* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Thursday, 1/29/09 - Liberty Union High School District
Photosynthesis wikipedia , lookup
Bremsstrahlung wikipedia , lookup
Chemical bond wikipedia , lookup
Bohr–Einstein debates wikipedia , lookup
Quantum electrodynamics wikipedia , lookup
Particle in a box wikipedia , lookup
Ultrafast laser spectroscopy wikipedia , lookup
X-ray photoelectron spectroscopy wikipedia , lookup
X-ray fluorescence wikipedia , lookup
Astronomical spectroscopy wikipedia , lookup
Tight binding wikipedia , lookup
Double-slit experiment wikipedia , lookup
Atomic orbital wikipedia , lookup
Matter wave wikipedia , lookup
Hydrogen atom wikipedia , lookup
Electron configuration wikipedia , lookup
Theoretical and experimental justification for the Schrödinger equation wikipedia , lookup
Bohr & New Atomic Model Arrangements of Electrons in Atoms Thursday, October 29th, 2015 1 The Development of a New Atomic Model •Rutherford’s model was incomplete •How were electrons distributed? •What prevents negative electrons from crashing into positive nucleus? •1900s-New Atomic Model (revolutionary) •Relationship between light & electrons 2 How many Electrons Can Fit in Each Energy Level? Max # e = 2n2 Big Question: Why? • Energy level is also called “principal quantum number” • Energy Level = n n 1 2 3 4 5 6 max. # electrons 2 8 18 32 50 72 3 There are two ways to state the location of an electron… • Quantum numbers – the location of a specific electron • Electron configurations – the addresses of every electron in an atom (read the periodic table like a book!) Quantum numbers • Principle quantum number (n) = identifies the energy level of the electron • Angular momentum quantum number (l) = identifies the sublevel • Spin quantum number = identifies the orientation in the pair 5 6 Properties of Light •Before 1900s-thought light behaved as a wave •During 1900s-also has particle characteristics •Electromagnetic Radiation 7 8 Electromagnetic Spectrum High Freq High E Low Freq Low E 9 Wavelength & Frequency •Speed of light = 3.00 x 108 m/s •Wavelength = distance between waves •Frequency = # waves that pass point in given time c c = as increases, decreases as increases, decreases (slinky analogy) 10 11 Photoelectric Effect •Emission of electrons from a metal when light shines on metal •Light of minimum frequency was required to eject electron from metal •Problem-wave theory of light predicted light of any frequency could eject electron 12 Is Light a Particle? •Max Planck-hot objects emit energy not as waves, but in packets called “quanta” •“quantum”-minimum energy that can be gained or lost by an atom •E = h •E = energy •h = Planck’s constant (6.626 x 10-34 Js) • = frequency (in s-1) 13 Light as a Wave & Particle •Einstein’s radical idea! •Dual wave-particle nature of light •Called particles of light “photons” (no mass) •Explained photoelectric effect (photon must have minimum energy) •Light is absorbed only in whole numbers of photons 14 15 Emission Spectra •White light -continuous spectrum •Hydrogen atoms -line-emission spectrum •Big Question -Why did hydrogen atoms only give off specific frequency (colors) of light? 16 The “Fingerprints” of Atoms 17 Atomic Spectra • The atom is quantized, i.e. only certain energies are allowed. 18 Neils Bohr Solved the Mystery • Electron circled nucleus in “orbit” of fixed energy • Absorption-electrons can “hop” from ground state to excited state • Emission-when electrons “fall” from excited state to ground state • Energy difference corresponds to hydrogen’s spectral lines • Ladder analogy 19 20 Remaining Mystery • Bohr’s model only explained hydrogen atom • What about other elements with more than one electron? 21