Honors Chemistry
Test Date: .
Chapter Six:

Electromagnetic Spectrum & Radiant Energy o know the order from gamma  radio
 be able to rank types of radiation (and colors) in order of increasing (or decreasing) energy, frequency, and wavelength
 know the relationships (and units of) between energy, frequency, and wavelength (directly or indirectly proportional?)

Energy Calculations (the constants & equations listed below will be given on the test) o Constants:
 c = 3.0 x 10 8 m/s = speed of light
 h = 6.63 x 10 -34 J•s = Planck’s constant
 R
H
= 2.18 x10 -18 J = Rydberg constant o Equations: (in all these equations,

must be in meters !)
 c =

= (wavelength)(frequency)
 E = h
 = (Planck’s constant)(frequency)

together, E = hc

 Rydberg  ΔE = R
H
1 - 1 .
n f
2 n i
2

Dual Nature of Light: behaves as both a particle & a wave o Photoelectric Effect
 one of the first pieces of evidence that supported the idea of light behaving as a particle (because in order to emit electrons from a metal atom, a minimum frequency must be reached – it has nothing to do with the intensity of light
(which correlates to amplitude of a wave) therefore, light must, in some part, be explained by the particle theory

Continuous vs. Line Emission Spectra o continuous spectra
 when which light is broken up by a prism, the entire rainbow is seen (all wavelengths present) o line emission spectra
 when colored light is put through a prism, only the distinct wavelengths which make up the light are seen (ex: 1 line of red at 725 nm, 1 yellow at 556 nm)
 line emission spectra are like a “fingerprint” for an elements because each element emits a different color when the different numbers of electrons in various energy levels throughout the atom are excited (flame test lab) which appear as all the separate wavelengths that make up that color.


Bohr Model o Electrons orbiting the nucleus of an atom in distinct paths – “planetary model”. Each path has a distinct amount of energy associated with it (quantized). o This model was truly just a stepping stone for the quantum mechanical model and electrons do not truly orbit the nucleus in these paths
 We took away two important things from his model: (1) electrons can only exist in discrete energy levels (2) energy is involved in moving electrons between the energy levels o electron “jumping”
 in order for the electron to move between energy levels, a specific quantity of energy must be gained or lost (like rungs on a ladder, only that exact amount of energy can cause the electron to “jump” levels)
 the lowest energy level, n = 1, is the ground state; any energy levels n = 2 or above are called excited states (the higher the n value, the more energy that level has and the more excited the electrons in that level are)
 this is an example of using the Rydberg equation to determine the energy absorbed or emitted when an electron moves between energy levels
 E emitted (-) as a photon when e
–
moves from higher to lower n
 E absorbed (+) by a photon of light causes e
–
to be promoted from lower to higher n = exciting the e
–

Heisenberg Uncertainty Principle o Says that you cannot know, with any great degree of certainty, both the location of an electron as well as its velocity at any given instant. The more certain you are of one factor, the less certain you are about the other factor.
 the smaller the mass, the greater the degree of uncertainty (thus, for electrons we cannot really be certain of its true location and speed)
 Bohr Model violates this principle

Quantum Mechanical Model of the Atom o This model combines the idea of electrons as both particles and waves, includes the distinct energy that was postulated by Bohr, and agrees with the Heisenberg
Uncertainty principle o Schrodinger
 developed wave equation which allows us to predict the likelihood (probability) of finding an electron in a general region of space around the nucleus o Principle Quantum Numbers
 n – principle quantum number = the main energy level … “the theatre”
 l – azmuthial quantum number (0  n -1)= the shape of a subshell (s,p,d,f)
(sphere, dumbbell, etc) … “the section”
 m l
– magnetic quantum number (l  + l )= the different # of orientations in space (tells you that an s has 1, p has 3, d has 5, and f has 7) … “the row”
 m s
– magnetic spin quantum number (+1/2, -1/2)= how is the electron moving in an orbital … “the seat”


Electron Configurations & Orbital Diagrams o Pauli Exclusion Principle, Aufbau Principle, & Hund’s Rule – know the definitions of each (use the memory tricks) and be able to distinguish which principle/rule is violated if there is an error in an electron configuration or orbital diagram o How is the periodic table set up? Where is 1s, 2s, 2p, etc? You need to know/memorize this order to correctly write the orbitals in order of increasing energy
(Aufbau principle). o The 6 exception elements and what the correct configurations for these elements would be o Orbital Diagrams
 drawing in the arrows to represent electrons
 must apply Pauli exclusion principle (one up, one down) and Hund’s Rule (for degenerate - know this word - p, d, and f orbitals) o Electron Configurations
 Standard – begins at 1s and work up
 Condensed/Shorthand – uses the last noble gas before the element to shorten the configuration
*** Problems in the text book are great for extra practice solving calculations!!! ***