Chapter 7: Quantum MechanicsClassical mechanics: laws describing the motion of macroscope objectsQuantum mechanics: principles of electrical and magnetic properties at the atomic and subatomic level, tied to the understanding of energy and matter. Wavelength: distance between crests. (Distance between tops) λFrequency: number of peaks per sound (how often waves come by) vλv=cHumans detect color with three different cone cells. (sensitivity in blue, green, red).Photoelectric effect:The PE requires a minimum of threshold frequency (vo) of light to occur.Metals exposed to light at or above the threshed frequency emit electrons.Einstein proposed that electromagnetic radiation (EMR) is a stream of particles (photons) Ephoton inversely proportional to vIntense EMR dislodges more electrons, it has more photons than dim EMR.Whether PE occurs, focus on frequency.v>>vov < voDim EMR Fewer electrons ejected higher velocityFewer electrons ejected No electrons ejectedBright EMR More electrons jected higher velocityNo electrons ejectedV<V0 = no electrons ejectedEnerg is quatized: it can only be released or absorbed in specific amounts, quanta (hv)Enegery per particle(for a single photon): Eproton=hv=hc/λΔEelectron= hv=hc/λ change of energy of an electron in an atom.Electrons can absorb a photon and move to a higher energy level.Evidence for electron transitions: atoms=line spectra, molecules=chemiluminscense.Electrons can jump from a lower quantum level to a higher one.Atoms- line spectra.When a pure element reflects llight and we refracts it, we only see certain light. The visible light is because the electron fell.Line spectra are not continuous spectrum because electron energies are quantized. Each line corresponds to an electronic transition.Four elements are names after prominent colors in their line spectra:Thalllium greek for sprout.Indium very strong indigo lineRubidium latin for deep redCesium latin for sky blueElectrons display various colors of light upon exposure to heat. Blue fireworks shouldn’t exist: all elements with able flame are toxic!In a given experiment, electrons act as either a particle or wave.50%- Electrons behave as particles because they have a mass50%- Electrons behave as waves because electrons ripple.λ=h/mvHeisenberg uncertainty Principle-An orbital is a volume of 3-D space with a probability of containing an electron. Δx*mΔv>/= h/4piΔx= uncertainty of particle positionmΔv=uncertainty of particle momentumv=velocityNiels Bohr: won 1922 Nobel Prize in Physic, son won it in 1975. Bohrium (Bh) named after him.Bohr Equation: (one electron system) E=-2.170*10-18J(Z2/n2)Z=nuclear charge=number of protons.N gives us orbital energy level and size. As n gets larger and larger the electron is further and further from nucleus. Largest value of n is 7. The higher the level the closer stacked.Lower energy jumping to higher= absorbing energy. Quantum numbers can be thought of as labels for an electron. Every electron in an atom has a unique set of four quantum numbers.The principal quantum number n corresponds to the shell in which the electron is located. Thus n can therefore be any integer. For example, an electron in the 2p subshell has a principal quantum number of n=2 because 2p is in the second shell.The angular momentum quantum number ℓ corresponds to the subshell inwhich the electron is located…ℓ can have integer values ranging from 0 ton−1ℓ= 1 2 3 4 5s p d f hRelative energies: ns < np < nd < nfThe magnetic quantum number mℓ corresponds to the orbital in which the electron is located. Instead of 2px,2py, and 2pz, the three 2p orbitals can be labeled −1, 0, and 1, but not necessarily respectively. As a rule, mℓ canhave integer values ranging from −ℓ to +ℓThe spin quantum number ms corresponds to the spin of the electron in the orbital. A value of 1/2 means an "up" spin, whereas −1/2 means a "down" spin.nlml90% inside orbital10% of the time electron is outside of the orbital.S orbitals are spherical and centered around the nucleus. As n gets bigger and bigger, orbital size and energy increase. Except for 1s, wavefunctions (_) change sign (+/-) at least once. The signs have nothing to do with charge. Nodes- planar or spherical surfaces where _ changes sign._=0 at nodeElectrons are never found at nodeFor any orbital, #nodes= (n-1)Radial distribution functions: probability electrons will be in certain areas. P orbitals, node is a plane. (x, y plane) dumbbell shaped, with node at nucleus.D orbitals, node is clover shaped. 5 orbitals ml= -2,-1,0,1,2Chapter 8Ground state: all electrons are in lowest energy orbitals possible.Electron configuration notation: for distribution of electrons in the orbitals of a ground state atom. Orbital diagram: electrons are arrow, orbitals are boxes.Cr: Half filled 4s, takes energy to move electron, but fills all shells stabilizing orbitals.Cu: promotes electron from 4s, but completely full 3d=stabilization.**Know electron configurations for main groups, and first transition series. Paulis exclusion principle1. In an atom no electrons have the same 4 quantum numbers (n,l,ml,ms).2. An orbital can hold 2 electrons and they must have opposite spins.s -1 orbital-2 electronsp -3 orbitals-6 electronsd -5 orbitals -10 electronsf -7 orbitals-14 electronsWhy is energy in a 1s orbital lower than in a 7s orbitals?n=size, electrons closer to nucleus at 1. The lower energy electrons experience less shielding.Periodic table trends: (relativity): Atomic radius, Ionic radius, ionization energy, electron affinity, metallic character, etc.Atomic radius: distance between nuclei and edge of atom. Increasing n=increasing sizeMove down: atomic radius increases., horizontally closer to Fr.Ion configuration/ion: Aufabau principle: energy shells fill from lowest to highest energy.Pauli Exclusion principle: orbitals hole 0,1, or 2 electrons if 2 they must have opposite spins.Hund’s rule: degenerate orbitals fill by placing one electron in each before completing pairs.Paramagnetic species have > 1 unpaired electron(s) and are attracted to magnetic fields.Diamagnetic species have no unpaired electrons and are slightly repelled by magnetic fields. to determine if a species is para or dia, consider its orbital diagram.Ionic radii increase going down a group. Larger positive charge= smaller cationLarger negative charge=larger anionCations are much smaller than their corresponding atoms.Anions are much larger than their corresponding