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Chem 105 Ch 3

Electromagnetic radiation any form of radiant energy in the electromagnetic spectrum
Electromagnetic spectrum spectrum of electromagnetic energy
Wavelength the distance from crest to crest or trough to trough
Frequency number of crests of a wave that pass a stationary point of reference per second
Amplitude height of wave from crest to midpoint or trough to midpoint
Atomic emission spectra characteristic patterns of bright lines on black background
Atomic absorption spectra characteristic patterns of dark lines produced when an external source of radiation passes through free, gaseous atoms
Blackbody radiation anything that acts as a perfect absorber and perfect emitter, wavelength depends on T
Quantum/quantized having values to whole number multiples of a specific base values, stairs
E = hν = hc/λ energy formula
Planck's constant 6.626x10^-34
Photons quanta of electromagnetic radiation
work function amount of energy needed to dislodge an election from the surface of a material
Photoelectric effect shows light has both wave and particle like waves
Atomic line spectra all lines
Bohr model not a valid but it explains quantum energy levels and nucleus concept
Rydberg equation used to find energy needed to move one electron to a different level
emission giving energy
absorption taking in energy
Energy level diagram showing the energy to levels
Ground state n=1
Excited states n>1
Electron transitions movement from one level to another
De Broglie equation λ=h/mu
Standing wave/nodes where there is not waves, no amplitude
Heisenberg uncertainty principle can't know position and momentum Δx*mΔu≥h/4π
Quantum mechanics/wave mechanics mathematic description of the wavelike bejavior of electrons and other particles
Schrödinger wave equation description of how electron matter wave varies with location and time around the nucleus
Wave functions solution to Schrödinger wave equation, define how varies both time and location
Quantum numbers describe electrons
Principal n energy level
Angular momentum n-1, s, p, d, f, g
Magnetic from +l to -l
Spin positive one half or negative
Pauli exclusion principle no two electrons can have the same four quantum numbers
Atomic orbitals the areas of propable electron location according to energy level
s, p, d orbital names
Nodes curved planar nodes, point node
Aufbau principle fill electrons from lowest to highest
Electron configurations/condensed electron configurations 1s2s2p3s3p4s3d4p… [Ne]3s3p
Atoms smallest unit of an element that still has the properties of that element
Anomalies cromium group and copper group
Ions usually corresponding to group number, transition metals usually 2+ or 3+
Core electrons electrons under the valence shell
Valence electrons not in core, anything higher than nearest noble gas
Valence shell outer shell where valence electrons probable location is
Hund’s rule fill all positive spin before negative spin
Orbital diagrams electron configuration with the little arrows representing spin
Isoelectronic same number of electrons (ions with noble gases)
Periodic trends trends presented by periodic table
Atomic radii decreases going left to right on a period, increases going down groups
Ionic radii cations < neutral, anions > neutral
Ionization energy increases going left to right on a period, decreases going down groups (anomalies between groups 2 and 13 and between N and O)
Electron affinity energy to add electron decreases as you go across a period
infinity true radius of an atom
Created by: ksheehy96



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