Modern Atomic Theory Notes

Electromagnetic radiation – energy that travels through space as waves. Waves have three primary characteristics: Wavelength (- lambda) – distance between two consecutive peaks or troughs in a wave. Unit = meter Frequency (  = nu) – indicates how many waves pass a given point per second. Unit = Hertz (Hz) Speed – velocity (c = speed of light = 3 x 108 m/sec) - indicates how fast a given peak moves in a unit of time c = 

Tell me what you know… Which wave has the greatest wavelength? Which wave has the greatest frequency? Which wave has the greatest speed?

Electromagnetic radiation (light) is divided into various classes according to wavelength.

Tell me what you know… Which color has the greatest wavelength? Which color has the shortest wavelength?

Wave- Particle Theory – Light as waves & Light as photons Photon/quantum – packet of energy OR a “particle” of electromagnetic radiation

Change in Energy of a photon = (Planck’s Constant) x (frequency) Energy - (E – change in energy) – Unit Joules (J) Planck’s Constant – (h = 6.626 x 10-34 J * s) Ephoton = h Change in Energy of a photon = (Planck’s Constant) x (frequency) c =  & Ephoton = h → Ephoton = hc  Ex: What is the wavelength of light with a frequency of 6.5 x 1014 Hz? What is the change in Energy of the photon? E = hc  ΔE = 4.3 x 10-19 J Given = 6.5 x 1014 Hz  = ? ΔE = ? = c  = 3 x 108 m/sec 6.5 x 1014 Hz = (6.626 x 10-34 J x s)(3 x 108 m/s) 4.6 x 10-7 m λ = 4.6 x 10-7 m

Tell me what you know… What does ΔE tell us about a photon?

Excited State – atom with excess energy Ground State – lowest possible energy state Wavelengths of light carry different amounts of energy per photon Only certain types of photons are produced (see only certain colors) Quantized – only certain energy levels (and therefore colors) are allowed

Emission and Absorption Spectra Emission Spectrum – bright lines on a dark background. Produced as excited electrons return to a ground state – as in flame tests. Absorption Spectrum – dark lines in a continuous spectrum. Produced as electrons absorb energy to move into an excited state, only certain allowable transitions can be made. Energy absorbed corresponds to the increase in potential energy needed to move the electron into allowed higher energy levels. The frequencies absorbed by each substance are unique. Nucleus Nucleus

Tell me what you know… How can we tell elements apart using emission spectra?

Bohr Model – suggested that electrons move around the nucleus in circular orbits Only Correct for Hydrogen Wave Mechanical Model – Described by orbitals gives no information about when the electron occupies a certain point in space or how it moves *aka – Heisenberg's Uncertainty Principle

Parts of the Wave Mechanical Model 1. Principle Energy Level (n) – energy level designated by numbers 1-7. -called principle quantum numbers 1 2 3 4 5 6 7 2. Sublevel – exist within each principle energy level -the energy within an energy level is slightly different -each electron in a given sublevel has the same energy -lowest sublevel = s, then p, then d, then f s p d f

Tell me what you know… Write the sublevels in order of highest energy to lowest energy.

Parts of the Wave Mechanical Model cont. 3. Orbital – region within a sublevel or energy level where electrons can be found s sublevel – 1 orbital p sublevel – 3 orbitals d sublevel – 5 orbitals f sublevel – 7 orbitals - ** No more than two electrons can occupy an orbital** -an orbital can be empty, half-filled, filled

Tell me what you know… How many total electrons are in each sublevel (s, p, d, & f)?

Sulfur = 1s2 2s2 2p6 3s2 3p4 Cd = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 Electron Configuration – arrangement of the electrons among the various orbitals of the atom Ex: 1s22s22p6 = Neon Sulfur = 1s2 2s2 2p6 3s2 3p4 Cd = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 Na = 1s2 2s2 2p6 3s1 Ne Na

Principle Energy Level Summary Principle Energy Level # of sublevels # of orbitals present s p d f Total # of orbitals Maximum # of electrons 1 1 1 - - - 1 2 2 2 1 3 - - 4 8 3 3 1 3 5 - 9 18 4 4 1 3 5 7 16 32

Shapes of orbitals All s orbitals are spherical as the principle energy level increases the diameter increases. All p orbitals are dumbbell or figure-8 shaped – all have the same size and shape within an energy level

4 of the d orbitals are 4-leaf clover shaped and the last is a figure-8 with a donut – all have the same size and shape within an energy level

f orbitals are complicated!!!!!

Electron Spin Spin – motion that resembles earth rotating on its axis– clockwise or counterclockwise Pauli Exclusion Principle – two electrons in the same orbital must have opposite spins Hund’s Rule – All orbitals within a sublevel must contain at least one electron before any orbital can have two Orbital Diagram – describes the placement of electrons in orbitals use arrows to represent electrons with spin line represents orbital (s=1, p=3, d=5, f=7) ____ full ____ half-full ____ empty

Orbital Diagrams Neon = 1s__ 2s__ 2p__ __ __ Ex: Carbon = 1s__ 2s__ Zinc = 1s__ 2s__ 2p__ __ __ 3s__ 3p__ __ __ 4s__ 3d__ __ __ __ __ Gallium = 1s__ 2s__ 2p__ __ __ 3s__ 3p__ __ __ 4s__ 3d__ __ __ __ __ 4p__ __ __

Tell me what you know… Summarize each in four words or less: Spin Pauli Exclusion Principle Hund’s Rule

Aufbau Order Aufbau Order – Tool to predict the order in which sublevels will fill OR use order on Periodic Table

Core Electrons – innermost electrons – not involved in bonding Noble Gas Configuration – Shorthand configuration that substitutes a noble gas for electrons Ex: Valence Electrons – Electrons in the outermost (highest) principle energy level in an atom, electrons use in bonding Core Electrons – innermost electrons – not involved in bonding Valence Configuration – shows just the valence electrons Na = 1s22s22p63s1 or [Ne]3s1 Sn = 1s22s22p63s23p64s23d104p65s24d105p2 or [Kr]5s24d105p2 Na = 3s1 3rd Shell/1valence electron Sn = 5s25p2 5th Shell/4 valence electrons