s orbital A 2p orbital Atomic Orbitals: A Brief Review In General, electrons are lower in energy if: 1. They are closer to a positive charge 2. They are in an orbital with fewer nodes 3. They have a larger orbital space to occupy 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 H or 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) !* H–H H H a !*-orbital (antibonding) H H H H a !-orbital (bonding) ! H–H Molecular Orbitals: A Brief Review !* H–H H H a !*-orbital (antibonding) y g H r H e n E 52 kcal/mol H H a !-orbital (bonding) ! 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. Molecular Orbitals: A Brief Review bond formation H + H H H exothermic by 104 Kcal/mol (!H = –104 Kcal/mol) bond cleavage (BDE) H H H + H endothermic by 104 Kcal/mol (!H = +104 Kcal/mol) The bond dissociation energy (BDE) is the energy required to break a bond homolytically (into a diradical). BDE (H2) = 104 Kcal/mol Molecular Orbitals: A Brief Review Mixing orbitals of opposite phase: leads C C to an antibonding interaction that is destabilizing (higher in energy than the atomic orbitals) C C or 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 C C a !*-orbital (antibonding) C C a !-orbital (bonding) C C ! C–C All mutiple bonds that we will encounter will be !-type 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 H O H this atom has the empty orbital to receive the electrons X incorrect H O + H H O H Resonance 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. O3 O O O O O O O O O ozone O O O O O O O O O +1 A more accurate single representation: –1/2 O –1/2 O O.
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