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, pages 90-116. Atomic Emissions/Abortions removed 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

Electron Configurations - overview • Bohr model • electrons exist in specific energy levels. • Electron orbitals (shapes) • Within each energy level, the 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. • 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, called orbitals. • 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” orbitals. • 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” orbitals. • “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

“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

Electron “Spin” • Electrons can be “spin up” or “spin down.” • (by convention, an electron that is alone is “spin up”) • Hund’s Rule • As electrons fill orbitals, they first fill each shape available with one electron before spin pairing. • Pauli’s Exclusion Principle • If two electrons share a shape, they must be spin-paired (one upandone down). • 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

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 two electrons? • Consider Calcium (Ca): • What if calcium lost two electrons?

Octets! • Octets: • Atoms with filled s and p orbitals in the same, highest level. • Have noble gas-like configurations • Have special stability • Both atoms and ions can have complete octets.

End of chapter 4 • Question: • Why do the atomic radii (size) of atoms decrease as electrons and protons 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

Electron Configurations and Quantum Chemistry