CVB102 Lecture 1 - Chemical Structure and Reactivity. Contact Information: Dr. Bill Lot Electronic Structure of Atoms

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1 CVB102 Lecture 1 - Chemical Structure and Reactivity Contact Information: Dr. Bill Lot b.lott@qut.edu.au Electronic Structure of Atoms Text: Blackman, et al Pp (Pp recommended) The periodic table of elements A systematic catalogue of elements o Elements are arranged in order of atomic number o Nonmetals are on the right side of the periodic table (with the exception of H) o Metalloids border the stair-step line (with the exception of Al and Po) Characteristics of both metal and nonmetal o Metals are on the left side of the table Ductile Malleability Classical Physics Vs. Quantum Theory Classical physics o Energy is continuous o Can be transferred in any quantity o Explains the transfer of energy at the macroscopic level o Does not explain energy transfer at the atomic level Quantum theory o Energy is not continuous o Can only be transferred in discreet packets (or quanta) o Atoms and molecules emit energy only in certain discrete quantities Waves Wave a vibrating disturbance by which energy is transmitted Lambda distance between identical points of successive waves V number of waves that pass through a point in one second Amplitude vertical distance between the midpoint of the wave and the peak (or trough) Longer wave length lower energy Shorter wave length higher energy High amplitude high intensity Low amplitude low intensity

2 Electromagnetic Radiation (Light) Electromagnetic radiation o Radiant energy o Transmitted in the form of light All forms of radiant energy move through a vacuum at 3x10 8 (speed of light) Particle duality of light Light can behave as stream of particles (photons) Albert Einstein deduced that each photon possesses energy E given by the equation E = hv = Where v is the frequency of light (in Hz or s -1 ) H is the Planck s constant (h = 6.63 x J.s) Lambda is wavelength (in nm) hc lambda Structure of Atoms Nucleus contains protons and neutrons Nucleus is surrounded by electrons o Occupy orbitals (classical physics) o Orbitals (quantum physics) around the nucleus The further an orbital lies from the nucleus the higher the energy of that orbital An atom with electrons in the lowest energy configuration is said to be in the ground state Absorbance An electron in lower energy orbital can be excited into a higher energy orbital by irradiation with light The energy of the light must be exactly equal to the energy difference between the two orbitals for the transition to occur Hv is exactly the energy difference between the orbitals Will only absorb the exact energy to move orbitals Emission When an electron in a higher energy orbital falls into a lower energy orbital a photon with energy hv is released Will only emit the exact energy it takes to move orbitals Continuous spectrum White light through a prism High energy violet Low energy red

3 Emission spectrum of hydrogen Only light specific colors emitted. Orbital energies must be quantized Hydrogen is not continuous Electrons have to go into discrete steps to high and low energy orbitals Bohr model of the atom Observations o Light is not emitted or absorbed in a continuous spectrum o Instead only light with certain energies is absorbed or emitted o In other words the energies of electronic orbits are quantized Conclusion o Only orbits with specific energies allowed o A discrete amount of energy (quantum) separates the energy levels of the orbits De Broglie s equation h lambda = mv Interferences Constructive interference o When waves match they combine Destructive interferences o When waves don t match they disrupt the wave CVB102 Lecture 2 Quantum Theory and Electron Configurations Quantum Mechanics Quantum mechanics is a branch of physics that is useful for describing the behavior of matter and energy on the minute scale of atoms and subatomic particles Quantum mechanics is fundamental to our understanding of all of the fundamental forces of nature except gravity Provides a detailed understanding of electronic structure This is turn leads to an understanding of chemistry Schrödinger equation Can only be solved exactly for 1 electron system

4 Solutions to the equation describe atomic orbitals Atomic Orbitals Atomic orbitals are the wave functions for electrons in atoms The square of the wave function is an expression of the probability of finding an electron in a volume surrounding the nucleus Quantum numbers Increase the principal quantum number = electron further from the nucleus = more energy = less stable Principal Quantum Number, n Must be a positive integer o N = 1, 2, 3, 4,, Azimuthal Quantum Number, l Allowed values are xero and positive integers less than n L = 0, 1, 2, 3, 4,, (n-1) L indicates the shape of the orbital Magnetic Quantum Number, ml Allowed values of ml are integers between l and +l Ml = -l, (-l+1),, 0,, (+l-1), +l if l = 0 then ml = 0 If l = 1 then ml = (2x1)+1 = 3 values Ml indicates the spatial orientation of directionality of the orbital Spin Quantum Number, ms Electron spin of their axis (similar to earth)

5 A spinning charge generates a magnetic field perpendicular to the direction of motion Electrons spin in one of either two opposing directions The electron spin quantum number indicates the direction of spin Ms = + ½, - ½ S Orbitals Spherical o 2s, 3s, etc. contain nodes Only one s orbital per principle quantum number o I.e. one 1s, one 2s, one 3s, P orbitals Two lobes o Three p orbitals per quantum shell starting at n = 2 I.e. three 2p, three 3p, three 4p etc. D orbitals Five d orbitals per quantum shell starting at n = 3 Electron configuration Electron configuration o Distribution of electrons among the various orbitals Ground state o Electronic configuration with the lowest possible energy Almost all of the physical and chemical properties of atoms are related to their electron configurations An understanding of electron configuration is vital Electron configurations are determined by adding the appropriate number of electrons to each orbital in order of increasing energy Aufbau principle o Electrons are assigned to orbitals from lowest energy to highest Pauli exclusion principle o No two electrons can have the same configuration o Cant have the same quantum numbers

6 Hund s rule o Electrons prefer to be unpaired (where possible)