Protons, Neutrons, Electrons
|Proton||Inside the nucleus||
|Neutron||Inside the nucleus||
|Electron||Outside the nucleus||
Charge of an electron = -1.6 x 10-19C
Mass of an electron = 9.11 x 10-31kg
Mass of a proton = 1.66 x 10-27kg
Much of the atom is empty space. The nucleus, which contains nearly all its mass has about 1/1000 of the radius of the atom.
Hydrogen is the only element that has special names for its forms.
Hydrogen Deuterium Tritium (Radioactive)
18 Neutrons (Number of neutrons = A - Z)
Relative Atomic Theory
In practice the Ar of an atom is the same of its mass number numerically.
For one atom
Ar = mass of the atom
1/12 of mass of Carbon-12 atom
For an element
e.g. 3/4 of chlorine atoms are Cl-35 and 1/4 is Cl-37. Find the Ar
Consider 100 atoms of Chlorine.
(75 x 35) + (25 x 37) = 3500
So average mass =
number of atoms
Ar (Cl) = 35.5
In a neutral atom, the number of electrons is equal to the number of protons.
Electrons occupy certain regions of an atom call Shells or Energy Levels.
The first energy level is full when it contains 2 electrons.
The second can hold 8 electrons.
The third an also hold 8 electrons.
The Periodic Table
This is the arrangement of elements in terms of their electronic structure.
From left to right, the proton number increases by 1.
Each horizontal row (a Period), energy levels build up with electrons.
Each vertical row (a Group), contains elements with similar properties.
When a reaction takes place, the shells furthest from the nucleus react first. This is why elements in the same group have similar properties.
Group 1 - Alkali Metals.
Group 2 - Alkaline Earth Metals.
Group 7 - Halogens.
Group 0 - Noble Gases.
History of the Periodic Table
There are 92 naturally occurring elements. Chemists had always wanted to find a way of fitting the elements into a pattern. In 1865, J. Newlons arranged the elements in order of increasing relative atomic mass.
He noticed that the eighth element resembled the first, the ninth, the second and so he arranged them in columns.
H Li Be B C N O
F Na Mg Al Si P S
Cl K Ca Cr Ti Mn Fe
The noble gases had not yet been discovered.
Newlon compared the elements with musical octaves, and called the resemblance, the law of octaves. The comparison with music was unfortunate and scientist laughed at him.
In 1869, a Russian Chemist, D.I Mendeleev also arranged the elements in order of increasing mass. He called it the periodic table which included improvements.
1 2 3 4 5 6 7 8
2 Li Be B C N O F Ni
3 Na Mg Al Si P S Cl Fe Co
4 K Ca Ti Vg Mn
5 Cu Zn As Se Br
Mendeleev introduced a long period for the elements which did not fit in. We now call these the Transitional Metals.
Mendeleev left spaces, for example, when he saw that Arsenic fitted in Group 5, he left two spaces after Zinc.
When two elements do not fit perfectly into slots in the Periodic Table suggested by their relative atomic mass, Mendeleev decided that the mass may be wrong, and more accurate methods were found.
New values for masses of Cr, In, Au, Pt justified Mendeleev's arrangement.
Mendeleev predicted that new elements found would fill the gaps. He had some outstanding success in predicting the properties they would have. This meant that faith in the periodic table soared.
The introduction of the Noble Gases
When the noble gases were discovered, they were found to fit between the halogens in Group 7 and the alkaline metals in Group 1. A separate group, 0 was added, however Argon has a higher relative atomic mass than Potassium ( (Ar)Ar = 40, (Ar)K = 39). It made more sense to put Potassium in Group 1 and to put Argon with the noble gases rather than stick to the relative atomic mass theory. Another kind of element Tellurium had the same problem with Iodine. Relative atomic masses placed tellurium under bromine and iodine under selenium.
Moseley's work with X-rays solved this problem. He discovered an new property of elements, proton number. This fixed the periodic table.
Trends in the Periodic Table
Elements in the same group have similar properties because they have the same number of electrons in the outer energy level.
Groups 1 and 2 are made up of metals, and metals appear in all the other groups apart from the noble gases. The metallic character increases down the group.
Elements in Groups 1 and 2 become more reactive down the group.
When they react, electrons are lost forming positive ions. This becomes easier down the group as the outer electrons are further away from the attraction of the nucleus and more importantly shielded from the nucleus by more shells of low energy electrons.
This means that Group 1 and 2 metals at the bottom of the group are better at gaining electrons thus making them more reactive than those at the top of the group.
The Group 7 metals are known as the Halogens. They are non-metals and are 1 short of an electron. When they react, chemically or covalently, they from a negative ion. When they are not in a compound or in a compound with a non-metal, each atom takes one covalent bond thus achieving a ""Noble Gas Electron Arrangement".
Elements in Group 7 become less reactive down the group.
When they react, electrons are gained forming negative ions. This becomes harder down the group as the incoming electrons are further away from the attraction of the nucleus and more importantly shielded from the nucleus by more shells of low energy electrons.
This means that Halogens at the top of the group are better at gaining electrons thus making them more reactive than those at the bottom of the group.
Metals at the bottom of the group are easier to cut and have lower melting points. The attractive forces between the Cations and the delocalised electrons are weak.