Chapter 3. MOLECULAR SHAPE AND STRUCTURE THE VSEPR MODEL 3.1 The Basic VSEPR Model 3.2 Molecules with Lone Pairs on the Central Atom 3.3 Polar Molecules VALENCE-BOND THEORY 3.4 Sigma and Pi Bonds 3.5 Electron Promotion and the Hybridization of Orbitals 3.6 Other Common Types of Hybridization 3.7 Characteristics of Multiple Bonds 2012 General Chemistry I 1 THE VSEPR MODEL (Sections 3.1-3.3) 3.1 The Basic VSEPR Model – Lewis structure: showing the linkages between atoms and the presence of lone pairs, but not the 3D arrangement of atoms Valence-shell electron-pair repulsion model (VSEPR model, first devized by Sidgwick and Powell, later modified by Gillespie and Nyholm) is based on Lewis structures. Molecular shapes are predicted by use of several rules that account for bond angles. 2012 General Chemistry I 2 Ideal Molecular Geometries According to VSEPR 2012 General Chemistry I 3 Generic “VSEPR formula”, AXnEm - A = central atom; Xn = n atoms bonded to central atom Em = m lone pairs on central atom E.g. BF3 (AX3), SO2 (AX2E), SO32- (AX3E), CH4 (AX4), PCl5 (AX5) Rule 1. Regions of high electron concentration (bonds and lone pairs on the central atom) repel one another and, to minimize their repulsions, these regions move as far apart as possible while maintaining the same distance from the central atom. This gives rise to the basic geometries, which are related to the total number (Xn + En) of electron pairs around the central atom (next slide). 2012 General Chemistry I 4 2012 General Chemistry I 5 Rule 2. There is no distinction between single and multiple bonds: a multiple bond is treated as a single region of high electron concentration. E.g. 2012 General Chemistry I 6 Self-Test 3.1B Predict the shape of a formaldehyde molecule, CH2O. Solution Although written, CH2O, carbon is the central atom in a H2CO skeleton, hence Lewis structure is O H C H C is bonded to 3 atoms, so the molecule is trigonal planar. 2012 General Chemistry I 7 3.2 Molecules with Lone Pairs on the Central Atom Rule 3. All regions of high electron density, lone pairs and bonds, are included in a description of the electronic arrangement, but only the positions of atoms are considered when identifying the shape of a molecule. E.g. Trigonal pyramidal 2012 General Chemistry I Angular (bent or V-shape) 8 - 2012 General Chemistry I 9 Rule 4 The strength of repulsions are in the order lone pair-lone pair > lone pair-atom > atom-atom. trigonal pyramidal - AX3E type more stable - AX4E type seesaw shaped axial equatorial - AX3E2 type - AX4E2 type T-shaped 2012 General Chemistry I square planar 10 Self-Test 3.3B (a) Give the VSEPR formula of the ClO2- ion. Predict (a) its electron arrangement and (c) its shape. Solution The Lewis structure is _ O .. Cl .. O , showing two 2 atoms and 2 lone pairs on the central Cl atom. (a) Hence its VSEPR formula is AX2E2. (b) The electron arrangement is tetrahedral. (c) Its actual shape is angular (bent or V-shape). 2012 General Chemistry I 11 Predicting the molecular shape of SF4 Step 1 Draw the Lewis structure. Step 2 Assign the electron arrangement around the central atom. Step 3 Identify the molecular shape. AX4E. Step 4 Allow for distortions. (Bent) seesaw shape 2012 General Chemistry I 12 2012 General Chemistry I 13 3.3 Polar Molecules Polar molecule: a molecule with a nonzero dipole moment. A diatomic molecule is polar if its bond is polar (if the electronegativity difference between the two atoms > 0.3) - see Chapter 2.12. E.g. HCl is a polar molecule (dipole moment of 1.1 D). H2 is a nonpolar molecule (zero dipole moment). - A polyatomic molecule is polar if it has polar bonds arranged in space in such a way that the dipole moments associated with the bonds do not cancel. See Table 3.1 for dipole moments of selected molecules. 2012 General Chemistry I 14 polar nonpolar polar 2012 General Chemistry I nonpolar 15 Molecular geometries giving non-polar or polar molecules: AX2 – AX3E2 (Fig. 3.7) 2012 General Chemistry I 16 Molecular geometries giving non-polar or polar molecules: AX4 – AX6 (Fig. 3.7) 2012 General Chemistry I 17 2012 General Chemistry I 18 VALENCE-BOND (VB)THEORY (Sections 3.4-3.7) -The VB theory is based on the Lewis model: each bonding electron pair is localized between two bonded atoms. It is a localized electron model, devized originally by Heitler and London, and later modified by Pauling, Slater and others. 2012 General Chemistry I 19 3.4 Sigma (s) and Pi (p) Bonds The s-bond is formed by ‘head-on’ overlap: there is no nodal surface containing the interatomic (bond) axis. - It is cylindrical or “sausage” shaped. E.g. H2, s(1s, 1s) HF, s(1s, 2pz) N2, s(2pz, 2pz) 2012 General Chemistry I See next slide 2012 General Chemistry I 21 The p-bond is formed by ‘sideways’ overlap: there is a nodal plane containing the interatomic (bond) axis. -It is composed of two cylindrical shapes (lobes), one above and the other below the nodal plane. N2 one s-bond with two perpendicular p-bonds - In General, single bond (one s-bond), double bond (one s- and one p-bond) triple bond (one s- and two p-bonds) 2012 General Chemistry I 22 23 Self-Test 3.6A How many s-bonds and how many p-bonds are there in (a) CO2 and (b) CO? p Solution Hs O s C (a) CO 2 is and two p-bonds. p (b) CO is C sO p 2012 General Chemistry I p O ,with two s-bonds with one s-bond and two p-bonds. 3.5 Electron Promotion and the Hybridization of Orbitals all C-H bonds equivalent ? Electron promotion: electron relocated to a higher-energy orbital promotion 2012 General Chemistry I hybridization 24 Hybrid orbitals: produced by hybridizing orbitals of a central atom C [He]2s12px12py12pz1 4 sp3 hybrids h1 = s + px + py + pz h2 = s - px - py + pz h3 = s - px + py - pz h4 = s + px - py - pz each C-H bond sp3 hybrid orbital 2012 General Chemistry I 25 Examples of sp3 hybrid orbitals in bonding - Ethane, C2H6 Each C atom uses sp3 hybrid orbitals. The C-C bond is formed as s(C2sp3, C2sp3). Each C-H bond is formed as s(C2sp3, H1s). - Ammonia, NH3 Four sp3 hybrid orbitals; one occupied by a lone pair, and the other three forming N-H s-bonds 2012 General Chemistry I 26 3.6 Other Common Types of Hybridization sp2 hybrid orbitals in BF3 sp hybrid orbitals in CO2 2012 General Chemistry I 27 - sp3d hybrid orbitals in PCl5 2012 General Chemistry I - sp3d2 hybrid orbitals in SF6 and XeF4 28 29 2012 General Chemistry I 30 Self-Test 3.8A Describe (a) the electron arrangement, (b) the molecular shape, and (c) the hybridization of the central atom in chlorine trifluoride. Solution (a) The Lewis structure of ClF3 is .. .. F Cl F F The VSEPR formula of ClF3 is AX3E2, hence the electron arrangement is trigonal bipyramidal. (b) The molecular geometry is T-shaped. (c) The hybridization on Cl is sp3d. 2012 General Chemistry I 3.7 Characteristics of Multiple Bonds - Ethene, CH2=CH2 C-C s bond, s(C2sp2, C2sp2) C-C p bond, p(C2p, C2p) each C-H bond formed as s(C2sp2, H1s) 2012 General Chemistry I 31 - Benzene, C6H6 C-C s bond, s(C2sp2, C2sp2) C-C p bond, p(C2p, C2p) Each C-H bond formed as s(C2sp2, H1s) - Ethyne (acetylene), C2H2 C-C s bond, s(C2sp, C2sp) Two C-C p bond, p(C2px, C2px), p(C2py, C2py) Each C-H bond formed as s(C2sp, H1s) 2012 General Chemistry I 32 33 Self-Test 3.9A Describe the structure of the carbon suboxide molecule, C3O2, in terms of hybrid orbitals, bond angles, and s- and p -bonds. The atoms lie in the order OCCCO. Solution .. The Lewis structure is : O C C C .. O: The VSEPR model predicts all three central C atoms to be of the type AX2, and hence bonding to all three is linear. All three C atoms are sp hybridized and oxygen is sp2 hybridized. All bond angles are 180o. Each carbon forms one s- and one p-bond, to each C or O neighbor. p p O s C s C s C s O p p 2012 General Chemistry I Chapter 3. MOLECULAR SHAPE AND STRUCTURE MOLECULAR ORBITAL THEORY 3.8 The Limitations of Lewis’s Theory 3.9 Molecular Orbitals 3.10 Electron Configurations of Diatomic Molecules 3.11 Bonding in Heteronuclear Diatomic Molecules 3.12 Orbitals in Polyatomic Molecules 2012 General Chemistry I 34 MOLECULAR ORBITAL THEORY (Sections 3.8-3.12) 3.8 The Limitations of Lewis’s Theory Valence Bond (VB) theory deficiencies - Cannot explain paramagnetism of O2 (existence of unpaired electrons) - Difficulty treating electron-deficient compounds such as B2H6 - No simple explanation for spectroscopic properties of compounds Molecular Orbital (MO) theory advantages - Addresses all of the above shortcomings of VB theory - Provides a deeper understanding of electron-pair bonds - Accounts for the structure and properties of metals and semiconductors - Universally used in calculations of molecular structures 2012 General Chemistry I 35 3.9 Molecular Orbitals Molecular orbitals (MOs) contain the valence electrons in molecules: they are delocalized over the whole molecule. - In VB theory, bonding electrons are localized on atoms or between pairs of atoms. - n molecular orbitals must be constructed from n atomic orbitals: MOs are formed by linear combination of atomic orbitals (LCAO-MO): - Bonding orbital (constructive interference) y = yA1s + yB1s → overall lowering of energy - Antibonding orbital (destructive interference) y = yA1s - yB1s → overall raising of energy 2012 General Chemistry I 36 Molecular orbital energy-level diagram MO energy-level diagrams are an integral part of MO theory: they give the following information. 1. Relative energies of original AOs and resulting MOs 2. Orbital location of electrons and electron spin (using arrows) H2 - In H2, two 1s-orbitals merge to form the bonding orbital s1s and the antibonding orbital s1s* 2012 General Chemistry I 37 3.10 Electron Configurations of Diatomic Molecules Constructing MOs (by combining AOs) must follow a number of rules,the first three of which are similar to those of the ‘Building-up Principle’ for constructing atomic electron configuration: 1. Electrons are accommodated in the lowest-energy MO, then in orbitals of increasingly higher energy. 2. According to the Pauli exclusion principle, each MO can accommodate up to two electrons. If two electrons are present in one orbital, they must be paired. 3. If more than one MO of the same energy is available, the electrons enter them singly and adopt parallel spins (Hund’s rule). 2012 General Chemistry I 38 39 4. Only atomic orbitals of the same symmetry along the bond axis can be combined. + + s (sideways) + p p (head-on) Not allowed: + p (head-on) or p p + 5. Atomic orbitals of the same or similar energy interact more strongly than those with very different energies. Low energy AOs contribute more to bonding MOs; high energy AOs contribute more to antibonding MOs. 2012 General Chemistry I 40 Homonuclear diatomic molecules of Period 2 elements (Li2 – F2) General Features -Two 2s atomic orbitals (one on each atom) overlap to form two s orbitals: one bonding (s2s-orbital) and the other antibonding (s2s*-orbital). - Six 2p atomic orbitals (three on each atom) overlap to form six MOs: two 2pz orbitals to form s bonding and antibonding (s2p, s2p*) MOs, and four 2px, 2py orbitals to form two p2p and two p2p*MOs. See next slide. 2012 General Chemistry I For O2 and F2 Four 2px, 2py AOs form two p2p and two p2p* MOs Two 2pz AOs form bonding s2p and antibonding s2p* MOs Two 2s AOs form one bonding (s2s) MO and one antibonding (s2s*) MO 2012 General Chemistry I 41 - For O2 and F2, the energy levels of 2s and 2p are well-separated. from Shriver 2012 General Chemistry I 42 - From Li2 to N2, the energy levels of 2s and 2p are close, and thus the 2s orbital also participates in forming s2p orbitals. from Shriver 2012 General Chemistry I 43 MO theory account of bonding in N2 - In N2, each atom supplies five valence electrons. A total of ten electrons fill the MOs. LUMO The ground configuration is, b = ½(8-2) = 3 HOMO Bond order (b): net number of bonds Bond order, b = 1 _ (Ne Ne*) 2 Ne = number of electrons in bonding MOs Ne* = number of electrons in antibonding MOs 2012 General Chemistry I 44 45 MO theory account of bonding in O2 - In O2, each atom supplies six valence electrons. A total of twelve electrons fill the MOs. The ground configuration is, LUMO b = ½(8-4) = 2 accounts for paramagnetism of O2 2012 General Chemistry I HOMO 46 Self-Test 3.10A s2p* Deduce the electron configuration and bond order of the ion C22-. p 2p* Solution s2p The ion C22- has the configuration of C2, but with two extra electrons, which enter the s2p bonding MO. p 2p Configuration is s2s2s2s*2p2p 4s2p2 Bond order is 1/2(8 _ 2) = 3 E s2s* s2s 2012 General Chemistry I 3.11 Bonding in Heteronuclear Diatomic Molecules A diatomic molecule built from atoms of two different elements is polar, with the electrons shared unequally by the two atoms. - In a nonpolar covalent bond, cA2 = cB2 - In an ionic bond, the coefficient belonging to one ion is zero. - In a polar covalent bond, the AO belonging to the more electronegative atom has the lower energy, and so it makes the larger contribution to the lowest energy (bonding) MO. - Conversely, the contribution to the highestenergy (most antibonding) orbital is greater for the higher-energy AO, which belongs to the less electronegative atom. 2012 General Chemistry I 47 - For HF, electronegativity of F (4.0), H (2.2) - For EO (E = C or N), mainly H1s-orbital → partial (+) charge on H mainly F2pz-orbital → partial (–) charge on F 2012 General Chemistry I 48 49 Self-Test 3.11B s2p* Write the configuration of the ground state of the cyanide ion, CN-, assuming that its molecular orbital energy diagram is the same as that for CO. p2p* Solution s2p There are 4 + 5 + 1 = 10 electrons to accommodate into the MOs of Fig. 3.35. Hence the configuration is 1s22s* 21p43s2 (s2s2s2s*2p2p4s2p 2) p2p E s2s* s2s 2012 General Chemistry I 3.12 Orbitals in Polyatomic Molecules - The MOs spread over all atoms in the molecule. Experimentally studied by using ultraviolet and visible spectroscopy. Water, H2O - A water molecule with six atomic orbitals (one O2s, three O2p, and two H1s) antibonding orbitals O2px nonbonding orbitals H1s-O2py-H1s bonding orbitals H1s-(O2s,2pz)-H1s 2012 General Chemistry I 50 Benzene, C6H6 - All thirty C2s-, C2p-, and H1s-orbitals contribute to MOs. - The orbitals in the ring plane: C2s-, C2px, C2py, and six H1s-orbitals → delocalized s-orbitals for C-C and C-H - six C2pz-orbitals perpendicular to the ring → delocalized p-orbitals spreading the ring From VB, each C atom with sp2 hybrid orbitals forming s-bonds and 120° angles. From MO, the six C2pz-orbitals form six delocalized p-orbitals. 2012 General Chemistry I 51 52 Great stability: the p-electrons occupy only orbitals with a net bonding effect. 2012 General Chemistry I Hypervalent compounds (a central atom forms more bonds than allowed by the octet rule) - From VB, SF6 has S with sp3d2 hybridization - From MO, four orbitals of S and six of F, a total of 10 AOs → 10 MOs 12 electrons occupy bonding and nonbonding Orbitals. - Average bond order of each S-F is 2/3. 2012 General Chemistry I 53 Colors of vegetation - highest occupied molecular orbital (HOMO) - lowest unoccupied molecular orbital (LUMO) → excited an electron from a HOMO to a LUMO, by the photons with the energy of visible light Beta-carotene and lycopene contain many conjugated C=C bonds. 2012 General Chemistry I 54 ULTRAVIOLET AND VISIBLE SPECTROSCOPY The Technique - The electrons in the molecule can be excited to a higher energy state, by electromagnetic radiation. Bohr frequency condition, DE = hn - UV-vis absorption gives us information about the electronic energy levels of molecules. i.e. Chlorophyll absorbs red and blue light, leaving the green light present in white light to be reflected. 2012 General Chemistry I 55 Chromophores are characteristic groups of atoms in molecules that absorb certain wavelengths of uv or visible light – p-p* transition in nonconjugated double bonds ~ 160 nm, but for molecules with many conjugated double bonds it is the visible region. – n-p* transition in the carbonyl group ~ 280 nm LUMO HOMO - d-to-d transition in d-metal complexes in visible ranges - charge transfer transition in d-metal complexes electrons migrate from the ligands to the metal atom or vice versa. E.g. deep purple color of MnO42012 General Chemistry I 56