Ch. 6 Electronic Structure of Atoms- Electronic structure: The arrangements of electrons in atomso Number of electrons in the atom as well as distribution around the nucleus and their energies6.1 The Wave Nature of Light- Electromagnetic radiation: Also known as radiant energy, light emitted or absorbed by substanceso All types move through vacuum at 2.998 x 108 m/s, speed of light- Periodic: Pattern of peaks and troughs repeats itself at regular intervalso Wavelength: Distance btwn two adjacent peakso Frequency: Number of complete wavelengths, or cycles, that pass a given point each second (units are hertz)- Inverse relationship between frequency and wavelength of EM radiationo λv= c (wavelength x frequency= speed of light)- Electromagnetic spectrum: Display showing various times of EM radiation arranged in order of increasing wavelength6.2 Quantized Energy and Photons- The emission of light from hot objects (blackbody radiation)- The emission of electrons from metal surfaces on which light shines (photoelectric effect)- The emission of light from electronically excited gas atoms (emission spectra)Hot Objects and the Quantization of Energy:- Max Planck (1858- 1947)o German physicisto 1900, proposed energy can be released or absorbed by atoms only in discrete “chunks”o Quantum: (“fixed amount”) Smallest quantity of energy that can be emitted or absorbedas EM radiationo E= hv E (energy of single quantum), v (frequency) Planck constant: 6.626 x 10-34 J-so Energy emitted or absorbed only in whole-number multiples of hvo Allowed energies are quantized- values restricted to certain quantitieso 1918 Nobel Prize in Physics for work on quantum theoryThe Photoelectric Effect and Photons:- Albert Einstein (1879- 1955)o Photoelectric effect: 1905, use of Planck’s theory to explain electrons emitted from metal surfaces Photon: Packet or particle of energy Energy of photon= E= hv Radiant energy is quantized Photons can transfer energy to electrons in metal Work function: Amount of energy required for electrons to overcome attractive forces holding them in metal Intensity (brightness) of light related to number of photons striking, but not E of each photon Nobel Prize in 1921 for photoelectric effect6.3 Line Spectra and the Bohr Model- 1913, Danish physicist Niels Bohr explained line spectra, using ideas of Planck and EinsteinLine Spectra:- Monochromatic: Radiation composed of a single wavelength- Polychromatic: Radiation composed of many different wavelengths (most sources)- Spectrum: Display of radiation separated into component wavelengths- Continuous spectrum: Rainbow of colors, containing light of all wavelengths- Ex. Red-orange glow from neon, yellow from sodium- Line spectrum: Spectrum containing radiation of only specific wavelengthso Line spectrum of H found in mid-1800s 1885, Swiss Johann Balmer found equation relating wavelengths to integerso Balmer equation extended to Rydberg equation- λ (wavelength), RH (Rydberg constant 1.096776 x 107 m-1), n1 and n2 (positive integers)Bohr’s Model:- Bohr assumed electrons in H atoms move in circular orbit around nucleus, but electron does not lose energy as it should.- Model based on three postulateso Only orbits of certain radii, corresponding to certain specific energies, are permitted for the electron in a H atom.o An electron in a permitted orbit is in an “allowed” energy state. It does not radiate energy and does not spiral into nucleus. o Energy emitted or absorbed by electron only as electron changes from one allowed energy state to another. Energy E= hvThe Energy States of the Hydrogen Atom:- Energies corresponding to allowed orbits for electron in H atomoo Principal quantum number: Integer, n, which can have whole-number values- Radius of orbit gets larger as n increases- The lower (most neg.) the energy is, the more stable the atom is. - Ground state: Lowest-energy state (n= 1)- Excited state: Higher-energy state (n=2 or higher)- Reference: Zero-energy, state in which electron is completely separated from nucleus- Energy absorbed to move to higher energy state and emitted when electron jumps to lower energy state-- The energy of the photon (Ephoton) must equal the difference in energy btwn the two states (ΔE)--Limitations of the Bohr Model:- Electrons exist only in certain discrete energy levels, which are described by quantum numbers.- Energy is involved in the transition of an electron from one level to another.6.4 The Wave Behavior of Matter- Louis de Broglie (1892- 1987)o French physicistoo Momentum: Quantity mv for any objecto Matter waves: Wave characteristics of material particleso Any object of mass and velocity would give rise to a characteristic matter wave. - X-ray diffraction: Interference pattern characteristic of wavelength properties of EM radiation that forms when X rays pass through a crystalThe Uncertainty Principle:- Werner Heisenbergo German physicisto Uncertainty principle: Impossible to simultaneously know both the exact momentum of the electron and its exact location in spaceo6.5 Quantum Mechanic and Atomic Orbitals- Erwin Schrodinger (1887- 1961)o Schrodinger’s wave equation: Incorporates both wave-like and particle-like behaviors of electrono Standing waves: Waves that do not travel in