Chapter 1

Chapter 1 - Covalent Bonding: Sharing Electrons Mass ratio...

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Unformatted text preview: Covalent Bonding: Sharing Electrons Mass ratio proton : electron = ~ 1800 Ionic Bonding: "No" Sharing of Electrons Who donates and who accepts? Partial Periodic Table Duet Octets Valence electrons Why Do Elements React? Noble Gas Configuration 1. "Ionic" Bonds Na2,8,1 F2,7 Therefore: 2 Na + F2 + 2 Na F -1e +1e [Na2,8]+ [F2,8]- Transfer of Valence Electrons 2. "Covalent" Bonds Elements in the "middle" of the periodic table have a problem with electron affinity (EA) and ionization potential (IP): Cov ale 4C nt Bon ds (He) C 4+ -4e C +4e Shows only valence e (Ne) Compromise: electron-sharing C 2 H2 + 4H + O2 H .. H C H .. H .. HOH .. .. .. .. .. 3. Most Bonds Are "Between" Covalent and Ionic: Polar Covalent H : F: : : + + + A:B + - I : Cl : : : - H3C : F : : : - Pauling electronegativity scale : : 1901-1994 Nobel prizes for chemistry and peace : : : : push pull Increase Decrease : 0.3 Covalent < 0.3 2.0 Polar Covalent < 2.0 Ionic The Shape of Molecules Controlled by valence electron repulsion : : : : Diatomics: linear (of course), e.g. H : H Triatomics: either linear (i.e. not bent), e.g. e.g : : : F: Be : F: : : : : : : : : : : Li : H : F: F: : : : : not : F: Be: H :O H : Cl: Tetraatomics: either trigonal, e.g. : Cl : B : Cl: or bent, when there are lone e-pairs, e.g. : : : : F: : : : : : : : : : : : : : Or pyramidal, when there are lone e-pairs, e.g. N H : H H H C H H H Pentaatomics: tetrahedral, e.g. But other shapes are possible, when there are more electrons, e.g. in transition metals (octahedral etc.). How to distribute valence electrons: Lewis Structures Rule 1: Draw molecular skeleton (given) CO2 OCO CH4 H HC H H Rule 2: Count total number of valence electrons .. . . : .C. :Br . H. O . . Rule 3: Octet (Duet) Rule : : : : : : : : Provide octets (duets for H) around all atoms : : F: F : :O H: H : : : : H :: H :: Rule 4: Take care of charges, if any. Charges occur when the formal "effective" electron count around the nucleus differs from valence electron count. "Effective" electron count: Each bond with two shared electrons counts as 1e; lone pairs count as 2e. Valence electron count: Rule 2 (# of valence electrons). + H + .. : N:::O : H:O:H .. + + H3O NO : C:::O: CO + Example: CO2 C O 2. Valence electrons: O 6e, C 4e 3. Octet rule 1. Atom arrangement: O 16e total Shortcut: 1. Connect all bonded atoms with "2e line". O C O .. .. 2. If there are e left, add them as lone pairs .. any .. to :O C O: atom to give it an octet until no e left. 3. If some atoms lack octet, move lone pairs into shared positions. .. .. .. .. :O C O: .. .. O C O .. .. Resonance Often several octet structures are possible for a molecule: Resonance forms Molecule is a superposition of these forms Form A .. :O: move electrons (pairs) form B .. : O: .. O .. .. Carbonate, CO32-. All forms are equivalent .. .. .. .. .. O .. .. .. .. - C O O .. .. .. .. .. .. O C O C O - Resonance Forms The carbonate ion is delocalized: symmetrical! .. :O : Electrostatic Potential Map: Red = relatively electron rich Blue = relatively electron poor .. .. .. O C O .. - .. Nonequivalent Resonance Forms Rules which ones are better? 1. Octet rule (wins over all other) N + O + N O major 6e H O C major H H O C+ H 6e 2. When there are two or more forms with complete octets: electronegativity rules. Example: enolate ion charge on more e-negative element H H C C major H H H O C C - H O But: N major + O 6e + N O When in doubt, rule 1 wins ! 3. Minimum charge separation O H C O H H O C + O major H Formic acid Note: -: ::: + C O: Rule 1 wins ! 1927 Schrdinger: Wave equations for an electron moving around the nucleus Orbitals: Solutions to wave equations (wavefunctions) Born: Square of a value of the wavefunction = probability of finding the electron there Orbitals Spherical (three-dimensional) shape. Contrast to mechanical waves (guitar string, rubber band) Two-dimensional wave: amplitude 0 + node - +/- = signs, not charges Most important orbitals (for us): node s Orbital "ball" p Orbital spherical "eight" + - Actual solutions: 1s, 2s, 2px, 2py, 2pz, 3s, 3px, 3py, 3pz, etc. #s are related to classical shells (increasing energy) The 1s Orbital The 2s Orbital The Three 3p Orbitals Aufbau Principle or: where the electrons go Energy diagram depicting solutions as energy "levels": 2s 1s px py pz electrons (He) There are rules for "filling up" levels with e 1. Lower energy orbitals filled first (closed shell) 2. Pauli: Exclusion principle: 2e max 3. Hunds rule: Equal energy orbital (i.e. px,y,z) filled with one e each first Energies of Orbitals Calculated for H Atom Bonding Bonding occurs by overlap of atomic orbitals to give molecular orbitals In phase overlap bonding molecular orbital Out of phase overlap antibonding molecular orb. Energy E diagrams antibonding (node) 1s 1s bonding 1s 2p no e bad e go down -Orbitals "split" energy levels when entering into overlap. -The better the overlap [e.g. same type of orbital, same energy (shell)], the larger the splitting. -For orbitals of unequal energy, the higher orbital "goes up", the lower "down". good 1s 2p The Types of Orbital Overlap Hybridization and Shape ++ -H : Be : H - is linear; but Be atom has filled shells (1s)2(2s)2! Li : H 2s + 1s no shape "issues" How does it bond? Use an empty p orbital: This allows for bonds, but gives wrong structure: Better: "Hybridization" of Orbitals Intraatomic overlap of 2s and one 2p orbitals generates two new hybrid molecular orbitals: s + p two sp hybrids, with linear arrangement: sp-Hybrid We shall see next that other combinations of intraatomic overlap (hybridization) are possible: s + p + p 3 sp2 with trigonal shape s + p + p + p 4 sp3 with tetrahedral shape Note: n atomic orbitals n new orbitals Example: Bonding in BH3. Hybridization to Trigonal sp2-Hybrid Bonding in Methane: Hybridization to Tetrahedral sp3-Hybrid Bonding in Ethane: Overlap of Two sp3 Hybrid Orbitals More sp3 Hybrids: NH3 and H2O Double and Triple Bonds: A Preview Molecular Models: Use Them!!! The Dashed-Wedged Line Notation ...
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This note was uploaded on 03/21/2012 for the course CHEM 140A taught by Professor Whiteshell during the Fall '04 term at UCSD.

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