The Structure of the Atom
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The Structure of the Atom
The protons and neutrons in each atom are tightly packed in a positively charged nucleus. Negatively charged electrons move around the nucleus. The number of protons in a nucleus defines the type of atom and is the same asthe atomic number. The number of neutrons is found by subtracting the atomic number from the mass number. In an atom because there is no overall charge the number of electrons equals the number of protons.
In chemical reactions the nucleus remains unchanged.
Geiger and Marsden bombarded a thin gold foil with a beam of alpha particles.
Most of the particles passed through the foil without deflection and were detected by a flash of light when the alpha particle struck a zinc sulphide screen, surrounding the gold foil.
A few were deflected and some of these were deflected at angles greater than 900, suggesting they had been repelled by large positive charges within the foil - nuclei of atoms of gold.
From GCSE you should be familiar with the Bohr model of electrons arranged around a nucleus. The electrons are in certain energy levels and each energy level can hold only up to a maximum number of electrons.
This is summarised in the table below:
|Energy level or 'shell'||Max no of electrons|
A sodium atom containing 11 electrons has an electron arrangement of 2,8,1. Two electrons filling the first shell, eight electrons filling the second shell and one electron in the outer third shell.
However, these models of electron arrangement are simple and a more advance done can now be used. It is possible to break these energy levels into sub-shells.
Electrons are impossible to locate exactly at any one time. It is however, possible to indicate a region or volume where the electron is most likely to be found. This region is called an Orbital.
Each orbital is capable of holding a maximum of 2 electrons. Orbitals can be divided into s, p, d, and f types. Each type has its own characteristic shape.
The shape of s and p orbitals are shown below:
The first energy level holds a maximum of 2 electrons in one s type orbital called 1s. There are no p, d, or f orbitals available at this energy level.
The second energy level consists of one s type orbital and three p type orbitals: 2s, 2px, 2py, 2pz.
Note: there are 3 p orbitals of identical energy, one along the x axis, one along the y-axis and one along the z-axis.
These four orbitals can hold a total of 8 electrons (i.e. 2 electrons each). There are no 2d or 2f orbitals.
The third energy level consists of: one s type orbital, three p type orbitals and 5 d type orbitals. These nine orbitals can hold a maximum of 18 electrons altogether (i.e. two electrons each).
Note: there are seven f type orbitals holding a maximum of 14 electrons in total.
When filling the available orbitals with electrons two important principles should be followed:
1. Electrons fill the lowest energy orbitals first and the other orbitals in order of ascending energy. It is incorrect to assume that an energy level is always completely filled before electrons enter the next energy level. The order of filling orbitals as shown below is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p.
2. Where there are several orbitals of exactly the same energy e.g. three 2p orbitals, electrons will occupy different orbitals whenever possible.
For example: nitrogen is 1s2 2s2 2px1 2py1 2pz1 and not 1s2 2s2 2px2 2py1.
This principle is Hund's rule. When an orbital only contains 1 electron then this electron is said to be unpaired.
a) The small number above the orbital refers to the number of electrons in the orbital: 1s2 means 2 electrons in a 1s orbital.
b) The electron arrangements are sometimes abbreviated.
For example: the electron arrangement for calcium may be written as 1s2 2s2 2p6 3s2 3p6 4s2.
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