Electron Configurations & Quantum Numbers

Standards SC3. Students will use the modern atomic theory to explain the characteristics of atoms. b. Use the orbital configuration of neutral atoms to explain its effect on the atom’s chemical properties. f. Relate light emission and the movement of electrons to element identification.

To Know: Know colors of light (ROYGBV). Quantum numbers (POMS). Red-high wavelength; low frequency (energy) Violet- low wavelength; high frequency (energy) Quantum numbers (POMS). Sublevels (s,p,d,f)-shapes Relationship between energy level, sublevel, slots, electrons. Electron configurations, orbital notations, electron dot diagrams.

Electrons and Their Energies Bohr model was a one dimensional model describing the structure of the atom. Atoms are 3 dimensional. Quantum numbers describe the orientation of electrons in atoms. Exact locations of electrons can not be determined. Which energy level an electron is on is only a highly probable or likely location.

Facts about electrons 1. They can only move certain distances from the nucleus. 2. They can only move at certain speeds. 3. They give off (release) energy when moving to a lower level and gain energy (absorb) when moving to a higher level. 4, They cannot jump down to a level with a full octet of electrons.

Photon- a bundle or packet of energy in the form of electromagnetic radiation. When an atom gains photons, it becomes excited and the electrons move to a higher energy level. When it loses photons, it gives off photons, e- move to lower level. Atoms become excited ( and gain photons) by heating them or passing electricity through them. When they lose these photons, the electrons give off color, depending on amount of energy given off (related to frequency and wavelength).

The color depends on distance the electron travels, or wavelength. When electron falls back to lower energy level, it releases energy (photons) in the form of color. Elements release a specific color when heated in a flame that is determined by their electron configurations. Ex. Na- yellow-orange Ne- red

Quantum Numbers- Describe the energy and arrangement of electrons 1. (P) Principal Quantum Number (n) 2. (O) Orbital Quantum Number (l) 3. (M) Magnetic Quantum Number 4. (S) Spin Quantum Number (+1/2, -1/2)

Principal Quantum Number Represented by n. Describes the distance from the nucleus. Identifies the particular energy level. Same as the period number. 4 is farther away than 3, etc. A positive, whole number.

Orbital Quantum Number Identifies the sublevel (s,p,d,f). Each sublevel (orbital) has its own shape. s- spherical p- dumbbell d- clover f- ?

Orbital (shape) is dependent on principal quantum number 1st level has 1 possible shape s 2nd level has 2 possible shapes s or p 3rd level has 3 possible shapes  s, p, or d 4th level and up has 4 possible shapes s,p,d,f N tells us what energy level AND how many possible sublevels or shapes.

Magnetic Quantum number Tells how many positions on the sublevels; the orientation in space. (x, y, or z axis) Sublevel # positions # electrons s 1 2 p 3 6 d 5 10 f 7 14 Ex. The 2nd energy level can have s and p sublevels for a total of 4 possible positions. S = 1 p = 3 = 4

The total number of possible positions in an energy level = n2.

Spin Quantum number Electrons can spin clockwise or counterclockwise. Only 2 electrons can occupy each position, one spins clockwise, the other spins counterclockwise.

Examples The first energy level can only have “s” sublevel which has only one position. represents that these electrons have opposite spins. 1s max. of 2 e- on each level. The second energy level has s and p sublevels. 2s 2p The maximum number of electrons in a certain energy level = 2n2.

No 2 electrons can have the same set of quantum numbers. No 2 electrons can have exactly the same amount of energy if they come from the same atom.

4 n – energy level (principal) sublevel (Shape) Total slots n2 Total electrons 2(n2) 1 2 3 4

Light Equations Speed = wavelenght (nm) X frequency (Hz) Frequency = 1/wavelength