Transcript Review
MIDTERM 2 Review Disclaimer Reviews do not cover all the material Atomic orbitals CHEMISTRY Most chemistry is in the electrons (the valence electrons) Atomic orbitals Atomic orbitals Representation of the 2p orbitals. Atomic orbitals Wolfgang Pauli Atomic orbitals The orbitals filled for elements in various parts intra molecular bonding CHEMISTRY Full shells make the most stable atoms intra molecular bonding intra molecular bonding The HCL molecule has a dipole moment intra molecular bonding The Pauling electronegativity values as updated by A.L. Allred in 1961. (cont’d) Arbitrarily set F as 4 molecular structure Skeletal Structure • Hydrogen atoms are always terminal atoms. • Central atoms are generally those with the lowest electronegativity. • Carbon atoms are always central atoms. • Generally structures are compact and symmetrical. molecular structure Exceptions to the Octet Rule • Molecules with an odd number of electrons. • Molecules in which an atom has less than an octet of electrons. • Molecules in which an atom has more than an octet of electrons. molecular structure Resonance Forms • Lewis structures that differ only in the placement of electrons are resonance forms. For O3: O O O = O O=O • Experimentally, it is found that both bonds are 0.128 nm long. • The Lewis structure of O3 must show both resonance forms. molecular structure Molecular Shapes AB2 Linear AB3 Trigonal planar AB5 Trigonal bipyramidal AB3E Angular or Bent AB4 Tetrahedral AB4E Irregular tetrahedral (see saw) AB6 Octahedral AB3E Trigonal pyramidal AB3E2 T-shaped AB6E Square pyramidal AB3E2 Angular or Bent AB2E3 Linear AB5E2 Square planar Dipole Moment Bond dipoles C O In H2O the bond dipoles are also equal in magnitude but do not exactly oppose each other. The molecule has a nonzero overall dipole moment. O Overall dipole moment = 0 O Nonpolar The overall dipole moment of a molecule is the sum of its bond dipoles. In CO2 the bond dipoles are equal in magnitude but exactly opposite each other. The overall dipole moment is zero. F k q1 q 2 d m=Qr 2 Coulomb’s law Dipole moment, m Bond dipoles H H Overall dipole moment Polar Bond Enthalpies and Bond Lengths As bond order increases, the bond enthalpy increases and the bond length decreases. D(C-C) = 348 kJ 0.154 nm D(C=C) = 614 kJ 0.134 nm D(CC) = 839 kJ 0.120 nm D(C-O) = 358 kJ 0.143 nm D(C=O) = 799 kJ 0.123 nm D(CO) = 1072 kJ 0.113 nm Molecular orbitals Hydrogen, H2 Hydrogen fluoride, HF Fluorine, F2 Molecular orbitals (a) Lewis structure of the methane molecule (b) the tetrahedral molecular geometry of the methane molecule. Molecular orbitals Hybrid Orbitals Types of Hybrid Orbitals sp Shapes: linear # orbitals: 2 sp2 sp3 sp3d sp3d2 triangular tetrahedral trig. bipyram. Octahedral 3 4 5 6 Molecular orbitals The relationship among the number of effective pairs, their spatial arrangement, and the hybrid orbital set required Molecular orbitals (a) Orbitals predicted by the LE model to describe (b) The Lewis structure for carbon dioxide Molecular orbitals energies The combination of hydrogen 1s atomic orbitals to form MOs Molecular orbitals energies (a) The MO energy-level diagram for the H2 molecule (b) The shapes of the Mos are obtained by squaring the wave functions for MO1 and MO2. Molecular orbitals energies The expected MO energy-level diagram for the combustion of the 2P orbitals on two boron atoms. The MO energy-level diagrams, bond orders, bond energies, and bond lengths for the diatomic molecules, B2 through F2. Molecular orbitals energies Intermolecular Forces Intermolecular forces • The covalent bond holding a molecule together is an intramolecular force. • The attraction between molecules is an intermolecular force. • Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl). • When a substance melts or boils the intermolecular forces are broken (not the covalent bonds). • When a substance condenses intermolecular forces are formed. 16a–26 Intermolecular forces Larger INTERmolecular forces → • Higher melting point • Higher boiling point • Larger enthalpy of fusion 16a–27 Intermolecular forces Larger INTERmolecular forces → • Higher melting point • Higher boiling point • Larger enthalpy of fusion •Larger viscosity •Higher surface tension •Smaller vapor pressure 16a–28 Intermp;ecular forces Intermolecular Table of Force Energies Type of Force Energy (kJ/mol) Ionic Bond 300-600 Covalent 200-400 Hydrogen Bonding 20-40 Ion-Dipole 10-20 Dipole-Dipole 1-5 Instantaneous Dipole/ Induced Dipole 0.05-2 Intermolecular Forces Intermolecular forces London Dispersion Forces • • • • • London dispersion forces increase as molecular weight increases. London dispersion forces exist between all molecules. London dispersion forces depend on the shape of the molecule. The greater the surface area available for contact, the greater the dispersion forces. London dispersion forces between spherical molecules are lower than between sausagelike molecules. Intermolecular forces H-Bonding Occurs when Hydrogen is attached to a highly electronegative atom (O, N, F). d+ N-H… N- O-H… N- F-H… N- N-H… O- O-H… O- F-H… O- N-H… F- O-H… F- F-H… F- dRequires Unshared Electron Pairs of Highly Electronegative Elements Intermp;ecular forces Intermolecular Intermolecular Forces Summary Copyright © Houghton Mifflin Company. All rights reserved. 16a–32 Intermolecular Intramolecular Intermolecular forces Which forces? London Dipole H-bond Xe CH4 CO2 CO HBr HF CH3OH NaCl CaCl2 ionic X X X X X X X X X X X X X 16a–33 Intermolecular forces Relative forces Larger London I2 > < H2O H-bond CH3OCH3 < CH3CH2OH H-bond CsBr > Br2 CO2 < CO SF2 > SF6 H2S ionic polar Cl2 16a–34 polar SOLIDS Bonding in Solids Copyright © Houghton Mifflin Company. All rights reserved. 16a–35 SOLIDS Examples of Three Types of Crystalline Solids Copyright © Houghton Mifflin Company. All rights reserved. 16a–36 crystals Copyright © Houghton Mifflin Company. All rights reserved. 16a–37 crystals Figure 16.11: Reflection of X rays of wavelength n λ = 2 d sin θ Copyright © Houghton Mifflin Company. All rights reserved. 16a–38 crystals Atoms in unit cell 1 ½ ¼ 1/8 Copyright © Houghton Mifflin Company. All rights reserved. 16a–39 crystals Cubic Unit Cells of Metals Simple cubic (SC) 1 atom/unit cell Bodycentered cubic (BCC) 2 atoms/unit cell Facecentered cubic (FCC) 4 atoms/unit cell Copyright © Houghton Mifflin Company. All rights reserved. 16a–40 crystals Copyright © Houghton Mifflin Company. All rights reserved. 16a–41 crystals Your eyes “see” 14 Cl- ions and 13 Na+ ions in the figure Ion Count for the Unit Cell: 4 Na+ and 4 Cl- Na4Cl4 = NaCl Can you see how this formula comes from the unit cell? Copyright © Houghton Mifflin Company. All rights reserved. 16a–42 crystals fcc V o lu m e : D en sity : R 8 4 x m ass R 8 4 x fra ctio n : Copyright © Houghton Mifflin Company. All rights reserved. 16a–43 3 4 3 R 3 R 8 3 3 0 .7 4 crystals Copyright © Houghton Mifflin Company. All rights reserved. 16a–44 Concentration Molarity = Moles of solute/Liters of Solution (M) Molality = Moles of solute/Kg of Solvent (m) Mole Fraction= Moles solute/total number of moles Mass %= Mass solute/total mass x 100 Copyright © Houghton Mifflin Company. All rights reserved. 17a–45 solutions solutions phase cahnges Figure 16.55: The phase diagram for water phase cahnges Qtotal = q1 + q2 + q3 + q4 + q5 Thermodynamics of Phase Changes B phasesolutions cahnges A Why does a liquid at A form a solid when the temperature is lowered to B Copyright © Houghton Mifflin Company. All rights reserved. 17a–49 phasesolutions cahnges Gases: Large Entropy Liquid: Smaller Entropy Solids: Smallest Entropy Copyright © Houghton Mifflin Company. All rights reserved. 17a–50 Thermodynamics for Phase Change ∆G = Negative for spontaneous process • • • • • • • ∆H Negative for liquid to solid phasesolutions cahnges - T∆S Positive for liquid to solid liquid→solid ∆H is negative (stronger intramolecular forces) ∆S is negative (more order) -T∆S is positive As T decreases, -T∆S becomes smaller ∆G goes to zero when ∆H = T∆S (at T = Tfusion) For T less than Tfusion, ∆G is negative, solid is stable. Copyright © Houghton Mifflin Company. All rights reserved. 17a–51 Factors Affecting Solubility solutions Gas – solvent: Pressure Effects Henry’s Law: C g kP g Cg is the solubility of gas, Pg the partial pressure, k = Henry’s law constant. Carbonated beverages are bottled under > 1 atm. As the bottle is opened, Pg decreases and the solubility of CO2 decreases. Therefore, bubbles of CO2 escape from solution. Copyright © Houghton Mifflin Company. All rights reserved. 17a–52 Raoult’s Law • Raoult’s Law: PA is the vapor pressure of A with solute PA is the vapor pressure of A alone A is the mole fraction of A o PA = XA PA o PTotal = XA PA + o X B PB solutions solutions Figure 16.44: Behavior of a liquid in a closed container solutions Colligative properties • • • • Vapor pressure – Freezing point depression – Boiling point elevation – Osmosis - Copyright © Houghton Mifflin Company. All rights reserved. 17a–55 Mole fraction molality molality Molarity semiconductors Figure 16.24: A representation of the energy levels (bands) in a magnesium crystal Copyright © Houghton Mifflin Company. All rights reserved. 16a–56 semiconductors Band structure of Semiconductors Copyright © Houghton Mifflin Company. All rights reserved. 16a–57 semiconductors Silicon Crystal Doped with (a) Arsenic and (b) Boron Copyright © Houghton Mifflin Company. All rights reserved. 16a–58 semiconductors Figure 16.34: The p-n junction involves the contact of a p-type and an n-type semiconductor. Copyright © Houghton Mifflin Company. All rights reserved. 16a–59 semiconductors Semiconductors – key points to remember • Band structure: Valence band – gap – conduction band •DOPING: Group V n type, Group III p type •n-p junctions •Devices: (LED, laser, transistor, solar cell) transition metal complexes What is a transition metal? “an element with valance d- or f-electrons” ie. a d-block or f-block metal d-block: transition elements 3d 4d 5d l=2 ml = -2,-1,0,1,2 6d l=3 ml = -3,-2,-1,0, 4f 1,2,3 5f f-block: inner transition elements transition metal complexes What is a coordination complex? charge on complex n+/- ligands X+/n metal ion counterion •Central metal ion or atom surrounded by a set of ligands •The ligand donates two electrons to the d-orbitals around the metal forming a dative or coordinate bond transition metal complexes -1 2+ transition metal complexes Common Coordination Numbers of Transition Metal Complexes Classes of isomers transition metal complexes transition metal complexes Isomers I and II Energy of 3d orbitals eg t2g transition metal complexes transition metal complexes Strong/weak fields, d6 Configuration Paramagnetic – 4 Unpaired Electron Spins Diamagnetic – No Unpaired Electron Spins Correlation of High and Low Spin Complexes With Spectrochemical Series t2g4eg2 t2g3eg3 t2g6 t2g5eg1 transition metal complexes