Structure 1.3 - Electronic configurations

• • The electromagnetic spectrum is a range of frequencies that covers all electromagnetic radiation and their respective wavelengths and energy
• • The spectrum shows the relationship between frequency, wavelength and energy
• • Frequency is how many waves pass per second, and wavelength is the distance between two consecutive peaks on the wave
• •  the higher the frequency, the shorter the wavelength
  • The equation that links them is c = νλ 
• • A continuous spectrum in the visible region contains all the colours of the spectrum
  • This is what you are seeing in a rainbow, which is formed by the refraction of white light through a prism or water droplets in rain
• a line spectrum only shows certain frequencies
• • Electrons move rapidly around the nucleus in energy shells
  • If their energy is increased, then they can jump to a higher energy level
  • The process is reversible, so electrons can return to their original energy levels
  • When this happens, they emit energy
  • The frequency of energy is exactly the same, it is just being emitted rather than absorbed:
• The difference between absorption and emission depends on whether electrons are jumping from lower to higher energy levels or the other way around
• • The energy they emit is a mixture of different frequencies
  • This is thought to correspond to the many possibilities of electron jumps between energy shells
  • If the emitted energy is in the visible region, it can be analysed by passing it through a diffraction grating
  • The result is a line emission spectrum
• • Line emission spectra Each line is a specific energy value
  • This suggests that electrons can only possess a limited choice of allowed energies
• • lines in the line emission spectrum get closer together towards the right end of the spectrum
  • This is called convergence and the set of lines is converging towards the higher energy end, so the electron is reaching a maximum amount of energy
  • This maximum corresponds to the ionisation energy of the electron
• • Niels Bohr applied the Quantum Theory to electrons in 1913, and proposed that electrons could only exist in fixed energy levels
  • The line emission spectrum of hydrogen provided evidence of these energy levels and it was deduced that the families of lines corresponded to electrons jumping from higher levels to lower levels
• • jumps to n3 go to the infrared region and the energy is low
• • jumps to n2 go to the visible region and the energy is increasing
• • jumps to n1 go the UV region and the energy is the highest there
• • The arrangement of electrons in an atom is called the electronic configuration
• • Electrons are arranged around the nucleus in principal energy levels or principal quantum shells
  • Principal quantum numbers (n) are used to number the energy levels or quantum shells
  • The lower the principal quantum number, the closer the shell is to the nucleus
  • The higher the principal quantum number, the greater the energy of the electron within that shell
• • Each principal quantum number has a fixed number of electrons it can hold
  • n = 1 : up to 2 electrons
  • n = 2 : up to 8 electrons
  • n = 3 : up to 18 electrons
  • n = 4 : up to 32 electrons
• • The principal quantum shells are split into subshells which are given the letters s, p and d
  • Elements with more than 57 electrons also have an f subshell
  • The energy of the electrons in the subshells increases in the order s < p < d
• principle quantum number and sub-shells

  
• • The ground state is the most stable electronic configuration of an atom which has the lowest amount of energy
  • This is achieved by filling the subshells of energy with the lowest energy first (1s)
  • This is called the Aufbau Principle
• The principal quantum shells increase in energy with increasing principal quantum number
• • The subshells increase in energy as follows: s ＜ p ＜ d ＜ f
  • The only exception to these rules is the 3d orbital which has slightly higher energy than the 4s orbital, so the 4s orbital is filled before the 3d orbital
• • Each shell can be divided further into subshells, labelled s, p, d and f
  • Each subshell can hold a specific number of orbitals:
  • s subshell : 1 orbital
  • p subshell : 3 orbitals labelled px, py and pz
  • d subshell : 5 orbitals
  • Each orbital can hold a maximum number of 2 electrons so the maximum number of electrons in each subshell is as follows:
  • s : 1 x 2 = total of 2 electrons
  • p : 3 x 2 = total of 6 electrons
  • d : 5 x 2 = total of 10 electrons
• • In the ground state, orbitals in the same subshell have the same energy and are said to be degenerate, so the energy of a px orbital is the same as a py orbital
• Pauli's Exclusion principle states that no two electrons in an atom can occupy the same orbital at the same time. This means that if one electron is present in an orbital then it must have opposite spin to any other electron already present in that orbital.
• • The s orbitals are spherical in shape
  • The size of the s orbitals increases with increasing shell number
  • E.g. the s orbital of the third quantum shell (n = 3) is bigger than the s orbital of the first quantum shell (n = 1)
• • The p orbitals are dumbbell-shaped
  • Every shell has three p orbitals except for the first one (n = 1)
  • The p orbitals occupy the x, y and z axes and point at right angles to each other, so are oriented perpendicular to one another
  • The lobes of the p orbitals become larger and longer with increasing shell number
• The electron configuration shows the number of electrons occupying a subshell in a specific shell
• • The transition metals fill the 4s subshell before the 3d subshell, but they also lose electrons from the 4s first rather than from the 3d subshell
• the periodic table is split up depending on their electronic configuration

  
• • An orbital can only hold two electrons and they must have opposite spin - the is known as the Pauli Exclusion Principle
• • This is because the energy required to jump to a higher empty orbital is greater than the inter-electron repulsion
  • For this reason, they pair up and occupy the lower energy levels first
• • Chromium and copper have the following electron configurations:
  • Cr is [Ar] 3d5 4s1 not [Ar] 3d4 4s2
  • Cu is [Ar] 3d10 4s1 not [Ar] 3d9 4s2
  • This is because the [Ar] 3d5 4s1 and [Ar] 3d10 4s1 configurations are energetically favourable
  • By promoting an electron from 4s to 3d, these atoms achieve a half full or full d-subshell, respectively
• —Emission spectra are produced by atoms emitting photons when electrons in excited states return to lower energy levels.
• —The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies
• 3—The main energy level is given an integer number, n, and can hold a maximum of 2n 2 electrons.
• —A more detailed model of the atom describes the division of the main energy level into s, p, d and f sublevels of successively higher energies.
• Each orbital has a defined energy state for a given electron configuration and chemical environment, and can hold two electrons of opposite spin. Sublevels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.
• In an emission spectrum, the limit of convergence at higher frequency corresponds to ionization.