Electron Energy Level Notes

Electron Energy Level Notes Electrons do not travel around the nucleus of an atom in orbits They are found in energy levels at different distances away from the nucleus. (kind of like shells or layers).

Hydrogen Atomic Orbitals Electrons cannot exist between energy levels (just like the rungs of a ladder). Principal quantum number (n) indicates the relative size and energy of atomic orbitals. n specifies the atom’s major energy levels, called the principal energy levels.

Electron Energy Level Notes Energy levels are broken up into sublevels: There are at least 4 possible types of sublevels—given labels: s, p, d, or f

Hydrogen Atomic Orbitals (cont.) Energy sublevels are contained within the principal energy levels.

Electron Energy Level Notes In each energy level, electrons fill sublevels in a certain order Level 1: only has one s sublevel (a spherical shape) 2 electrons may fit in this sublevel--each one has an opposite “spin”, allowing them to take up the same space Pauli exclusion principle—no more than 2 electrons may be found in the same orbital (“orbital” means a particular location)

Electron Energy Level Notes has two sublevels: s and p 2 electrons in s there are 3 different p orbitals, and may hold 2 electrons each—6 total. total of 8 overall in Level 2

Electron Energy Level Notes has 3 sublevels: s, p, and d 2 electrons max in s 6 electrons max in p there are 5 different d orbitals, and 2 electrons can fit in each—total of 10. total of 18

Level 4: has 4 sublevels: s, p, d, and f 2 electrons max in s 6 electrons max in p 10 electrons max in d. 7 types of orbitals in f, each with 2 electrons = 14 electrons total of 32

Hydrogen Atomic Orbitals (cont.) Each energy sublevel relates to orbitals of different shape.

Aufbau Principal: electrons must enter into the lowest energy level first.

Electron Energy Level Notes An easy way to remember this is to use the periodic table--it is arranged to show how these orbitals are filled.

Order of Orbitals—Periodic Table

The aufbau principle states that each electron occupies the lowest energy orbital available.

Ground-State Electron Configuration (cont.) The Pauli exclusion principle states that a maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins. Hund’s rule states that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same energy level orbitals.

Transition metals Ground-State Electron Configuration (cont.) See below (and on p. 242 of your text) that the pattern you just learned is not so perfect: Any time a transition metal in row 4 and 5 of the periodic table has a chance to HALF FILL the D orbital, it will steal those electrons from the outer s orbital to do that. Once you hit row 5, it gets a bit more complicated. The general idea above still applies, but now you have to explain Ru, Rh, and Pd. This is due largely to the closeness of the d and s orbital energies. Once you get into the inner transition metals, and all elements in the d block AFTER the f block, the problem only gets compounded by the presence of f orbitals! Sc Ti V Cr Mn Fe Co Ni Cu Zn 4s23d1 4s23d2 4s23d3 4s13d5 4s23d5 4s23d6 4s23d7 4s23d8 4s13d10 4s23d10

Evidence for electron configs How do we know that all this information is true? What laboratory evidence is there to support all this? When electrons are mobile. When the “pump” up to higher energy levels, then fall back down into lower levels, they release light energy. This is call “emission spectra” (emission of color).

Hydrogen makes a good case study Hydrogen only has 1 electron, so we like to pump this guy full of energy and watch what happens as the electron slowly falls.

Moving electrons around

Xray pics of orbitals