Electron Configuration

The Electromagnetic Spectrum Know what is in the red boxes High frequency Short wavelength High energy lower frequency longer wavelength lower energy

Jumping Electrons normally electrons exist in the ground state, meaning they are as close to the nucleus as possible when an electron is excited by adding energy to an atom, the electron will absorb energy and "jump" to a higher energy level heating a chemical with a Bunsen burner is enough energy to do this

Emission line spectrum energy is applied to a specific element this “excites” the element visible light, infrared, and UV can be emitted when electrons fall back down to the ground state unique for every element and are used to identify atoms (much like fingerprints are used to identify people) http://www.youtube.com/watch?v=QI50GBUJ48s

Give off energy when falls back down to ground energy level More on emission line spectrum Give off energy when falls back down to ground energy level

this is a repetitive slide- just couldn’t bear to delete it an electron in the atom gains (absorbs) energy from heating electron jumps up an energy level. electron is now unstable (unwelcome) in this level and is “kicked out” when the electron loses the energy and come back to the original level, light is emitted

Scandium 3-D video (2:31) 3-D Graphic Examples of Atomic Orbitals

Quantum Numbers (however, actual numbers are often not used) each electron in an atom is described by four different quantum numbers think of the 4 quantum numbers as the address of an electron… country > state > city > street electrons fill low energy orbitals before they fill higher energy ones

Principle quantum number (n) Quick intro, more later. Principle quantum number (n) describes the SIZE of the orbital or ENERGY LEVEL (shell) of the atom. Angular quantum number (l) a SUB-LEVEL (shell) that describes the type or SHAPE of the orbital Magnetic quantum number (m) the NUMBER of orbitals describes an orbital's ORIENTATION in space Spin quantum number (s) describes the SPIN or direction (clockwise or counter-clockwise) in which an electron spins

Principle Quantum # (n) LEVEL/SIZE 1 2 3 4 Angular Quantum # (l) ORBITAL SHAPE or SUBLEVEL s s p s p d s p d f Magnetic Quantum # (m) AXIS/ ORIENTATION or ORBITALS 1 orbital 1 3 4 total orbitals 1 3 5 9 total orbitals 1 3 5 7 16 total orbitals Spin Quantum # (s) DIRECTION OF ELECTRON SPIN 2 e- 8 e- 18 e- 32 e-

4f 4d 4p 4s 14 (7) 10 (5) 6 (3) 2 (1) 32 = level and sub-level = max. # of electrons = # of electrons = number of orbitals 3d 3p 3s 10 (5) 6 (3) 2 (1) 18 2p 2s 6 (3) 2 (1) 8 1s 2 (1) 2

Principle Quantum Number (n) or Energy Level values 1-7 used to specify the level the electron is in describes how far away from the nucleus the electron level is the lower the number, the closer the level is to the atom's nucleus and less energy maximum # of electrons that can fit in an energy level is given by formula 2n2 (n = energy level)

Angular Quantum Number (l) or Sub-Levels determines the shape of the sub-level number of sub-levels equal the level number ex. the second level has two sub-levels, the third has three and so on… they are numbered but are also given letters referring to the sub-level type l=0 refers to the s sub-level l=1 refers to the p sub-level l=2 refers to the d sub-level l=3 refers to the f sub-level just know this

Magnetic quantum number (m) or Orbitals Electron Orbitals YouTube 1:37 the third of a set of quantum numbers tells us how many sub-levels there are of a particular type and their orientation in space of a particular sub-level only two electrons can fit in an orbital = electron

S sub-level has only 1 orbital only holds two electrons

P sub-level has 3 orbitals holds up to six electrons

D sub-level has 5 orbitals holds up to 10 electrons

F sub-level has 7 orbitals holds up to 14 electrons

Spin quantum number (s) the fourth of a set of quantum numbers number specifying the direction of the spin of an electron around its own axis. only two electrons of opposite spin may occupy an orbit the only possible values of a spin quantum number are +1/2 or -1/2.

Principle Quantum # (n) LEVEL/SIZE 1 2 3 4 Angular Quantum # (l) ORBITAL SHAPE or SUBLEVEL s s p s p d s p d f Magnetic Quantum # (m) AXIS/ ORIENTATION or ORBITALS 1 orbital 1 3 4 total orbitals 1 3 5 9 total orbitals 1 3 5 7 16 total orbitals Spin Quantum # (s) DIRECTION OF ELECTRON SPIN 2 e- 8 e- 18 e- 32 e-

Principle energy level (n) Type of sublevel Table 3-6b Orbitals and Electron Capacity of the First Four Principle Energy Levels Principle energy level (n) Type of sublevel Number of orbitals per type Number of orbitals per level(n2) Maximum number of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7

“Rules” for Writing Electron Configurations a method of writing where electrons are found in various orbitals around the nuclei of atoms. three rules in order to determine this: Aufbau principle Pauli exclusion principle Hund’s rule

Aufbau Principle electrons occupy the orbitals of the lowest energy first each written represents an atomic orbital (such as or or or ….) electrons in the same sublevel/shell have equal energy ( same energy as ) principle energy levels/shells (1,2,3,4..) can overlap one another ex: 4s orbital has less energy than a 3d orbital

Pauli Exclusion Principle Hamster video 1:00 only two electrons in an orbital must have opposite spins represents one electron represents two electrons in an orbital actually incorrect as well, see next slide

Hund’s Rules every orbital in a subshell must have one electron before any one orbital has two electrons all electrons in singly occupied orbitals have the same spin.

Writing Orbital Diagrams

Energy

Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) s orbitals d orbitals p orbitals f orbitals

Boron Atomic # 5 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

Boron ion (3+) Atomic # 5 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

Neon Atomic # 10 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

Bromine Atomic # 35 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

Bromine ion (1-) Atomic # 35 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

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Orbital diagrams Electron Configurations 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 1s2 2s2 2p6 3s2 3p1 1s2 2s2 2p6 3s2 3p2 1s2 2s2 2p6 3s2 3p3 1s2 2s2 2p6 3s2 3p4 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6 Na Mg Al Si P S Cl Ar 1s 2s 2p 3s 3p Orbital diagrams Electron Configurations

Electron Configurations 2p4 Number of electrons in the sublevel Energy Level Sublevel 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

Writing Electron Configurations To write out the electron configuration of an atom: use the principal quantum number/energy level (1,2,3, or 4…) use the letter term for each sub-level (s,p,d, or f); don’t worry about orientation such as x,y,z axis but you do have to be able to draw these for IB use a superscript number indicates how many electrons are present in each sub-level hydrogen =1s1. Lithium =1s22s1. don’t write anything for spin

Sometimes levels are switched in order to keep the level together. Order of Electrons Sometimes levels are switched in order to keep the level together. I hate when they do that! 4s requires less energy and I think it should be before 3d. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14… electron configuration video (3:24)

exceptions (don’t need to know this, just be aware that there are exceptions) orbitals “like” to be empty, half filled, or full therefore, an electron leaves the 4s (leaving it half full) and goes to the 3d in order to make it full Cr we would predict: 1s2 2s2 2p6 3s2 3p6 4s2 3d4 but it is actually: 1s2 2s2 2p6 3s2 3p6 4s13d5 Cu 1s2 2s2 2p6 3s2 3p6 4s2 3d9 1s2 2s2 2p6 3s2 3p6  4s1 3d10

Noble Gas Shortcut same