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Atomic Structure Electron Configuration. Scandium 3-D video (2:31). 3-D Graphic Examples of Atomic Orbitals. Topic 12.1.3 – 12.1.6. Review. Jumping Electrons. normally electrons exist in the ground state , meaning they are as close to the nucleus as possible
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Atomic StructureElectron Configuration Scandium 3-D video (2:31) 3-D Graphic Examples of Atomic Orbitals Topic 12.1.3 – 12.1.6
Review 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
Review • after a short time, this electron will spontaneously "fall" back to a lower energy level, giving off a quantum of light energy called a photon • the key to Bohr's theory was the fact that the electron could only "jump" and "fall" to precise energy levels, thus emitting a limited spectrum of light. • quantum is the amount of energy required to move an electron from one energy level to another
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 • the first three of these quantum numbers (n, l, and m) represent the three dimensions in which an electron could be found • the fourth quantum number (s) refers to a certain magnetic quality called spin
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
4f 4d 4p 4s 14(7) 10(5) 6(3) = level and sub-level = max. # of electrons = # of electrons = number of orbitals 2 (1) 32 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
Angular Quantum Number (l) or Sub-Levels • determines the shapeof the sub-level • number of sub-levels equal the level number • ex. the second level has two sub-levels • 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
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
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.
“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 oror ….) • 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 actually incorrect as well, see next slide • only two electrons in an orbital • must have opposite spins • represents one electron • represents two electrons in an orbital
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.
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
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