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 λλλλ  nm  Frequency Frequency  Number of waves that pass a given point in a specified time νννν  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 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!**

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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  OHRQD.html OHRQD.html 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  /quantum/ /quantum/ /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

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!

Electron Configuration

Representing Electron Configurations  Three Notations  Orbital Notation  Electron Configuration Notation  Electron Dot Notation

Orbital Notation  Uses a series of lines and arrows to represent electrons  Examples

Orbital Notation  More examples

Electron Configuration Notation  Eliminates lines and arrows; adds superscripts to sublevels to represent electrons  Long form examples

Electron Configuration Notation  Short form examples – “noble gas configuration”

Electron Dot Notation  Outer shell e -  Inner shell e -  Highest occupied energy level / highest principle quantum number  Valence electrons – outermost e -  Examples

Electron Dot Notation  More examples

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