Electrons in Atoms Chapter 5

5.1 Wave Nature of Light  Electromagnetic Radiation is a form of energy that exhibits wavelike behavior as it travels through space.

Waves  Wavelength is the shortest distance between equivalent points on a continuous wave (λ)  Frequency is the number of waves that pass a given point per second, measured in Hz or s -1 (ν)  Amplitude is the wave’s height from origin to crest or trough

Waves  All electromagnetic waves travel 3.00x10 8 m/s  Wavelength and frequency can be calculated using the following equation: c = λν  The electromagnetic spectrum encompasses all forms of radiation; long wavelength low frequency on one end and short wavelength high frequency on the other

Particle Nature of Light  The wave model of light does not explain light’s interactions with matter.  Planck concluded that matter can gain or lose energy in specific amounts called quanta. He proposed that the energy of a quantum (minimum amount of energy that can be gained or lost by an atom) is related to the frequency of the emitted radiation: h = 6.626x Js = Planck’s constant

Example problem 

Einstein  Einstein proposed that electromagnetic radiation has both wavelike and particlelike natures.  Light has many wavelike characteristics AND can also be thought of as a stream of tiny particles called photons.  A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

5.2 Quantum Theory and the Atom  Bohr Model The lowest energy state of an atom is called its ground state. When an atom gains energy, it is in an excited state. Bohr related energy states to the motion of the electron within the atom.

Bohr’s Model and de Broglie’s Equation  Bohr’s model explained hydrogen well, but failed to explain the spectra of other elements. Electrons do NOT move around the nucleus in circular orbits!  de Broglie predicted that all moving particles have wave characteristics (including electrons) λ = h / νm

The Heisenberg Uncertainty Principle  Heisenberg studied interactions between photons and electrons and determined that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

Schrodinger  Schrodinger derived an equation that treated the hydrogen atom’s electron as a wave.  This model applied equally well to atoms of other elements!  This atomic model became known as the quantum mechanical model of the atom.  Schrodingers equation results in a solution known as a wave function, which is related to the probability of finding an electron within a particular region of space around the nucleus.

Orbitals  Electrons occupy three- dimensional regions of space called atomic orbitals.  These orbitals describe an electron’s probable location.  There are four types of orbitals, denoted by the letters s, p, d, and f.

5.3 Electron Configurations  The arrangement of electrons in an atom is called the atom’s electron configuration.  Electron configurations are described by three rules: The aufbau principle The Pauli exclusion principle Hund’s rule

Aufbau Principle  All orbitals related to an energy sublevel are of equal energy. (ex. All 2p orbitals are same energy)  The energy levels within a principal energy level have different energies. (ex. 2p higher than 2s)  The sequence of energy sublevels within a principal energy level is s, p, d, and f.  Orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal level. (ex. 4s is lower than 3d)

Aufbau Principle

Pauli Exclusion Principle  A maximum of two electrons may occupy a single atomic orbital, but only if the electrons have opposite spins.  Arrows are used to indicate electrons in an orbital.

Hund’s Rule  Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbital.

Valence Electrons  Electrons related to the atom’s highest principle energy level are referred to as valence electrons.  Valence electrons determine the chemical properties of an element.

Electron Configurations  Electron configurations may be represented using orbital diagrams, electron configuration notation, and electron- dot structures.