Electrons in Atoms Unit
Outline and Notes
I. Review Atomic models
a. Dalton — solid sphere
b. Thomson — Plum pudding
c. Rutherford — Planetary
d. Bohr — e
- have circular orbits of specific energy around the nucleus; energy is quantizede. Schroedinger — e
- travel in probability clouds outside the nucleus1) Never find e
- in nucleus2) Area of probability of finding e- called orbital; Heisenberg theory says you can know location of e
3) A quantum of energy is needed for an electron to move from one level to another
SO: this model is called the "Quantum Mechanical Model"
II. Atomic Orbitals and Sublevels
a. Areas of high probability of finding electrons
b. Four orbital shapes
1) s — sphere
2) p — dumbell
3) d — cloverleaf
4) f — nondescript
III. Electron configuration, Orbital diagrams, and electron dot structure
a. Electron configurations are used to organize electrons from lowest to highest energy
1) Whole numbers (1,2,3,4...) represent the main energy level
2) Letters (s,p,d,f) represent sublevels
3) Superscripts represent the number of electrons in that sublevel
` C has 6 electrons: 1s
22s22p2sublevel # orbitals maximum # e-
s 1 2
p 3 6
d 5 10
f 7 14
1. Aufbau principle — electrons enter orbital of lowest energy first
2. Pauli Exclusion principle — only two electrons with opposite spin are allowed in each orbital
3. Hund’s rule — each orbital within a sublevel must get an electron before doubling up
b. Orbital diagrams — same as electron configurations except show orbitals
C has 6 electrons: [X] [X] [/][/][ ]
1s 2s 2p
1. Valence electrons
— those electrons in highest energy levelC has 6 electrons: 1s
22s22p2 but only 4 electron are in the valence shell!2. Heisenberg uncertainty principle — can never know both velocity and position of electrons at the same time; so it is uncertain exactly where the electrons are at given moment in time
*Exceptions: Cu and Cr prefer half-filled and filled sublevels instead of partially filled sublevels; (Cu ends in 3d
IV. Spectra
a. Light behaves like waves as electromagnetic radiation
b. Elements emits light when electric current or heat is passed through giving an atomic emission spectrum unique to that element; electrons are energized to an excited state; when they drop back to the ground state, they emit this light
c. Speed of light = wavelength x frequency
c = l x v
E.M. Spectrum
Radio Micro IR VIS X-ray
/ \
ROYGBIV
800nm 400nm
GT:
1. Quantum concept — energy changes occur as discrete units called quanta; whereas classical physics predicts that energy changes are continuous; classical physics does not explain emission spectra of elements that are single lines; Max Planck explained that when a piece of steel is heated, it glows black, then red, yellow, white, and blue as temperature increases; this is due to energy changes of discrete units
2. Energy = Planck’s constant x frequency
E = h x n
where h = 6.6 x 10-34 J sE= h x n and n =c/x \ E = hc
l
3. Photoelectric effect — light quanta are called photons; electrons are ejected by metals when light shine on them; can only be explained by quantum theory not classical physics; there is a threshold frequency which must be attained in order for electrons to be ejected (not continuous); this means that not just any light will do; the frequency of the light must be at or above the threshold frequency to contain enough energy (E=h x n ) to eject the electrons; this converts light energy into electrical energy
4. Wave motion and quantum mechanics
— Louis de Broglie (1924) theorized that since light behaves as particles (and wave), perhaps particles can behave as waves;l= h
mv where m= mass of particle and v =velocity
we can calculate the velocity of a moving electron because m=9.11 x 10
-28g and l =2 x 10-8cm (about the diameter of an atom); according to de Broglie’s equation, all matter exhibits wavelike motion, but we don’t observe it in everyday life because the objects we encounter are too largeQuantum #’s