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- 1. Some Foundations of Inorganic Chemistry Part 1 Assoc. Prof. Dr. Pham Chien Thang Department of Inorganic Chemistry
- 2. Textbooks 2
- 3. Outline 1. Introduction to Inorganic Chemistry 2. Atomic Structures 3. Molecular Structures & Bonding 3
- 4. What is Inorganic Chemistry? • Chemistry of inorganic substances (‘everything else’ excluded organic compounds) 4 Inorganic Chemistry Organic Chemistry Physical Chemistry Materials Analytical Chemistry Biochemis try
- 5. Outline 1. Introduction to Inorganic Chemistry 2. Atomic Structures 3. Molecular Structures & Bonding 5
- 6. Element & atom • Chemical element: substance that cannot be broken down into simpler substances & made of only one type of atom • Chemical element has its own chemical symbol • Atom: smallest constituent unit of ordinary matter that has properties of a chemical element.1 6 1 ‘Atom’. Compendium of Chemical Terminology (IUPAC Gold Book) (2nd ed.). IUPAC. Retrieved 2015-04-25
- 7. History of atomic structures 7
- 8. Atomic structure • Atom: nucleus (center) & electron shells (outside) • Nucleus: protons & neutrons • Electron shells consist of electrons • Most of atom is empty space !!! 8 nucleus electron shells proton neutron electron Particle Symbol Mass /kg Relative mass Charge /C Relative charge Proton p 1.673 × 10–27 1 +1.602 × 10–19 +1 Neutron n 1.675 × 10–27 1 0 0 Electron e 9.109 × 10–31 1/1836 (~0) –1.602 × 10–19 –1 • Most of mass of atom concentrated in nucleus !!!
- 9. Atomic & mass number • Proton/atomic number (Z): number of protons • Each element has a unique Z • Nucleon/mass number (A): number of protons & neutrons • Isotopes: atoms with same Z & different A • Isotopes have same chemical properties 9
- 10. • Relative atomic mass ( ): weighted mean of mass numbers of all naturally occurring isotopes • Example: Mg: 78.99% 24Mg 10.00% 25Mg 11.01% 26Mg masses 23.99 24.99 25.98 Isotopes 10
- 11. Isotopes 11 • Determination of accurate Ar from mass spectrometry Mass spectrometer Mass of each isotope (mass/charge ratio) How much of each isotope (relative abundance) Isotopic mass Abundance /% 20 90.9 21 0.3 22 8.8 Mass spectrum
- 12. Electronic structures of atoms • Schrödinger equation: • Solution of the equation provides: Wavefunctions → atomic orbitals (AOs) Energy E associated with particular 12 Why ?
- 13. AOs & Quantum Numbers • AOs defined by 3 quantum numbers: Principal quantum number ( ): related to size & energy AO Orbital quantum number ( ): determines shape of orbital Magnetic quantum number ( ): related to orientation of AO 13 principal orbital magnetic
- 14. AO shapes 14
- 15. Ground State Electronic Configuration • aufbau (building-up) principle: lowest energy AOs filled first • Pauli exclusion principle: maximum of 2 electrons in an AO • Hund’s rule: maximum total spin Example: ground state electronic configuration of N (Z = 7), Cu (Z = 29) 15
- 16. Classic Periodic Table 16 • Most successful classification by Dmitri Mendeleev in 1869 Dmitri Ivanovich Mendeleev (1834 –1907)
- 17. Classic Periodic Table 17 • Arrangement of elements in order of atomic weight • Elements with similar chemical properties in same group Annalen der Chemie und Pharmacie, VIII, Supplementary Volume for 1872, page 511
- 18. Modern Periodic Table 18 • Arrangement of elements in order of atomic number row ~ period column ~ group
- 19. The Periodic Table & Electronic Configuration 19 • Period: elements with same electron shells • Period number: number of electron shells • Group: elements with same valence electronic configuration • Group number: number of valence electrons with ‘1–18’ or ‘I – VIII’ numbering system
- 20. Outline 1. Introduction to Inorganic Chemistry 2. Atomic Structures 3. Molecular Structures & Bonding 20
- 21. Basic concepts • A molecule is an electrically neutral group of two or more atoms held together by chemical bonds.1 • Classification of chemical bonds: Ionic bonds: electron transfer between atoms to form ions Covalent bonds: electron sharing between atoms Metallic bonds 21 1 ‘Molecule’. Compendium of Chemical Terminology (IUPAC Gold Book) (2nd ed.). IUPAC. Retrieved 2015-04-25
- 22. Molecular structure & Bonding 1. The Octet Rule 2. Lewis Diagram 3. Molecular Shapes & the VSEPR Model 4. Valence Bond (VB) theory 5. Molecular Orbital (MO) theory 22
- 23. The octet rule • Atoms gain, lose or share electrons to give an outer shell with 8 electrons (an octet) • Elements in Period 2 strictly obey the octet rule 23 If atoms is not only 8 electron in outer shell ?
- 24. The octet rule • Example: MgO Mg (Z = 12) 1s2 2s22p6 3s2 Mg – 2e → Mg2+ (1s2 2s22p6) O (Z = 8) 1s2 2s22p4 O + 2e → O2– (1s2 2s22p6) → MgO ionic compound 24 O x x x x x x x Mg x x x + x x O2– x x x x x x Mg2+ x x x x x x
- 25. The octet rule • Example: O2 O (Z = 8) 1s2 2s22p4 → O2 covalent compound 25 O x x x x x x O + O x x x x x x O
- 26. Molecular structure & Bonding 1. The Octet Rule 2. Lewis Diagram 3. Molecular Shapes & the VSEPR Model 4. Valence Bond (VB) theory 5. Molecular Orbital (MO) theory 26
- 27. Lewis Electron-dot Diagrams • Using dots to describe number & arrangement of valence electrons in molecules • Example: • Bonding pairs: pairs of electrons involved in bonding & represented by a single line (–) • Non-bonding electrons: electrons not involved in bonding Atoms Molecule
- 28. Lewis Electron-dot Diagrams • Constructing Lewis structure: 1. Determine overal number of valence electrons 2. Write arrangement of atoms bonded together 3. Distribute electrons in pairs so that each atom has an octet • Example: CO2, SCN– 28
- 29. Resonance structures • More than one valid Lewis structure with a given atomic arrangement → resonance structures • Resonance indicated by double-headed arrow • Actual structure ~ a resonance hybrid of all resonance structures • Example: CO3 2– 29 a given chemical formula resonace structure describe the position of elctrons in compound.
- 30. Higher electron counts • Number of electrons around central atom over 8 → expanded shell & hypervalent molecules • Frequently observed for elements of 3rd & higher periods due to the d orbitals • Example: PCl5, SF6 30
- 31. Molecular structure & Bonding 1. The Octet Rule 2. Lewis Diagram 3. Molecular Shapes & the VSEPR Model 4. Valence Bond (VB) theory 5. Molecular Orbital (MO) theory 31
- 32. VSEPR Model & Molecular Shapes • Molecular shape: 3D arrangement of atoms that constitute a molecule • Experimentally determined by spectroscopic methods • Valence-Shell Electron-Pair Repulsion (VSEPR) model • Simple but useful to predict shape of small molecules of main group elements 32
- 33. VSEPR Model & Molecular Shapes • Valence-Shell Electron-Pair Repulsion (VSEPR) model • Assumption: valence electron pairs adopt arrangements that minimize repulsion between them 33 Electron pairs 2 3 4 5 6 Geometry linear trigonal planar tetrahedral trigonal bipyramidal octahedral Standard bond angle 180° 120° 109.5° 90°, 120° 90°
- 34. VSEPR Model & Molecular Shapes • Valence-Shell Electron-Pair Repulsion (VSEPR) model • Electron–electron repulsions decrease in sequence: lone-lone pair > lone-bonding pair > bonding-bonding pair triple -single bond > double -single bond > single -single bond 34 VSEPR model describe the 3D shape in space so of compound
- 35. Electron pairs Bonding pairs Non- bonding pairs Electronic geometry Molecular shape Shape 2 2 0 Linear Linear 3 3 0 Trigonal planar Trigonal planar 3 2 1 Trigonal planar Bent VSEPR Model & Molecular Shapes 35
- 36. Electron pairs Bonding pairs Non- bonding pairs Electronic geometry Molecular shape Shape 4 4 0 Tetrahedral Tetrahedral 4 3 1 Tetrahedral Trigonal pyramidal 4 2 2 Tetrahedral Bent VSEPR Model & Molecular Shapes 36
- 37. Electron pairs Bonding pairs Non- bonding pairs Electronic geometry Molecular shape Shape 5 5 0 Trigonal bipyramidal Trigonal bipyramidal 5 4 1 Trigonal bipyramidal Seesaw 5 3 2 Trigonal bipyramidal T-shaped 5 2 3 Trigonal bipyramidal Linear VSEPR Model & Molecular Shapes 37 equatorial axial
- 38. Electron pairs Bonding pairs Non- bonding pairs Electronic geometry Molecular shape Shape 6 6 0 Octahedral Octahedral 6 5 1 Octahedral Square pyramidal 6 4 2 Octahedral Square planar VSEPR Model & Molecular Shapes 38
- 39. Molecular structure & Bonding 1. The Octet Rule 2. Lewis Diagram 3. Molecular Shapes & the VSEPR Model 4. Valence Bond (VB) theory 5. Molecular Orbital (MO) theory 39
- 40. Valence Bond (VB) theory • Principle: chemical bonds ~ overlap of valence AOs on two different nuclei so that both nuclei share pair of electrons • Localized electron model (localized AOs) • Homonuclear diatomic molecules 40 x y z ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ F F σ bond H H ↑↓ σ bond H2 F2 F (Z = 9) 1s2 2s2 2p5
- 41. Valence Bond (VB) theory • Homonuclear diatomic molecules 41 x y z O (Z = 8) 1s2 2s2 2p4 O2 π bond ↑↓ ↑↓ ↑↓ ↑ ↓ O O σ N (Z = 7) 1s2 2s2 2p3 N2 π ↑↓ ↓ ↑ ↑ ↓ N N σ π observe empty orbital in atom and
- 42. VB theory & Hybridization of AOs Polyatomic molecules – Hybridization of AOs • Hybridization of AOs: combinations of AOs to mathematically obtain new AOs for molecular shapes in terms of σ-bonds • Number of hybrid orbitals = number of AOs mixed • Label of hybrid orbitals reflects contributing AOs • Type of hybrid orbitals varies with types of AOs mixed 42
- 43. sp3 hybridization ↑↓ ↑ ↑ ↑ ↑ ↑ ↑ ↑↓ ↑ ↑ ↑ ↑ ↑ ↑ ↑↓ ↓↑ ↑↓ ↑↓ s px py pz sp3 43 E E
- 44. sp2 hybridization ↑↓ ↑ ↑ ↑ ↑ ↑↓ ↑ ↑↓ ↓↑ ↑↓ ↑ ↑ ↑ s px py pz sp2 pz 44 E E
- 45. sp hybridization ↑↓ ↑ ↑ ↑↓ s px py pz sp pz py ↑ ↑ ↓↑ ↑↓ 45 E E
- 46. Hybridization related to d-orbitals sp3d sp3d2 s p p p sp3d d d d d d d d d d s p p p sp3d2 d d d d d d d d 46 E E
- 47. Double & Triple bonds in VB theory C C ↓↑ ↓↑ ↓↑ ↓↑ ↓↑ ↓↑ ↓↑ ↓↑ ↑ ↑ ↓ ↓ ↑ ↓ 47
- 48. Molecular structure & Bonding 1. The Octet Rule 2. Lewis Diagram 3. Molecular Shapes & the VSEPR Model 4. Valence Bond (VB) theory 5. Molecular Orbital (MO) theory 48
- 49. Molecular Orbital (MO) theory • Generalization AO description of atoms to MO description of molecules • MOs: space spread whole molecule, in which a single electron can occupy → delocalized electron model • Linear Combination of Atomic Orbital (LCAO) approximation 49 when 2 atom bond form bond, and valence atomic orbital is wave function and the property of wwave fuction is the same when 2 atom bond form bond, and valence atomic orbital is wave function and the property of wwave fuction is the same
- 50. Molecular Orbital (MO) theory • MO arises from interactions between AOs if interactions are: Allowed if symmetries of AOs are compatible Efficient if region of overlap is significant Efficient if AOs are relatively close in energy • Number of MOs equal number of contributing AOs 50 s s px py px s pz s
- 51. Molecular Orbital (MO) theory • Homonuclear diatomic molecules destructive interaction ↑↓ ↑ H H H2 51 E anti-bonding MO constructive interaction bonding MO ↑
- 52. Molecular Orbital (MO) theory 52
- 53. Orbital energies for main group elements 53 J. B. Mann, T. L. Meek, L. C. Allen, J. Am. Chem. Soc., 2000, 122, 2780
- 54. MO diagram of X2 (X = O, F, Ne) X X X2 2s 2s 2p 2p 54 E
- 55. MO diagram of X2 X X X2 2s 2s 2p 2p 2s 2s 2p 2p X X X2 55 E (X = O, F, Ne) (X = Li, Be, B, C, N)
- 56. Molecular Orbital (MO) theory • Heteronuclear diatomic molecules antibonding MOs close to less electronegative atom bonding MOs close to more electronegative atom 56 X Y XY less electronegative more electronegative E
- 57. MO diagram of CO 2s 2s 2p 2p C O CO (-32.38 eV) (-15.85 eV) (-10.66 eV) (-19.43 eV) ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 57 E
- 58. Molecular Orbital (MO) theory • Very unequal energies → resulting MOs with energies & shapes closer to original AOs 58 ∗ X Y XY E less electronegative more electronegative
- 59. MO diagrams of HF & LiF 𝜎 𝜎 𝜎∗ 𝜋 2s 1s 2p H F HF (-40.2 eV) (-18.7 eV) (-10.7 eV) ↑↓ ↑ ↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑↓ 𝜎 𝜎 𝜎∗ 2s 2s 2p Li F LiF (-40.2 eV) (-18.7 eV) (-5.4 eV) ↑↓ ↑ ↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑↓ 𝜋 59 E E