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Transcript
Lecture 1 Chemical Bonds: Atomic Orbital Theory and Molecular Orbital Theory Dr. A.K.M. Shafiqul Islam 14.07.08 Summary of Modern Atomic Theory Atoms have an internal structure consisting of one or more subatomic particles: protons, neutrons, and electrons. proton positive charge mass = 1.673 x 10-27 kg neutron no charge mass = 1.675 x 10-27 kg electron negative charge mass = 9.109 x 10-31 kg Summary of Modern Atomic Theory Most of the mass of an atom is concentrated in the nucleus. The nucleus contains one or more positively charged protons, and one or more neutrons with no electrical charge. Summary of Modern Atomic Theory One or more negatively charged electrons are in constant motion somewhere outside the nucleus. The number of electrons is equal to the number of protons; the atom has no overall electrical charge. Summary of Modern Atomic Theory An atom is mostly free space because the volume of the nucleus and the electrons outside the nucleus are extremely small compared to the overall volume of the atom. Summary of Modern Atomic Theory Heisenberg Uncertainty Principle • It is not possible to determine both the position and the momentum of an electron. Wave Function y • Describes the energy of an electron and the probability of finding the electron in a region around the nucleus. Atomic Orbital • The probability distribution about one atomic nucleus (i.e., wave function = atomic orbital). – Amplitudes – numerical magnitudes – Signs – positive or negative – Nodes – values of wavefunction equals zero (given by quantum numbers) Summary of Modern Atomic Theory Heisenberg Uncertainty Principle • It is not possible to determine both the position and the momentum of an electron. Wave Function y • Describes the energy of an electron and the probability of finding the electron in a region around the nucleus. Atomic Orbital • The probability distribution about one atomic nucleus (i.e., wave function = atomic orbital). – Amplitudes – numerical magnitudes – Signs – positive or negative – Nodes – values of wavefunction equals zero (given by quantum numbers) Summary of Modern Atomic Theory Heisenberg Uncertainty Principle • It is not possible to determine both the position and the momentum of an electron. Wave Function y • Describes the energy of an electron and the probability of finding the electron in a region around the nucleus. Atomic Orbital • The probability distribution about one atomic nucleus (i.e., wave function = atomic orbital). – Amplitudes – numerical magnitudes – Signs – positive or negative – Nodes – values of wavefunction equals zero (given by quantum numbers) Summary of Modern Atomic Theory Heisenberg Uncertainty Principle • It is not possible to determine both the position and the momentum of an electron. Wave Function y • Describes the energy of an electron and the probability of finding the electron in a region around the nucleus. Atomic Orbital • The probability distribution about one atomic nucleus (i.e., wave function = atomic orbital). – Amplitudes – numerical magnitudes – Signs – positive or negative – Nodes – values of wave function equals zero (given by quantum numbers) Atomic Orbitals s orbital p orbitals px py pz Hybrid Atomic Orbitals hybrid sp3() sp2 () sp () bonding head-head sideways sideways shape rotation bond tetrahedral free single triangular rigid double linear rigid triple Carbon Atomic Number = 6 2p 2s 1s Carbon Atomic Number = 6 2p 2s 1s Oxygen Atomic Number = 8 2p 2s 1s Oxygen Atomic Number = 8 2p 2s 1s Chemical Bonding Chemical Bonds • The forces holding atoms together in compounds. Valence Electrons • The electrons in the outer shell. Lewis Dot Representation of Atoms • Dots around the chemical symbol of an atom represent the valence electrons. Chemical Bonding Chemical Bonds • The forces holding atoms together in compounds. Valence Electrons • The electrons in the outer shell. Lewis Dot Representation of Atoms • Dots around the chemical symbol of an atom represent the valence electrons. Chemical Bonding Chemical Bonds • The forces holding atoms together in compounds. Valence Electrons • The electrons in the outer shell. Lewis Dot Representation of Atoms • Dots around the chemical symbol of an atom represent the valence electrons. Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P Molecular Orbital Theory Molecular orbitals – delocalized over entire molecule. First Principle • The total number of molecular orbitals is always equal to the total number of atomic orbitals contributed by the atoms that have combined. Molecular Orbitals for H2 Bonding molecular orbital • addition of two atomic orbitals, leads to the probability of finding an e- between the atoms. Antibonding molecular orbital • substration of one atomic orbital from the other, leads to reduced probability of finding an e- between the nuclei, increased in other regions. Molecular Orbital Theory Second Principle • The bonding molecular orbital is lower in energy than the parent orbitals, and the antibonding orbital is higher in energy. Third Principle • The electrons of the molecule are assigned to orbitals of successively higher energy (according to Pauli exclusion principle and the Hund’s rule). The Ionic Bond Transfer of Electrons from One Atom to Another Example: • Sodium chloride Na Cl + 1 e Na + 1e Cl The Ionic Bond The electrostatic attraction between oppositely charged ions. The Covalent Bond Some atoms do not transfer electrons from one atom to another to form ions. Instead they form a chemical bond by sharing pairs of electrons between them. A covalent bond consists of a pair of electrons shared between two atoms. The Covalent Bond Some atoms do not transfer electrons from one atom to another to form ions. Instead they form a chemical bond by sharing pairs of electrons between them. A covalent bond consists of a pair of electrons shared between two atoms. The Covalent Bond Some atoms do not transfer electrons from one atom to another to form ions. Instead they form a chemical bond by sharing pairs of electrons between them. A covalent bond consists of a pair of electrons shared between two atoms. Hydrogen, H2 * E* 1s 1s E Fluorine, F2 Fluorine, F2 F + F F F Fluorine, F2 F + F F F F2 * F 2p F 2p F2 Nitrogen, N2 Nitrogen, N2 N O + + C + N N N O O C O Nitrogen, N2 N + N N N Carbon Dioxide, CO2 O + C + O O C O Nitrogen, N2 N + N N N Carbon Dioxide, CO2 O + C + O O C O Molecular Orbitals for He2 He2 has no net stability – two He atoms have no tendency to combine.