Unit 3: Electronic Structure and Periodic Trends
Light
Spectrum & Range
longer wavelength = lower frequency & energy
equations
speed of light = wavelength (m) times frequency
c=λν
energy = Planck's constant (J-s) times frequency
h (Plank's constant) = 6.62.x10^ -34 Jxs
Planck
discovered radiation profiles emitted by solid bodies heated to incandescence could not be explained with existing knowledge
energy can only be lost in whole number multiples of quantity
Energy of a photon: E = hν = (hc)/λ
Planck's Constant
6.626x10^-34 Jxs
diffraction patterns can only be explained in terms of waves
continuous spectrum contains all the wavelengths of light
shorter wavelength = higher frequency & energy
Electrons & Light
light is emitted from an element when electrons move from a higher energy level to a lower one - releases energy
as an electron gets closer to the nucleus, the the attraction goes stronger
electron's natural state is it's ground state
excited states are ones that electrons move up to
when electrons fall from excite to ground states they give off energy as light
further distance > more energy > shorter wavelength
longer distance > less energy > longer wavelegth
Electrons
E-Configs
Bohr's model was flawed > math only worked for hydrogen
ideology of quantized energy levels was correct
n = number of electron energy levels
number of electrons that fit into a shell = 2n^2
different electron orbital shapes
s
p
d
f
one orientation
3 orientations
5 orientations
7 orientations
orbitals:
click to edit
each orbital only has 2 electrons
one spins up, the other spins down
energy levels > sub shells > orbitals > electrons
orbitals within the same shell & have the same energy are degenerate orbitals
Orbital Filling
ex. : all 2p orbitals have the same energy
1) Aufbau Principle > an electron occupies the lowest energy level it can
2) Pauli Exclusion Principle: no 2 electrons in the same atom can have the same spin. one spins up, the other spins down
3) Hund's Rule: orbitals of equal energy are each occupied by one electron before any electron is occupied by a second electron
Ions
Periodic Trends
Ionic Bonds
atomic radius decreases as you go across a period (left to right)
atomic radius increases down a group
Columbic attraction increases as atomic radius decreases
ionization energy tends to increase across a period and tends to decrease down a group
oyer electrons farther from nucleus, lower Columbia attraction
electronegativity decreases down a group and increases across a period
cation > positive ion
anion > negative ion
metals and nonmetals bond ionically
mainly formed by metals
mainly famed by nonmetals, halogens always form anions
ionic compound forms crystals by assembling nanny pairs
ionic formulas tell us the ratio of electrons not the number in the compound total
tightly covalently bonded compounds (polyatomic ions) can act as one unit to ionically bond with other elements