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CHEM 2331H (13)
Chapter 2

# Chapter 2 Book Notes

8 Pages
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School
University of Minnesota Twin Cities
Department
Chemistry
Course
CHEM 2331H
Professor
T.Andrew Taton
Semester
Fall

Description
Chapter 2 Book Notes 2.1) Wave Properties of Electrons in Orbitals  Electrons should be treated more like waves than particles  2 types of waves: traveling and standing o Traveling wave: sound waves, water waves o Standing wave: organ pipe or guitar string vibrating o Electron in orbital is more like standing wave  Wavefunction (ψ): mathematical description of wave shape as it vibrates o Wave is positive for one instant then negative (not charges) 2 o Electron density at any point = ψ ; o 1s orbital is spherically symmetric  Wave can also have node (part that does not move)  wave is out of phase (one part is positive while other is negative and vice versa)  2p orbital is like 1 harmonic of string wave  2.1A) Linear Combination of Atomic Orbitals o Linear combination of atomic orbitals (LCAO): when atomic orbitals overlap to create more complex standing waves; # new orbitals generated = # starting orbitals o 1) When orbitals on different atoms act, produce molecular orbitals (MOs)  bonding and antibonding interactions o 2) When orbitals in same atom act, give hybrid atomic orbitals  define geometry of bonds 2.2) Molecular Orbitals  Covalent bonds = stable b/c in bonding region, electrons are close to both nuclei  lowers energy; electrons come between positive nuclei so less nuclei repulsion  Bond length: optimum distance between 2 bonded nuclei; has minimum energy  2.2A) The Hydrogen Molecule; Sigma Bonding o When 2 H atoms approach near each other, their 1s wave functions add constructively (reinforce each other) or destructively (cancel each other out)  constructively bond  increase electron density in bonding region  bonding molecular orbital (bonding MO) o Sigma bond (σ bond): cylindrically symmetrical bond; most of electron density = centered along line connecting nuclei  Most common bond in organic compounds  All single bonds in organic compounds = sigma bonds  Every double and triple bond contains 1 sigma bond o Antibonding molecular orbital: occurs when 2 hydrogen 1s orbitals overlap out of phase (destructively) and cancel each other out where they overlap   canceled out parts are nodes or nodal planes separating 2 atoms  Represented by σ* o  Bonding MOs have lower energy than 1s orbital; antibonding MOs have higher energy than 1s  Electrons have paired spins  more stable  All molecules have bonding and antibonding MOs but antibonding MOs are usually empty; often participate in reactions  2.2B) Sigma Overlap in p Orbitals o Similar as 1s bonding MOs 2.3) Pi (╥) Bonding  Occurs when 2 p orbitals oriented perpendicularly to line connecting nuclei overlap sideways  Electron density is centered above and below line connecting nuclei  Overlap is parallel, not linear like sigma bond  pi bonds are not cylindrically symmetrical  2.3A) Single and Double Bonds o When there is double bond, first pair of electrons goes into sigma bonding o 2 pair of electrons go into pi bonding (can’t go into same orbital/space as 1st pair)  Electron density is above and below sigma bond o 1 sigma bond + 1 pi bond = double bond o Hybrid atomic orbitals: result of mixing orbitals on same atom used to form bonds  more common than 1 sigma and 1 pi bond in organic chemistry 2.4) Hybridization and Molecular Shapes  Valence-shell electron-pair repulsion theory (VSEPR theory): a way to predict bond angles based on how electrons repel each other o If we didn’t have hybridization, we’d just have 90 degree angles since s orbitals are nondirectional and p orbitals have 90 degree angles o 4 pairs of electrons: 109.5 degrees; 3 pairs: 120; 2 pairs: 180  Hybridization allows more electron density in bonding region between nuclei  2.4A) sp Hybrid Orbitals o We can add/subtract orbitals from same atom just as we do with orbitals from different atoms o  sp hybrid orbital: mix of s and p orbitals; electron density’s concentrated toward one side of atom  1 s and 1 p orbital become 2 sp orbitals pointed in opposite directions  the 2 sp orbital results if we add the p orbital with the opposite phase  enhanced electron density in bonding region for sigma bond toward left and right of atom  180 bond angle  linear bonding arrangement 2  2.4B) sp Hybrid Orbitals o Composod of 1 s orbital and 2 p orbitals o 120 bond angle  trigonal geometry o The 3 p orbital is unhybridized  2.4C) sp Hybrid Orbitals o 1 s and 3 p orbitals o o bond angle is 109.5  tetrahydral geometry 2.5) Drawing Three-Dimensional Molecules  Wedge lines means coming out or forward  Dotted lines mean going back into the page  Straight, normal bond lines mean in the plane of the page  Ex: Ethane is like combining 2 tetrahedral molecules together 2.6) General Rules of Hybridization and Geometry  Rule 1: Both sigma bonding electrons and lone pairs can occupy hybrid orbitals. The # hybrid orbitals on an atom is computed by adding the # lone pairs of electrons on that atom.  Rule 2: Use the hybridization and geometry that give the widest possible separation of the calculated # bonds and lone pairs  # hybrid orbitals = # atomic orbitals combined.  Lone pairs take up more space than bonded pairs  Rule 3: For multiple bonds, 1 bond = sigma; 2 and rest = pi bonds 2.7) Bond Rotation  If a bond rotates easily, the same molecule can rotate through different angular atomic arrangements  If not, the different arrangements might be different compounds  2.7A) Rotation of Single Bonds o Conformations: molecular structures that differ only by single bond rotations o Eclipsed: when symmetric molecules have parallel bonds o Staggered: when you take an eclipsed molecule and then twist it a bit o Magnitude of orbital overlap in sigma bond remains same as bond rotates  2.7B) Rigidity of Double Bonds o Double bonds don’t rotate easily because pi bonds then lose overlap o  Different arrangements of atoms are different compounds  different properties 2.8) Isomerism  Isomers: different compounds w/ same molecular formula  2 large classes of isomers: constitutional isomers and stereoisomers  2.8A) Constitutional Isomerism o Constitutional isomers differ in bonding sequence (atoms are connected differently) o ex: n-butane and isobutane; or pentane o There might be different ways to branch of carbon chains o Position of double/triple bonds may be different o May have a ring instead of chain  2.8B) Stereoisomers o Stereoisomers are different b/c of how atoms are oriented in space (same order of bonding still) o Cis-trans isomers are one type of stereoisomers  also called geometric isomers because differ by geometry of groups on double bond  Cis: similar groups are on same side of double bond  Trans: similar groups are on opposite sides of bond  there must be 2 different groups for this to happen 2.9) Polarity of Bonds and Molecules  2.9A) Bond Dipole Moments o Can be nonpolar covalent, polar covalent, or ionic; even each category has differences in magnitude o Bond dipole moment (BDM): measurement of polarity of an individual bond  µ = δ x d  µ = bond dipole moment, δ = amt of charge at either end of dipole, d = distance between the charges -30  units: debye (D); 1 debye = 3.34*10 coulomb meters (must divide the answer you get for µ calculation by this constant to have D units  rule of thumb: µ(in D) = 4.8 x δ x d(in angstroms) o Dipole moments are measured experimentally o Crossed end of arrow is at positive end; arrow head is at negative end  2.9B) Molecular Dipole Moments o Molecular dipole moment (MDM): dipole moment of the entire molecule o Can be measured directly (bond dipole moments are estima
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