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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 (colors) at specific wavelengths when energy is added Properties of Light • Wave-Particle Nature of Light – early 1900’s – A Dual Nature • It was discovered that light and e- both have wavelike 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 • All the forms of electromagnetic radiation – Speed of light in a vacuum • 3.0 x 108 m/s Wave Nature of Light • Wavelength – Distance between two corresponding points on adjacent waves –λ – nm • Frequency – Number of waves that pass a given point in a specified time -u – Hz - Hertz Wave Nature of Light • Figure 4-1, page 92 • Equation – c=λu – Speed = wavelength * frequency – Indirectly related! • Spectroscope – Device that separates light into a spectrum that can be seen Particle Nature of Light • Quantum – Minimum quantity of energy that can be lost or gained by an atom • Equation – E=hu • Direct relationship between quanta (particle nature) and frequency (wave nature) • 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 – 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 – ex. voltage – and the ejumps from ground state to excited state – Atoms return to original energy, e- drops back down to ground state – The energy is transferred out of the atom in a NEW FORM • Continuous spectrum – Emission of continuous range of frequencies • Line Emission Spectrum – Shows distinct lines Bohr Model of the Hydrogen Atom • Described electrons as PARTICLES – 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 – E increases as e- moves to farther orbits – http://chemmovies.unl.edu/ChemAnime/BOHRQD/B OHRQD.html • Bohr Model (cont) – ONLY explained atoms with one e• Therefore – only worked with hydrogen!! • The principles of his work is applied to the models of other atoms, but the models do not perfectly fit the experimental data. • Orbits = The circular paths electrons followed in the Bohr model of the atom • 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 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 – 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 Angular Momentum Magnetic Spin Principle Quantum Number • • • • • Which main energy level? (“shell”) The distance from the nucleus Symbol- n n is normally 1-7 Greater n value means farther from the nucleus Angular Momentum Quantum Number • What is the shape of the orbital? • Symbol – l • l = s,p,d,f Magnetic Quantum Number • Orientation of orbital around nucleus • Symbol – ml • s–1 p–3 d–5 f–7 • Every orientation can hold 2 e-!! • A “subshell” is made of all of the orientations of a particular shape of orbital • Figures 4-13, 4-14, 4-15 on page 102-103 Spin Quantum Number • Each e- in one orbital must have opposite spins • Symbol – ms • +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 – place one electron in each equal energy orbital before pairing – 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 Lags 1 behind Lags 2 behind 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 • Long Form: 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- - Outermost electrons; In highest principle quantum # • Inner shell e- - not in the highest energy level • Highest occupied energy level / highest principle quantum number • Valence electrons – outermost e• Examples Electron Dot Notation • More examples Summary Questions 1. How many orbitals are in a d subshell? 2. How many individual orbitals are found in Principle Quantum #3 (the third main energy level) 3. How many orbital shapes are found in Principle Quantum #2? 4. How many electrons can be found in the fourth energy level? 5. A single 4s orbital can hold how many electrons? Back to show Back to show Back to show Back to show Back to show Back to show Back to show Back to show Back to show Back to show