Presentation on theme: "Chapter 4: Arrangement of Electrons in Atoms Chemistry."— Presentation transcript:
Chapter 4: Arrangement of Electrons in Atoms Chemistry
Development of a New Atomic Model There were some problems with the Rutherford model…It did not answer: Where the e - were located in the space outside the nucleus Why the e - did not crash into the nucleus Why atoms produce spectra at specific wavelengths
Properties of Light Wave-Particle Nature of Light – early 1900’s A Duel Nature It was discovered that light and e - both have wave-like and particle-like properties
Wave Nature of Light Electromagnetic radiation – form of energy that exhibits wave-like behavior as it travels through space Electromagnetic spectrum Electromagnetic spectrum Electromagnetic spectrum All the forms of electromagnetic radiation Speed of light in a vacuum 3.0 x 10 8 m/s
Wave Nature of Light Wavelength Wavelength Distance between two corresponding points on adjacent waves λλλλ nm Frequency Frequency Number of waves that pass a given point in a specified time νννν Hz - Hertz
Wave Nature of Light Figure 4-1, page 92 Equation c=λν Indirectly related! Spectroscope Device that separates light into a spectrum that can be seen Diffraction Grating – the part of the spectroscope the separates the light
Particle Nature of Light Quantum Minimum quantity of energy that can be lost or gained by an atom Equation E=hν Direct relationship between quanta and frequency Planck’s Constant (h) Planck’s Constant (h) Planck’s Constant (h) h=6.626 x 10 -34 Js
Particle Nature of Light Photon Individual quantum of light; “packet” The Hydrogen Atom Line emission spectrum (Figure 4-5, page 95) Ground State Lowest energy state (closest to the nucleus) Excited State State of higher energy **When electron drops from its excited state to its ground state, a photon is emitted! This produces a bright-line spectrum. Each element has a characteristic bright-line spectrum – much like a fingerprint!**
Particle Nature of Light Why does an emission spectrum occur? Atoms get extra energy – voltage The e - jumps from ground state to excited state Atoms return to original energy, e - drops back down to ground state Continuous spectrum Continuous spectrum Continuous spectrum Emission of continuous range of frequencies
Particle Nature of Light Bohr Model of the H atom 1913 – Danish physicist – Niels Bohr Single e - circled around nucleus in allowed paths or orbits e - has fixed E when in this orbit (lowest E closest to nucleus) Lot of empty space between nucleus and e - in which e - cannot be in space space E increases as e - moves to farther orbits orbits http://chemmovies.unl.edu/ChemAnime/BOHRQD/B OHRQD.html http://chemmovies.unl.edu/ChemAnime/BOHRQD/B OHRQD.html http://chemmovies.unl.edu/ChemAnime/BOHRQD/B OHRQD.html
Particle Nature of Light Bohr Model (cont) ONLY explained atoms with one e - Therefore – only worked with hydrogen!!
Particle Nature of Light Spectroscopy Study of light emitted by excited atoms Bright line spectrum
The Quantum Model of the Atom e - act as both waves and particles!! De Broglie 1924 – French physicist e - may have a wave-particle nature Would explain why e - only had certain orbits Diffraction Bending of wave as it passes by edge of object Interference Occurs when waves overlap overlap
The Quantum Model of the Atom Heisenberg Uncertainty Principle 1927 – German physicist It is impossible to determine simultaneously both the position and velocity of an e - 12:28-14:28
The Quantum Model of the Atom Schrodinger Wave Equation 1926 – Austrian physicist Applies to all atoms, treats e - as waves Nucleus is surrounded by orbitals Laid foundation for modern quantum theory Orbital – main energy level; 3D region around nucleus in which an e - can be found Cannot pinpoint e - location!!
Quantum Numbers Quantum Numbers Solutions to Schrodinger’s wave eqn Probability of finding an e - “address” of e - Four Quantum Numbers Principle Anglular Momentum Magnetic Spin
Principle Quantum Number Which main energy level? (“orbital” “shell”) Symbol- n n is normally 1-7 (corresponds to period on periodic table) Higher the n, the greater the distance from the nucleus
Angular Momentum Quantum Number What is the shape of the orbital? shape shape F shape F shape F shape Symbol – l l = s,p,d,f When n = 1, l = s n = 2, l = s,p n = 2, l = s,p n = 3, l = s,p,d n = 3, l = s,p,d n = 4, l = s,p,d,f n = 4, l = s,p,d,f http://www.chemeng.uiuc.edu/~alkgrp/mo/gk12 /quantum/ http://www.chemeng.uiuc.edu/~alkgrp/mo/gk12 /quantum/ http://www.chemeng.uiuc.edu/~alkgrp/mo/gk12 /quantum/
Magnetic Quantum Number Orientation of orbital around nucleus Orientation Symbol – m s – 1 p – 3 p – 3 d – 5 d – 5 f – 7 f – 7 Every orientation can hold 2 e - !! Figures 4-13, 4-14, 4-15 on page 102-103
Spin Quantum Number Each e - in one orbital must have opposite spins Symbol – s + ½, - ½ Two “allowed” values and corresponds to direction of spin
Electron Configuration Electron configurations – arrangements of e - in atoms Rules: Aufbau Principle – an e - occupies the lowest energy first Hund’s Rule – each orbital is filled with 1e - first and then the 2 nd e - will fill Pauli Exclusion Principle – no 2 e - in the same atom can have the same set of QN 14:30-18:25
Electron Configuration Representing electron configurations Use the periodic table to write! Know the s,p,d,f block and then let your fingers do the walking!