Understanding Atomic Structure
Atomic structure is actually remarkably simple to understand; despite the fact a definite atomic structure hadn’t being cemented until the early 19th century. There are three main components to the atom: electrons, neutrons and protons. All three are summarised below.
>>Protons<<
Protons are positively charged sub-atomic particles which have a relative atomic weight of 1 and a charge of +1. They exist in the centre of the atom in the “nucleus”. The entire nucleus is held together by nuclear forces (sometimes called “residual strong force”) and contains within it 99.9% of the atom’s mass concentrated into one tiny spot.
>>Electrons<<
Electrons have a miniscule relative atomic mass (approximately 1/1836th of a neutron) however they do have a negative charge of -1. This means that they effectively “cancel out” the positive charges of the protons when examining the atom as a whole, however contribute nothing to the overall mass. Electrons are organised into shells/orbits around the the central nucleus in what is commonly compared to planets orbiting a star. The positions of these electrons are maintained by electromagnetic forces created due to the differentce between the charges of protons and electrons. Whilst the “reality” of electrons is that they exist in a quantum “cloud” around the nucleus the more conventional way of thinking allows us to make extremely accurate predictions about how an atom will react with other elements and the characteristics the element will have.
Whilst electronic configuration is a complex enough subject to be addressed in an entirely different article (or indeed a series of articles) it can simplified into some commonly used notations that are effective for common chemical applications. The electron configuration for Argon is shown below, the latter of which takes into account sub-orbitals.
– Electron per Shell: 2, 8, 8
– Electron Configuration: 1s2, 2s2, 2p6, 3s2, 3p6
>Key for Electron Configuration:
+ 1st Number- the shell number
+ 1st Letter- the type/location of the sub-orbital into which shells are divided named “blocks”.
+ 2nd Number- the number of electrons found in this sub-orbital.
[It should be noted that Argon is one of the last elements before the transition metals in which electronic configuration becomes more complicated by the inclusion of “d-blocks” and thus has been left out of this article.]
In very general terms atoms with 1 or 2 electrons in their outer shell will be very reactive because they are trying to lose their outer electrons and become more stable. Inversely atoms with 7 outer electrons will also be very reactive as they are trying very hard to gain extra electrons from other atoms. This happens because atoms are constantly trying to achieve a full outer shell of electrons as this is the most stable configuration. Atoms which already have full outer shells are the “Noble Gases” which are the most unreactive group of all elements, including Neon, Helium and Argon.
The two of the most important types of bonding that arise from electronic interactions are covalent and ionic. Covalent bonds are formed by two atoms “sharing” electrons one another. Ionic bonds meanwhile are caused by one atom donating it’s electrons to another atom and the imbalance in electrical charges which is created results in the two atoms attracting and sticking together.
>>Neutrons<<
The neutrons are another key part of the atomic nucleus. They have a relative atomic mass of 1 (the same as a proton); however they do not have any electric charge. Despite their non-existent charge they play a key role in atomic structure as they give the ability for atoms to form isotopes, which are outlined below.
>>Isotopes<<
Because of the neutral stance neutrons hold within atomic composition the number of them at are present can vary within from atom to atom within the same element. For example a carbon atom can have 6, 7, or even 8 neutrons without the atom becoming overly unstable and changing into something else through the processes of decay. When the mass of protons are counted too this makes 3 different plausible isotopes of carbon; carbon-12, carbon-13 and carbon-14. This can be applied to almost all of the elements throughout the periodic table. The reactivity or physical properties are not affected, only the mass of the atom.
>>Understanding Atomic Structure from Notation<<
There is a common way in which elements are written and atomic structure can be interpretted from this easily. If we continue to use Carbon as our example it is written in the periodic table as a capital C with a twelve above it and a 6 below it. For Carbon the “6” is the Atomic Number of the element. This represents the number protons which are in the atom, and because protons attract an equal number of electrons we can now determine that there are also 6 electrons.
The “12” of the Carbon is the Mass Number, which tells us how heavy an atom is in Relative Atomic Mass. Because we have already determined we have 6 protons we can take 6 away from 12 and this gives us the number of neutrons, which is also 6.
Further Reading:
http://www.chemguide.co.uk/atoms/properties/gcse.html
http://web.jjay.cuny.edu/~acarpi/NSC/3-atoms.htm
