CVB102 – Lecture 1 - Chemical Structure and Reactivity Contact ...

CVB102 – Lecture 1 - Chemical Structure and Reactivity Contact Information: Dr. Bill Lot [email protected] Electronic Structure of Atoms Text: Blackman, et al Pp. 127-147 (Pp. 148 -159 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

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 3x108 (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 10-34J.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

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 𝒍𝒂𝒎𝒃𝒅𝒂 =

𝒉 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

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Can only be solved exactly for 1 electron system

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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

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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)

   

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

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