Chapter 4 Electron Configurations and Quantum Chemistry

Chapter 4Electron Configurations and Quantum Chemistry Electron configurations determine how an atom behaves in bonding with other atoms! Topics rearranged from your text. You should read pages 90-116. Atomic Emissions/Abortions deleted Anyone who says that they can contemplate quantum mechanics without becoming dizzy has not understood the concept in the least. -Niels Bohr

The Bohr Model • Niels Bohr • rebuilt the model of the atom placing the electrons in energy levels. • Quantum chemistry • a discipline that states that energy can be given off in small packets or quanta of specific size. • What would happen to an electron if the right sized quanta of energy was added to it? • What would happen when the electron came back down to its ground state? EXCITED STATE Ground state

Atomic Absorption / Emission • Atomic Absorption • Electron is given a specific amount of energy, it “quantum jumps” to a higher energy level (excited state). • Atomic Emission • Electron returns to its ground state, it emits energy equal to the amount of energy required to raise it to the higher level. • The difference in energy of the levels produces photons of differing energy • (blue = higher energy, red = lower energy) Internet Animation of Atomic Absorption / Emission

Atomic Absorption / Emission • White light is composed of all visible wavelengths (colors) of light. (Electromagnetic Radiation) • The continuous spectrum • having all wavelengths / colors of visible light (white)

Emissions Spectrum What happens next? • When electrons in a gas are charged with energy • those electrons make their quantum jumps… • Then return to their ground states… • resulting (unique) pattern of light given off is called an emission spectrum. • Each element has a different emission spectrum, a sort of quantum fingerprint, that we can use to identify elements in unknown samples.

Some bright-line emission spectra: • Some emission spectra of elements • from www.webelements.com: • Sodium, 11 electrons in 3 ground-shells: • Mercury, 80 electrons in 6 ground-shells:

Absorption Spectrum • Absorption spectrum • Given off when white light is shined through a gas • electrons in the gas absorb some wavelengths of the light • Viewing light from distant stars through the gasses of our planets allows us to know what chemicals make up those planets.

Bright-Line Spectra • Max Planck, German physicist • calculated that energy required to make a quantum jump for specific energy levels and colors. • Planck’s Constant, 6.6262x10-27 • multiplied by the frequency of the desired emission color equals energy required to produce that jump.

Electron Configurations - overview • Bohr model • electrons exist in specific energy levels. • Electron orbitals (shapes) • Within each energy level, orbits the electrons can occupy. • Within each orbital • electrons can be set “spin up” or “spin down” • Electron configuration • The configuration of electrons in their levels, orbitals, and spins gives us an atom’s. • Modern Quantum Model • Electron exists in electron configurations

Energy Levels (n) Old School: “KLM notation” • The electrons exist in energy levels or shells. • The first energy shell can hold only 2 electrons. • Hydrogen and Helium in their ground state have electrons that occupy this shell. • The second shell can hold 8 electrons. • The third can hold 18 electrons. 2 8 18 32 Shells All shells after three can hold 32 electrons.

Orbitals (Shapes) • Orbitals • electrons travel in set paths. • These paths form shapes. • Each “shape” can hold 2 electrons • The smallest orbital is the “s” orbital. The “s” orbital: • Has only 1 shape (holds 2 e-) • Is spherical in shape • Is the lowest energy orbital S-2

p-Orbitals • The 2nd orbital shape is the “p” orbital shape. • There are 3 “p” shapes, each holding 2 electrons, for a total of 6 electrons in the “p” orbital. • The “p” orbitals are: • Dumbbell-shaped • Higher in energy than the “s” S-2 P-6

d-Orbitals • The 3rd orbital shape is the “d” orbital shape. • There are 5 “d” orbital shapes, for a total of 10 electrons in the “d” orbital. • “d” orbitals are higher in energy than “p” orbitals. S-2 P-6 D-10

f-Orbitals • The last orbital shape is the “ f ” orbital shape. • “ f ” orbitals have irregular shapes due to quantum tunneling. • There are 7 “ f ” shapes, for a total of 14 electrons. Electrons in f orbitals are very high in energy F-14 S-2 P-6 D-10

Electron “Spin” • Electrons can be “spin up” or “spin down.” • (by convention, an electron that is alone is “spin up”) • Pauli’s Exclusion Principle • If two electrons share a shape, they must be spin-paired (one upandone down). • Hund’s Rule • As electrons fill orbitals, they first fill each shape available with one electron before spin pairing. • For instance: take a “p” orbital…it has three orbital-shapes that can hold 2 e- each. • It would fill like this:

Writing Electron Configurations s  low energy • The Aufbau principle • electron will fill lower energy orbitals first. • Energy of electrons: • low energys < p < d < fhigh energy • low energynearer < fartherhigh energy • low energylevel 1 < level 7high energy • Total energy of an electron: • Product of energy of its shell and the energy of its orbital. • Guess: Which is lower in energy, an electron found in 3d or one found in 4s? d  high energy close  low energy far  high energy Total energy = Shell x orbital shape The 4s electrons are lower in energy!

Writing Electron Configurations • Orbital filling diagram • Shows how electrons fill into levels and orbitals  Don’t Copy this

Building the Orbital Filling Diagram • Begin by listing the shells 1, 2, 3, 4, 5, 6, 7 vertically. • These are your “s” orbitals. • Next, add another column of number, beginning with 2. • These are your “p” orbitals. • Do the same for “d” and “f” orbitals, beginning with “3” for the “d” orbitals and “4” for the “f” orbitals. • Next, add your orbital letters. • Finally, draw diagonal lines as shown. 1 s 2 s 2 p 3 s 3 p 3 d 4 s 4 p 4 d 4 f s p d f 5 5 5 5 6 s 6 p 6 d 6 f 7 s 7 p 7 d 7 f s p d f

Electron Configurations of Some Atoms • Consider Fluorine, with 9 electrons • What about Copper, with 29 electrons? Notice the position of the last electron… Both used

“Blocks” of the periodic table… • The periodic table tells us in which orbital the last electron should be found. • The last electron in an atom is found in the… p orbitals s orbitals d orbitals f orbitals

Noble Gas Shorthand • Notice the configurations of the noble gases: • We can shorten the electron configuration of larger elements with NGS. • Consider Mg: • We can substitute Neon’s e- config, and write Mg: • Similarly, Titanium’s (Ti) e- config: • Can be shortened to:

Ion e- configurations • Ions(elements with more/less electrons) also have electron configurations. • Consider Sulfur (S): • What if sulfur gained an electron? • Consider Calcium (Ca): • What if calcium lost two electrons?

End of chapter 4 • Question: • Why do the atomic radii of atoms decrease as electrons are added to the atom, as you move from left to right across a period? • electrostatic attraction • attraction between the electrons (-) in the shells and the protons(+) in the nucleus – pulls the electrons in This is what we call a periodic trend

Chapter 4 Electron Configurations and Quantum Chemistry