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Disclaimer Reviews do not cover all the material MIDTERM 2 Review.

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1 Disclaimer Reviews do not cover all the material MIDTERM 2 Review

2 CHEMISTRY Most chemistry is in the electrons (the valence electrons) Atomic orbitals

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4 Representation of the 2p orbitals. Atomic orbitals

5 Wolfgang Pauli Atomic orbitals

6 The orbitals filled for elements in various parts Atomic orbitals

7 CHEMISTRY Full shells make the most stable atoms intra molecular bonding

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9 The HCL molecule has a dipole moment intra molecular bonding

10 The Pauling electronegativity values as updated by A.L. Allred in (cont’d) Arbitrarily set F as 4 intra molecular bonding

11 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

12 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

13 Resonance Forms Lewis structures that differ only in the placement of electrons are resonance forms. For O 3 : Experimentally, it is found that both bonds are nm long. The Lewis structure of O 3 must show both resonance forms. O O O   O O O  = = molecular structure

14 Molecular Shapes AB 2 Linear AB 3 Trigonal planar AB 4 Tetrahedral AB 5 Trigonal bipyramidal AB 6 Octahedral AB 3 E Angular or Bent AB 3 E Trigonal pyramidal AB 3 E 2 Angular or Bent AB 4 E Irregular tetrahedral (see saw) AB 3 E 2 T-shaped AB 2 E 3 Linear AB 6 E Square pyramidal AB 5 E 2 Square planar molecular structure

15 Dipole Moment Nonpolar Polar.. HH O C OO Bond dipoles Overall dipole moment = 0 Bond dipoles Overall dipole moment The overall dipole moment of a molecule is the sum of its bond dipoles. In CO 2 the bond dipoles are equal in magnitude but exactly opposite each other. The overall dipole moment is zero. In H 2 O the bond dipoles are also equal in magnitude but do not exactly oppose each other. The molecule has a nonzero overall dipole moment. Coulomb’s law  = Q r Dipole moment, 

16 Bond Enthalpies and Bond Lengths As bond order increases, the bond enthalpy increases and the bond length decreases. D(C  C) = 348 kJ nm D(C=C) = 614 kJ nm D(C  C) = 839 kJ nm D(C  O) = 358 kJ nm D(C=O) = 799 kJ nm D(C  O) = 1072 kJ nm

17 Hydrogen, H 2 Hydrogen fluoride, HFFluorine, F 2 Molecular orbitals

18 (a) Lewis structure of the methane molecule (b) the tetrahedral molecular geometry of the methane molecule. Molecular orbitals

19 Hybrid Orbitals spsp 2 sp 3 sp 3 dsp 3 d 2 Types of Hybrid Orbitals Shapes: linear triangular tetrahedral trig. bipyram. Octahedral # orbitals: Molecular orbitals

20 The relationship among the number of effective pairs, their spatial arrangement, and the hybrid orbital set required Molecular orbitals

21 (a) Orbitals predicted by the LE model to describe (b) The Lewis structure for carbon dioxide Molecular orbitals

22 The combination of hydrogen 1s atomic orbitals to form MOs Molecular orbitals energies

23 (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

24 The expected MO energy-level diagram for the combustion of the 2 P orbitals on two boron atoms. Molecular orbitals energies

25 The MO energy-level diagrams, bond orders, bond energies, and bond lengths for the diatomic molecules, B 2 through F 2. Molecular orbitals energies

26 16a–26 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. Intermolecular forces

27 16a–27 Larger INTERmolecular forces → Higher melting point Higher boiling point Larger enthalpy of fusion Intermolecular forces

28 16a–28 Larger INTERmolecular forces → Higher melting point Higher boiling point Larger enthalpy of fusion Larger viscosity Higher surface tension Smaller vapor pressure Intermolecular forces

29 Table of Force Energies Type of ForceEnergy (kJ/mol) Ionic Bond Covalent Hydrogen Bonding20-40 Ion-Dipole10-20 Dipole-Dipole1-5 Instantaneous Dipole/ Induced Dipole Intermp;ecular forcesIntermolecular forces

30 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 sausage- like molecules. Intermolecular forces

31 H-Bonding Occurs when Hydrogen is attached to a highly electronegative atom (O, N, F). 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- ++ -- Requires Unshared Electron Pairs of Highly Electronegative Elements Intermolecular forces

32 Copyright © Houghton Mifflin Company. All rights reserved. 16a–32 Intermolecular Forces Summary Intermolecular Intramolecular Intermp;ecular forcesIntermolecular forces

33 16a–33 Which forces? LondonDipoleH-bondionic Xe CH 4 CO 2 CO HBr HF CH 3 OH NaCl CaCl 2 X X X XX XX XX XX X X Intermolecular forces

34 16a–34 Relative forces I2I2 Cl 2 H2SH2SH2OH2O CH 3 OCH3CH 3 CH 2 OH CsBrBr 2 CO 2 CO SF 2 SF 6 > Larger London < H-bond < < polar > > ionic Intermolecular forces

35 Copyright © Houghton Mifflin Company. All rights reserved. 16a–35 Bonding in Solids SOLIDS

36 Copyright © Houghton Mifflin Company. All rights reserved. 16a–36 Examples of Three Types of Crystalline Solids SOLIDS

37 Copyright © Houghton Mifflin Company. All rights reserved. 16a–37 crystals

38 Copyright © Houghton Mifflin Company. All rights reserved. 16a–38 Figure 16.11: Reflection of X rays of wavelength n λ = 2 d sin θ crystals

39 Copyright © Houghton Mifflin Company. All rights reserved. 16a–39 1 ½ ¼ 1/8 Atoms in unit cell crystals

40 Copyright © Houghton Mifflin Company. All rights reserved. 16a–40 Cubic Unit Cells of Metals Simple cubic (SC) Body- centered cubic (BCC) Face- centered cubic (FCC) 1 atom/unit cell 2 atoms/unit cell 4 atoms/unit cell crystals

41 Copyright © Houghton Mifflin Company. All rights reserved. 16a–41 crystals

42 Copyright © Houghton Mifflin Company. All rights reserved. 16a–42 Ion Count for the Unit Cell: 4 Na + and 4 Cl - Na 4 Cl 4 = NaCl Can you see how this formula comes from the unit cell? Your eyes “see” 14 Cl - ions and 13 Na + ions in the figure crystals

43 Copyright © Houghton Mifflin Company. All rights reserved. 16a–43 fcc crystals

44 Copyright © Houghton Mifflin Company. All rights reserved. 16a–44 crystals

45 Copyright © Houghton Mifflin Company. All rights reserved. 17a–45 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 Concentration solutions

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47 Figure 16.55: The phase diagram for water phase cahnges

48 Q total = q1 + q2 + q3 + q4 + q5 phase cahnges

49 Copyright © Houghton Mifflin Company. All rights reserved. 17a–49 Thermodynamics of Phase Changes AB Why does a liquid at A form a solid when the temperature is lowered to B solutions phase cahnges

50 Copyright © Houghton Mifflin Company. All rights reserved. 17a–50 Gases: Large Entropy Liquid: Smaller Entropy Solids: Smallest Entropy solutions phase cahnges

51 Copyright © Houghton Mifflin Company. All rights reserved. 17a–51 Thermodynamics for Phase Change ∆G = ∆H - T∆S 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 = T fusion ) For T less than T fusion, ∆G is negative, solid is stable. Negative for spontaneous process Negative for liquid to solid Positive for liquid to solid solutions phase cahnges

52 Copyright © Houghton Mifflin Company. All rights reserved. 17a–52 Factors Affecting Solubility Gas – solvent: Pressure Effects Henry’s Law: C g is the solubility of gas, P g the partial pressure, k = Henry’s law constant. Carbonated beverages are bottled under > 1 atm. As the bottle is opened, P g decreases and the solubility of CO 2 decreases. Therefore, bubbles of CO 2 escape from solution. solutions

53 Raoult’s Law Raoult’s Law: P A is the vapor pressure of A with solute P A  is the vapor pressure of A alone  A is the mole fraction of A P A = X A P A o P Total = X A P A o + X B P B o solutions

54 Figure 16.44: Behavior of a liquid in a closed container solutions

55 Copyright © Houghton Mifflin Company. All rights reserved. 17a–55 Colligative properties Vapor pressure – Mole fraction Freezing point depression – molality Boiling point elevation – molality Osmosis - Molarity solutions

56 Copyright © Houghton Mifflin Company. All rights reserved. 16a–56 Figure 16.24: A representation of the energy levels (bands) in a magnesium crystal semiconductors

57 Copyright © Houghton Mifflin Company. All rights reserved. 16a–57 Band structure of Semiconductors semiconductors

58 Copyright © Houghton Mifflin Company. All rights reserved. 16a–58 Silicon Crystal Doped with (a) Arsenic and (b) Boron semiconductors

59 Copyright © Houghton Mifflin Company. All rights reserved. 16a–59 Figure 16.34: The p-n junction involves the contact of a p-type and an n-type semiconductor. semiconductors

60 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) semiconductors

61 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 f-block: inner transition elements 3d 4d 5d 6d l = 2 m l = -2,-1,0,1,2 4f 5f l = 3 m l = -3,-2,-1,0, 1,2,3 transition metal complexes transition metal complexes

62 n+/- What is a coordination complex? 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 metal ion ligands charge on complex X +/- n counterion transition metal complexes transition metal complexes

63 transition metal complexes transition metal complexes 2+

64 Common Coordination Numbers of Transition Metal Complexes transition metal complexes transition metal complexes transition metal complexes

65 Classes of isomers transition metal complexes transition metal complexes transition metal complexes

66 Isomers I and II transition metal complexes transition metal complexes transition metal complexes

67 Energy of 3d orbitals t2gt2g egeg transition metal complexes transition metal complexes transition metal complexes

68 Strong/weak fields, d 6 Configuration Paramagnetic – 4 Unpaired Electron Spins Diamagnetic – No Unpaired Electron Spins transition metal complexes transition metal complexes transition metal complexes

69 Correlation of High and Low Spin Complexes With Spectrochemical Series t 2g 4 e g 2 t 2g 3 e g 3 t 2g 6 t 2g 5 e g 1 transition metal complexes transition metal complexes transition metal complexes

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