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The three particles in an atom are protons, neutrons and electrons. The details of these particles are,
The nucleus of an atom contains protons and neutrons, and is therefore positively charged.
The electrons in an atom form a cloud of electron density around the nucleus.
The numbers of protons, electrons and neutrons in a particular isotope of a particular atom are given in a periodic table by the mass number and atomic number. For example,
atomic number = number of protons
mass number = number of protons + number of neutrons
So if both of these numbers are given in an exam question, the numbers of protons, electrons and neutrons can be calculated.
|number of protons = 9|
|number of neutrons = 19 - 9 = 10|
|number of electrons = 9|
N.B.: In a neutral atom the number of protons = number of electrons. For a (charged) ion the number of electrons differs from the number of protons by the number of the charge, i.e. number of electrons = atomic number - charge.
|number of protons = 11|
|number of neutrons = 23 - 11 = 12|
|number of electrons = 11 - 1 = 10|
|number of protons = 17|
|number of neutrons = 37 - 17 = 20|
|number of electrons = 17 - (-1) = 18|
Isotopes of an element differ only in the number of neutrons in the nucleus
|number of protons = 92||number of protons = 92|
|number of neutrons = 235 - 92 = 143||number of neutrons = 238 - 92 = 146|
|number of electrons = 92||number of electrons = 92|
An atom is made up of a nucleus, containing protons and neutrons, and shells of electrons moving around the nucleus.
The rows, called periods, of a periodic table correspond to the gradually increasing energy shells, or energy levels.
From the second period onwards the energy level is split up into sub-levels. These splits are called atomic orbitals.
The first period is not split and has just one atomic orbital - the s orbital.
The second period is split into an s orbital and three p orbitals.
The third period is split into an s, three p and five d orbitals.
N.B.: The energy of the 3d orbital is greater than that of the 4s orbital.
Each separate atomic orbital may have a maximum of two electrons in it. Therefore the first period may have 2 electrons in it; the second period may have 8 electrons and the third period may have 18 electrons and so on.
The s orbital has the shape of a sphere,
The p orbital shapes are more complex. They are shaped like dumbbells, with the lobes placed along the three different principle axes in space.
The ionisation energy for an atom is the energy required to remove an electron from that atom. The first ionisation energy is the energy required to remove one electron from gaseous atoms of the element,
Element(g) → Element+(g) + e-
A graph of first ionisation energies vs. atomic number shows some interesting patterns,
The general trend in the energy values is a decrease with increasing atomic number. This is because as the number of electrons increases so does the number of shells, which means the electrons gradually get further away from the nucleus. The pull on the electrons by the nucleus decreases therefore the ionisation energy decreases.
N.B.: As the atomic number increases the number of protons in the nucleus increase, which increases the attraction of the electrons to the nucleus; however, this increase in positive charge is overridden by the increased distance of the negative electrons from the nucleus.
Also as the number of shells increases the electrons in the inner shells shield the outer shell electrons from the attraction of the nucleus. This further lowers the ionisation energy of the outer shell electrons.
There are various blips in this general trend caused by the various atomic orbitals - s, p and d. As you fill an atomic orbital it becomes more stable and therefore harder to remove an electron from.
Taking a closer look at the second period i.e. atomic number 3-10 :
At beryllium (atomic number 4) the 2s orbital is full and therefore the energy needed to remove an electron is higher than that for lithium (atomic number 3). At boron (atomic number 5) the 2p orbital has only one electron in it and this means the energy dips down a little. As the 2p orbitals are gradually filled the ionisation energy increases because of the increased stability.
This can also be seen with a graph of log10(ionisation energy) for successive ionisation energies for a particular element e.g. for aluminium,
As electrons are lost from a complete orbital the energy jumps up a little disrupting an otherwise linear change in ionisation energy.back to top
The electron configuration for an element is a representation of the positions of the atom's electrons in the various atomic orbitals, s-, p-, d-, etc., see below,
|Li||1s22s1 or [He]2s1|
|Be||1s22s2 or [He]2s2|
|B||1s22s22p1 or [He]2s22p1|
|C||1s22s22p2 or [He]2s22p2|
|N||1s22s22p3 or [He]2s22p3|
|O||1s22s22p4 or [He]2s22p4|
|F||1s22s22p5 or [He]2s22p5|
|Ne||1s22s22p6 or [He]2s22p6|
|Na||1s22s22p63s1 or [Ne]3s1|
|Mg||1s22s22p63s2 or [Ne]3s2|
|Al||1s22s22p63s23p1 or [Ne]3s23p1|
|Si||1s22s22p63s23p2 or [Ne]3s23p2|
|P||1s22s22p63s23p3 or [Ne]3s23p3|
|S||1s22s22p63s23p4 or [Ne]3s23p4|
|Cl||1s22s22p63s23p5 or [Ne]3s23p5|
|Ar||1s22s22p63s23p6 or [Ne]3s23p6|
N.B.: As can be seen above it is sometimes acceptable to write a shorthand version of the electron configuration, by using "[noble gas]" to represent the electron configuration of the nearest previous noble gas.
The electron configurations of ions differ from their atom's configuration by the number of electrons that have either been added, for a negative ion, or removed, for a positive ion.
written by Dr Richard Clarkson : © Saturday, 1 November 1997
Updated : Thursday 30th August, 2012
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