2.2 Electron configuration
The electromagnetic spectrum
Light is made of discrete packets of energy called photons. Photons carry momentum, have no mass, and travel at the speed of light. All light has both particle-like and wave-like properties.
FREQUENCY The number of crests that pass a given point within one second is de- scribed as the frequency of the wave. One wave—or cycle—per second is called a Hertz (Hz), after Heinrich Hertz who established the existence of radio waves. A wave with two cycles that pass a point in one second has a frequency of 2 Hz.
All electromagnetic waves travel at the same speed(c) but can be distinguished by their different wavelengths(λ). Different colours of visible light have different wavelengths; red light, for example, has a longer wavelength than blue light.The wavelength and frequency are related by the equation: c=νλ
c (speed of light) = 3 x 108m/sec
Bohr model
Unlike planets orbiting the Sun, electrons cannot be at any arbitrary distance from the nucleus; they can exist only in certain specific locations called allowed orbits. This property, first explained by Danish physicist Niels Bohr in 1913.
The atom then remains in its lowest-energy state, which is also called the ground state. | The state of an atom after one or more electrons have been brought to a higher energy level is called an excited state. | When an electron “falls” from an outer shell to an inner shell, a photon is emitted. The energy of the photon corresponds to the energy difference of the two shells. |
Ephoton=∆Eelectron=hν
This equation and the value of h(the Planck constant) are given in sections 1 and 2 of the IB data booklet.
Because different orbits have different energies, whenever a quantum leap occurs, the energy possessed by the electron will be different after the jump. For example, if an electron jumps from a higher to a lower energy level, the lost energy will have to go somewhere and in fact will be emitted by the atom in a bundle of electromagnetic radiation. This bundle is known as a photon, and this emission of photons with a change of energy levels is the process by which atoms emit light.
Bohr’s model was a tremendous success in explaining the spectrum of the hydrogen atom. Unfortunately, when the mathematics of the model was applied to atoms with more than one electron, it was not able to correctly predict the frequencies of the spectral lines. While Bohr’s model represented a great advancement in the atomic model and the concept of electron transitions between energy levels is valid, improvements were needed in order to fully understand all atoms and their chemical behavior.
The hydrogen atom gives out energy when an electron falls from a higher to a lower energy level. Hydrogen produces visible light when the electron falls to the second energy level (n= 2). | The transitions to the first energy level (n= 1) correspond to a higher energy change and are in the ultraviolet region of the spectrum. Infrared radiation is produced when an electron falls to the third or higher energy levels |
(2.2 - The hydrogen emission spectrum – youtube!!!)
Atomic orbitals
Atomic orbitals are commonly designated by a combination of numerals and letters that represent specific properties of the electrons associated with the orbitals—for example, 1s, 2p, 3d, 4f. The numerals, called principal quantum numbers, indicate energy levels as well as relative distance from the nucleus. A 1selectron occupies the energy level nearest the nucleus. A 2s electron, less strongly bound, spends most of its time farther away from the nucleus. The letters, s, p, d, and f designate the shape of the orbital.
Now that we have defined the physical spaces that electrons can occupy, we need to determine the order of electron orbital filling. There are three major rules that we need to follow when filling electron orbitals.
3 RULES FOR CONFIGURATIONS
Aufbau Principle: Electrons are filled according to their lowest energy possible
Pauli exclusion principle: Electrons must differ in one of 4 quantum numbers (max 2 electrons in one orbital)
Hund’s Rules: Electrons want to have maximum SPIN
Energy Filling Diagram. Orbitals with the lowest energy are filled with electrons before orbitals at higher energy levels.
Thenth energy level of the Bohr atom is divided into nsub-levels. For example, the fourth level (n= 4) is made up from four sub-levels. The letters s, p, d, and f are used to identify different sub-levels. Each main level can hold a maximum of 2n2electrons. The 3rd energy level, for example, can hold a maximum of 18 electrons (2 × 32 = 18).
s sub-levels can hold a maximum of 2 electrons.
p sub-levels can hold a maximum of 6 electrons.
d sub-levels can hold a maximum of 10 electrons.
f sub-levels can hold a maximum of 14 electrons.
(Quantum Numbers, Atomic Orbitals, and Electron Configurations–youtube!!!)
