Chpt 7 - Atomic Structure Electromagnetic Radiation Atomic Spectrum - Bohr Model Quantum Mechanical Model Orbital Shapes and Energies Electronic Structure & Periodic Table Periodic Trends HW: Chpt 7 - pg. 329-338, #s 23-27, 37-43, 54, 62, 65, 67-68, 70, 71, 74, 76, 82, 86, 98, 100, 102, 104, 115, 116, 119, 126

Electromagnetic Radiation λ (lamba) = wavelength (m) ν (nu) = frequency (Hertz, Hz or s-1) E = energy c = speed of light, 2.9979 x 108 m/s c = λν they are inversely related Know the relative order of radiation in E, λ, ν

1900s Death of Classical Physics Black Body Radiation Planck’s hypothesis… energy is quantized E = hν or ΔE = nhν n = integer h = 6.626x10-34 J.s Photoelectric effect Einstein proposed EM radiation is quantized A stream of “particles” called photons E = hν = hc/λ deBroglie λ = h/mv (wavelength of a particle) velocity in m/s mass in kg - so units cancel with J

Photoelectric Effect Light with frequency lower than a specific threshold have no electrons emitted (no matter how intense it is) Light with frequency greater than threshold emits electrons and number of electrons increases with intensity

Diffraction Pattern in a Crystal Electron beam is diffracted off of a crystal. Electron exhibits wave behavior!!! Davisson Germer experiment - They shared Nobel prize with GP Thomson which did similar type experiment.

Continuous vs Discrete Spectrum Continuous spectrum vs. discrete spectrum (line spectrum) Absorption vs emission spectrum Only certain energies are allowed for the electrons in any atom

Hydrogen Atom The observed spectrum was explained by Bohr by proposing the electrons move around the nucleus in certain allowed circular orbits.

Bohr Energy Expression Calculated from hydrogen atom spectrum E = -2.178x10-18 J (Z2/n2) Z = atomic number, 1 for hydrogen n = orbital that the electron is located ultimately only good for hydrogen atom spectrum

Quantum Mechanics Schrodinger solved the problem mathematically (no real physical significance) treating electrons as waves. Hψ = Eψ ψ is the wave function of the electron’s coordinates in 3 dimensions Heisenberg - uncertainty principle Δx * Δ(mv) >= h/4π position momentum See Heisenberg laser slit video

Orbital shapes and Energies Orbitals are simply then a probability distribution of where the electron could be found. (left) probability function for s-orbital (below) Radial probability function for s-orbital

Shapes of p and d orbitals

What can we know about electron? 4 Quantum numbers describe the electron in an orbital. n is principle quantum number - relates to size of the orbital, n = 1, 2, 3, 4,… l is angular momentum q.n. - relates to shape of orbital, l = 0, 1, 2, …, n - 1 s-orbital is l = 0 p-orbital is l = 1 d-orbital is l = 2 f-orbital is l = 3 ml is magnetic q.n. - relates to orientation in space ml = -l,…,0,…, +l ms is electron spin q.n. - relates to spin of electron ms = - 1/2 or +1/2 (called spin up & spin down or clockwise/counter clockwise)

Quantum numbers Examples of valid quantum numbers for various orbitals. In addition, spin +/- 1/2 for each individual orbital.

Energy Levels of orbitals As we keep adding energy levels, we see as the principle quantum number, n, increases the number of sublevels (types of orbitals) increases. In addition the energy spacings get closer together 1s - 2s - 3s - 4s - etc. So the energy of the 4s orbital comes lower than the 3d. The order need not be memorized because the elements in Periodic Table shows it with its s,p,d,f blocks.

Electron Configuration rules Electron’s occupy lowest energy level first - aufbau principle Maximum of 2 electrons in any orbital - Pauli exclusion principle If 2 electrons occupy the same orbital they have opposite spins. +1/2 or -1/2 also called spin up / down or clockwise / counter-clockwise For degenerate orbitals (the same energy like the three p, five d, or seven f) use Hund’s rule, also known as the bus rule - only pair up the electrons if necessary.

General s,p,d,f blocks The periodic table clearly shows that after the 3p orbital, the 4s fills before the 3d. Likewise, 6s 4f 5d 6p is the order when the lanthanides start.

Electron Configurations A couple of exceptions Cr and Cu groups in the transition metals promote an s electron to achieve a half-filled and fully-filled set of d-orbitals because they have more stability.

Mendeleev’s Original Periodic Table Organized by increasing atomic mass and put in columns by similar properties and reactivities Left spaces for undiscovered elements together with predicted properties - these were confirmed by experimental results!!!

Periodic Table Trends some are found in Chpt 8 Ionization Energy Electron Affinity Atomic Radius Ionic Radius Electronegativity

Ionization Energies The ionization energy is the energy necessary to remove an electron completely from an atom. X --> X-1 + e- The 2nd ionization energy would remove the next electron, etc. Notice the trends in this chart a) across the period - general and detailed b) 1st ion E, 2nd ion E, etc. large jumps associated with core electrons.

1st Ionization Energy Chart

Electron Affinity EA is the energy change with adding an electron to an atom X + e- --> X-1 This energy is correlated to thermodynamics, thus atoms that have a high EA (like to gain e-) the associated E change is negative (exothermic) the higher EA the more exo it is. Generally, EA increases up a group and across a period.

Atomic Radius Radii are estimated from actual spacing in metals or molecules

Ionic Radius Trends Ionic radius of most common ion reported in picometers. The size typically decrease across the period with a large jump when going from anion to cation. Also, cations are smaller than their atoms and anions are larger than their atoms.

Electronegativity Trends Electronegativity is the ability of an atom to attract electrons to itself in a chemical bond. It generally increases across a period and decreases down a group.

Alkali Metals Periodicity The alkali metals are shown below with various physical properties. These are expected trends for other groups of metals as well.