Chapter 4

Arrangement of Electrons in Atoms

Objectives

  1. The Development of the New Atomic Model
    1. Rutherford’s model needed to be improved
      1. Had no place for the negative charges
      2. What prevented the negative electrons form being drawn into the positive nucleus?
    2. New discoveries about the atom where made using the absorption and emission of light.
  2. Properties of Light
    1. Light behaves as both a wave and a particle
      1. Wave Description of light
        1. Electromagnetic radiation- a form of energy that exhibits wavelike behavior as it travels through space. LIGHT
        2. Electromagnetic spectrum- the combination of all the forms of electromagnetic radiation.
          1. X-rays
          2. Ultraviolet and infrared
          3. Microwaves
          4. Radio waves
        3. All forms of electromagnetic radiation move at constant speed of 3.0 X 108 m/s
        4. The important characteristic of waves is the wavelength
          1. The distance between corresponding points on adjacent waves.
          2. Units are meter, centimeter, nanometer
        5. Frequency- as the number of waves that pass a given point in a specific time, usually one second.
        6. Equation- c=l v
    2. The photoelectric Effect
      1. Definition- refers to the emission of electrons from a metal when light shines on the metal.
      2. Solar panels make use of the photoelectric effect.
      3. Believed to be due to the particle effect of light.
      4. Quantum- the minimum quantity of energy that can be lost or gained by an atom.
      5. Deals with planks constant.
      6. Photon- is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.
    3. The hydrogen-Atom Line-Emission Spectrum
      1. Ground State- The lowest energy state of an atom
      2. Excited State-a state in which an atom has a higher potential energy than it has in its ground state.
      3. When electron falls back light of a given wavelength is emitted and color is produced.
    4. Bohr model of the Hydrogen Atom
      1. Due to the discovery of the emission of a photon at given quantum’s of energy.
      2. Bohr discovers the idea of energy levels.

Objectives

  1. The Quantum Model of the Atom (Section 4-2)
    1. Electrons as Waves-
      1. Investigations into the photoelectric effect and hydrogen atomic emission revealed that light could behave as both a wave and a particle.
        1. De Broglie- related Bohr’s model to the behavior of waves.
        2. Later proved that electrons could be diffracted and interfere with each other.
      2. The Heisenberg Uncertainty Principle-States that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.
      3. Schrodinger Wave Equation
        1. Took the idea of dual particle
        2. Backed up Bohr’s theories
    2. Heisenberg and Schrodinger together started the quantum model.
      1. Quantum Theory- describes mathematically the wave properties of electrons and other very small particles.
      2. Orbitals- a three-dimensional region around the nucleus that indicated the probable location of an electron.
  2. Atomic Orbitals and Quantum Numbers
    1. Quantum numbers-Specify the properties of atomic orbitals and the properties of electrons in orbitals.
      1. Principal Quantum Number-
        1. Symbol- n
        2. Indicates the main energy level occupied by the electron
        3. As n increases the electron’s energy and distance from the nucleus increases.
        4. Total number of orbitals = n2
      2. Angular Momentum Quantum Number
        1. Symbol- l
        2. Definition-indicates the shape of the orbital.
        3. Sublevels- orbitals of different shapes.
        4. Number of sublevels = n
        5. Table 4-1

        Orbital Letter Designations According to Values of l

        l

        Letter
           
           
           
           

      3. Magnetic Quantum Number-
          1. Symbol- m
          2. Indicates the orientation of an orbital around the nucleus.
      4. Spin Quantum Number-
          1. Two possible values (+1/2, -1/2)
          2. Indicates the spin of the electron.

Objectives
  1. Electron Configurations (Section 4-3)
    1. Electron configuration- The arrangement of electrons in an atom.
    2. Rules for Electron configurations
      1. Aufbau principle- an electron occupies the lowest-energy orbital that can receive it
      2. Pauli exclusion principle-no two electrons in the same atom can have the same set of four quantum numbers.
      3. Hund’s rule-orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in single occupied orbitals must have the same spin.
    3. Representing Electron Configurations
      1. Orbital Notation
        1. Orbitals represented by a line
        2. The name of the orbital written under the line
        3. Arrows represent the electrons
      2. Electron-Configuration Notation
        1. Eliminates the lines and the arrows
        2. Number of electrons is shown by a superscript
      3. Noble-Gas notation
        1. Use of the nearest noble gas to shorten the notation
      4. Stability
        1. Eight electrons in outer shell
        2. A half full shell next most stable.
    4. As always there are exceptions.