Chapter 11: Electrons in atoms

1. The development of atomic models
2. The quantam mechanical model of the atom
3. Atomic orbitals
4. Electron configurations [Practice Orbital Filling Diagrams] [Practice Inert Core Notation] [Reference]
5. Exceptional electron configurations
6. Light and atomic spectra
7. The quantam concept
8. Light as particles: the Photoelectric Effect
9. Atomic spectra
10. Wave motion of matter & quantam mechanics

Chapter 11

11-1 The Development of Atomic Models

- Dalton's Atomic Theory: 1) All elements are composed of tiny indivisible particles called atoms; 2) Atoms of the same element are identical. The atoms of any one element are different from those of any other element; 3) Atoms of different elements can combine with one another in simple whole number ratios to form compounds; 4) Chemical reactions occur when atoms are seperated, joined, or rearranged. However, atoms of one element are not changed into atoms of another by a chemical reaction.

- Dalton mostly wrong.

- Thomson: electrons.

- Rutherford: nucleus.

- Bohr: orbits.

- Energy levels - like ladder; electrons can move from one energy  level to the next. A quantum is the amount of energy needed to move one electron to the next energy level.

11-2 The Quantam Mechanical Model of the Atom

- Quantam Model: based on probability.

- Erwin Schroedinger.

11-3 Atomic Orbitals

- Not actually orbitals per se.

- There are energy levels around the nucleus.

- "Within" each energy level there are up to 4 sublevels: S, P, D, and F.

- Within each sublevel there are a certain number of orbitals. S=1, P=3, D=5, F=7.

- S orbitals are spherical, P orbitals are dumbbell-shaped, D orbitals are four-leaf-clover shaped, the F orbitals are too strange to describe with words.

11-4 Electron Configurations [Practice Orbital Filling Diagrams] [Practice Inert Core Notation] [Reference]

- Aufbau Principle: electrons enter orbitals of lowest energy first. This is the order: 1s, 2s, 2p, 3s, 3p, 4s, 4d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. This order also appears in Chempire's Online Orbital Filling Order Guide.

- Pauli Exclusion Principle: there may be either 0, 1, or 2 electrons in any one orbital.

- Hund's Rule: orbitals in a sublevel each get one electron before any get two.

- Notation: you can either use boxes with arrows to represent electrons (each orbital is one box) or shorthand: 1s^2, 2s^1 means that the 1s orbital is full and that the 2s orbital has one electron in it.

11-5 Exceptional Electron Configurations

- Sometimes it is better to have 2 half-filled sublevels than 1 filled and 1 partially filled.

11-6 Light and Atomic Spectra

- Amplitude: center line to crest.

- Wavelength (l): the length of one complete wave (both the up and the down).

- Frequency (n ):the number of complete waves that pass in a unit of time.

- n = c / l  (c equals the speed of light: 3.0 x 10^10).

- All light is in a spectrum:

 Low Energy                                                                                                     High Energy

 Low Frequency                                                                                          High Frequency

 [ radio waves, microwaves, infrared, red, green, blue, ultraviolet, x-rays, gamma rays]

 High Wavelength                                                                                     Low Wavelength

- Atoms emit light when they lose energy. This light can be passed through a prism and the resulting bands of color can be used to identify the atom.

11-7 The Quantam Concept

- Planck: amount of energy absorbed or emitted by a body is proportional to the frequency of the radiation: E = h x n where h is Planck's constant [6.6262 x 10^(-34)].

11-8 Light as Particles: the Photoelectric Effect

- Einstein: light is quanta of energy that act like particles. These light quanta are called photons.

- Photoelectric Effect: electrons called photoelectrons are ejected by metals when light shines on them.

11-9 An Explanation of Atomic Spectra

- Black lines on absorbtion spectra occur where electrons absorb the energy to jump to the next energy level. Colored lines on emission spectra occur where electrons release the energy when they move down a level.

11-10 The Wave Motion of Matter and Quantam Mechanics

- de Broglie's equation: all matter exhibits wavelike properties.

- Heisenberg Uncertainty Principle: knowledge of the position and velocity of a particle are inversely proportional. I.e., you can't know exactly where a particle is and how fast it's going at the same time.

Go to Chapter: 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20