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Chemical principles: Atomic orbital

Introduction

 

Negatively charged electrons in an atom orbit around the positively charged nucleus to which they are attracted. Because of this attraction it takes energy to move the electrons away from the nucleus, and the closer they are the more energy is required. This attraction is affected by repulsion from other electrons. Only a fixed number of electrons can exist at a given energy level before this repulsion becomes so strong that further electrons have to exist at a higher energy level. The orbit of the electrons around the nucleus is also influenced by interaction with other atoms, but this article will focus on atomic orbital in individual atoms.

 

Electron Clouds

 

In schools, electron shells are usually represented as circles drawn around the nucleus like planets around the sun. These diagrams are simple to interpret but misleading as electron orbitals are not always spherical nor do the electrons have a fixed position in them.

Electrons are continually moving and at degree level you will most often see them represented in atomic orbitals as electron clouds. These clouds do not pinpoint the specific location of an electron but rather represent the probability of finding an electron within that volume of space around the nucleus. Where the cloud is most dense (darkest) the probability of finding an electron is highest.

 



Energy Levels

 

Energy levels, also called energy shells, are the name for the fixed energy levels electrons can exist at. Electrons fill the shells starting at the lowest energy levels closest to the nucleus and then filling the outer shells. When the electrons are in their lowest energy configuration this is called “ground state.” Introducing energy to the system can temporarily excite an electron into a higher shell. They must always absorb fixed amounts of energy as electrons cannot exist between energy potentials. However these energy levels are not evenly spaced out. Less energy is needed to excite an electron from a high energy level to the next one up than a low energy level to the next shell up.

 

Energy is released when the electron drops back down a level although often not in the same form. For example if an atom absorbs wavelength of light that excites it two energy levels higher in one go it may drop back one energy level at a time and each time release light at a different wavelength to the one it absorbed (known as fluorescence).

 



Valence electrons

The chemical behaviour of an atom depends mostly on the valence electrons. These are the electrons on the outermost electron shell.  Atoms with the same number of electrons in the valence shell tend to display chemically similar properties. For example fluorine and chlorine both have seven valence electrons and react in a very similar way when forming compounds.

 

The valence shell is the name given to the outer shell of electrons which can contain anything from just one electron to being completely full. The less free spaces for electrons there are in the valence shell the less reactive an atom tends to be. For example helium is described as an “inert” element because its valence shell is completely full making it very unreactive with other atoms.  Atoms with gaps in their valence shell fill them by reacting with other atoms which can involve electron sharing or loss and gain of electrons as shown in the diagram below. In both cases the free places in the valence shell become occupied.



Orbitals

 

An orbital is defined as the space in which an electron can be located 90% of the time. Each orbital can hold a maximum of two electrons and each energy shell has a specific number of orbitals. The sublevels, or subshells, of an energy level  contain orbitals of a specific shape. The orbitals can be remembered by the mnemonic Silly Penguins Don’t Fly for s (sharp), p (principle), d (diffuse) and f (fine) orbitals. Each of these orbitals has a specific shape, shown below. There is one spherical S orbital per shell and three dumbbell shaped p orbitals per shell. Four of the five d orbitals per shell have the same daisy shape but different orientations in space. The shape of f orbitals is more complex but there are seven f orbitals (up to fourteen electrons per shell).

 



 

The lowest energy level will always fill with electrons first. For example the first shell has one s orbital (called 1s) so a maximum of two electrons. This shell will fill up first. The second shell also has an s orbital (called 2s) and three p orbitals (called 2p). The p orbitals are higher energy so the shells will fill up in the order 1s (2 electrons), then 2s (2 electrons) and then 2p (6 electrons). This means two total electrons in the first shell, and eight in the second shell meaning a total of ten electrons in the atom. Where this gets a bit confusing is with the third shell when you get d orbitals. The d orbital is higher energy than the s orbital above it, for example the 4s shell fills before the 3d shell.

 

When there are equal energy places for the available electrons to fill, they will spread themselves out as far as possible. For example if there are three electrons to occupy the empty 2p orbitals one will go into 2px, one into 2py and one into 2pz instead of two electrons occupying the same orbital.

 

Writing out the electron configuration of an atom

 

Because they always fill up lowest energy level first it is possible to write out the electron configuration just from the atomic number. The atomic number can be found in the periodic table and is the number of protons in an atom, which is therefore equal to the number of electrons in an uncharged atom. To work out how many electrons in a charged atom simply add or subtract the charge. For example the atomic number of carbon is six. If it were C- you would add one electron for a total of seven and if it were C2+ you would subtract two electrons for a total of four. The diagram to the right shows the nomenclature for writing out the shells. Next we will show two examples:

 

Carbon (6) 1s2 2s2 2p2

Oxygen (8) 1s2 2s2 2p4



Remember that the 4s orbital has a lower energy than 3d orbitals  so the 4s orbital fills first:

 

Iron (26) 1s2 2s2 2p6 3s2 3p6 4s2 3d6 or 1s2 2s2 2p6 3s2 3p6 3d6 4s2

 

You may see them written both ways. The important thing is that the s orbital fills first but to be safe, when revising for exams check which way your lecturer does it and practise doing the same!

 

Electron configurations and the periodic table

 

Elements are grouped into the periodic table based on their number of valence electrons (columns) their outer energy level (rows) and their outer orbital type (blocks). The s block, p block, d block and f blocks are shown below.

 



References

 

 

  • Stryer, L., Berg, J. and Tymoczko, J. (2002) Biochemistry Revised 5th Edition. W.H.Freeman and Co Publishing.
  • Clayden, J., Greeves, N., Warren, S. and Wothers, P. (2001) Organic Chemistry. OUP Oxford Publishing.
  • eDewcate (2009) Shapes of atomic orbitals video tutorial. http://www.youtube.com/watch?v=F-xLQ1WBIl. Accessed 12/06/2012
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