Lecture Notes: Atomic Structure & Periodic Trends PDF

Quantum-mechanical Model of an Atom Periodic Trends Lecture 1 Dr. Khatuna Barbakadze Niels Bohr’s Model of the Atom (1913)  Louis de Broglie (1924) de Broglie hypothesis - wave–particle dual...

Quantum-mechanical Model of an Atom Periodic Trends Lecture 1 Dr. Khatuna Barbakadze Niels Bohr’s Model of the Atom (1913)  Louis de Broglie (1924) de Broglie hypothesis - wave–particle duality concept electrons and all matter have both wave & particle properties. This suggestion is a central part of the theory of quantum mechanics h deBroglie = mv  – wavelength m – mass of electron v – velocity of electron  Heisenberg uncertainty principle (1927) it is impossible to know simultaneously both the momentum p (defined as mass times velocity, p=mv) and the position of a particle with certainty where Δx and Δp are the uncertainties in measuring The position and momentum of the particle, respectively.  (1925-1926) Schrodinger Wave equation – a compact version: E = H Three Quantum Numbers: 1. Principal Quantum Number 2. Orbital Quantum Number 3. Magnetic Quantum Number A fourth quantum number - the spin quantum number describes the behavior of a specific electron and completes the description of electrons in atoms  Principal Quantum Number, n Indicates main energy levels n = 1, 2, 3, 4… Each main energy level has sub-levels - The number of orbitals on given energy level 2 - The total number of electrons on given energy level  Orbital Quantum Number, ℓ (Angular Momentum Quantum Number) Indicates shape of orbital sublevels ℓ = n-1 to 0  Magnetic Quantum Number, ml equal to –l, 0 to +l in integer increments identifies number of orbitals within a sublevel describes spatial orientation orbitals within a sublevel for a certain l, there are (2l+1) integral values of ml Atomic Orbital s for s orbital l=0 spherical in shape (2ℓ +1) = 1 so one spatial orientation 2s and ml =0  Quantum Numbers / Resume n =1 l=0 m=0 s = + ½ ან – ½ n =2 l=0 m=-1 m = 0 s = + ½ ან – ½ l=1 m=+1 n =3 l=0 m = -2 l=1 m=-1 m = 0 s = + ½ ან – ½ l=2 m=+1 m = +2 m=-3 n =4 l=0 m = -2 l=1 m=-1 s = + ½ ან – ½ l=2 m = 0 l=3 m=+1 m = +2 m=+3 The distribution of electrons in energy levels is subject to three basic principles: I. Lowest Energy Principle (Kletchkovsky's Rule) II. Pauli Exclusion Principle III. Hund’s Rule I. II. 3 III. the distribution of electrons in subshells proceeds to give the maximum absolute value of a spin quantum number: Condensed ground-state electron configurations in the first two periods Hund’s rule Trends in the Periodic Table Covalent radii  Atomic radii of Periodicity of the main-group elements atomic radius Atomic radii of the transition elements  Trends in the Periodic Table Periodicity of atomic radius  Ionic radii rcation