The atom is the smallest particle of a chemical element that can exist while retaining the same properties as the original element.
Atoms are made up of three main components, neutrons, protons and electrons.
- Protons are positively charged and are found in the nucleus.
- Neutrons are the particles that keep the atom together. without neutrons, the protons would repel each other, and these would act as a barrier to the protons. These do not have a charge.
- Electrons are negatively charged particles which are orbiting around the nucleus.
The mass of the protons and neutrons is more or less equivalent while the mass of the electron is very lower, and thus only the mass of the protons and neutrons when the mass of the atom is considered.
For more information on how to calculate the electrons and how to put electrons in shells click here.
The nucleus can be considered to be the heart of the atom, where the protons indicate the composition of the atom while the neutrons indicate the radioactivity of the nucleus.
This statement shows that a change in the proton number would change a nucleus from one element to another while a change in the neutron number will only change the stability of the nucleus. An isotope is when an element would have a nucleus with the same number of protons but a different number of neutrons. This explains the fact that no element has got a whole mass number due to the fact that this would be the average of all of the elemental nuclei.
For information on how to work the RAM from isotoped using Mass Spectrometry please click here.
Radioactivity has found its way in both archaeological and medical industries. In medicine, isotopes are used in several medical tests while in archaeology Carbon 14 is a very important tool in dating their findings.
For more information about Carbon14 click here.
The atom is made up of different orbitals, with each orbital being divided into sub-orbitals. These sub-orbitals are further divided into groups, which are s, p, d and f.
These are the shapes of the orbitals, with p-orbitals being in 3 subgroups and d orbitals being dividing into 5 subgroups, all being at the same energy but in different orientations in space.
When filling up orbitals with electrons there are a few rules that one must note. These are:
- The lowest energy level orbitals must be filled first.
- If there are any degenerate sub-orbitals (same energy) then these would be filled by putting one electron each with the same spin.
- Each sub-orbital must not have more than two electrons, which would have opposite spins.
This configuration would then divide the elements into groups or blocks, depending on the final orbital that is filled.
Sodium is an s-block element while Carbon is a p-block element and iron is an f-block element. Can you show that this statement is correct?
Yttrium: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d1
Bismuth: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p3
Periodicity is the study of the periodic table. The basic concepts re explained in this topic, with further reactions being discussed with their respective groups.
Changes across a period
Atomic radius decreases
Going across a period the number of protons increases while the number of shells remains the same. This means that each electron in the outer shell is being pulled towards the centre by more protons, making the radius smaller.
Electron affinity increases
As one progresses across a period it becomes easier for an atom to gain an electron, since the radius is smaller and more protons are available.
Ionisation energy increases
The energy required to lose an electron become bigger since the radius is smaller and the number of protons increases, making it more difficult to lose an electron.
As one can see from this graph there is a decrease in energy between Magnesium and Aluminium, and Phosphorus and Sulfur. This is due to the fact that Magnesium and Phosphorus are fully filled orbitals and half-filled orbitals. These increase stability, and therefore these would require more energy to lose an electron.
Changes down a group
Atomic radius increases
Doing down a group there is an increase in the number of shells, and therefore the radius would increases.
Electron affinity decreases
The bigger radius down a group would mean that the nuclear pull decreases down a group, making it more difficult to attract an electron.
Ionisation energy decreases
The bigger radius down a group would mean that the nuclear pull decreases down a group, making it easier to lose an electron from the outer shell, and thus decreasing the ionisation energy down a group.
When an atom gains an electron the radius increases.
When an atom loses an electron the radius decreases.
When an atom loses an electron the number of protons would be bigger than the number of electrons and therefore each electron would have a larger nuclear pull making the radius smaller.
When an atom gains an electron the number of protons would be smaller than that of the electrons, and therefore each electron would have a smaller nuclear pull making the radius bigger.
The change in size between the atom and the ions can be seen in the diagram below. Anions become bigger while cations become smaller.
A relationship within the periodic table by which certain elements in the second period have a close chemical similarity to their diagonal neighbours in the next group of the third period. This is particularly noticeable with the following pairs.
Lithium and magnesium:
(1) both form chlorides and bromides that hydrolyse slowly and are soluble in ethanol;
(2) both form colourless or slightly coloured crystalline nitrides by direct reaction with nitrogen at high temperatures;
(3) both burn in air to give the normal oxide only;
(4) both form carbonates that decompose on heating.
Beryllium and aluminium:
(1) both form highly refractory oxides with polymorphs;
(2) both form crystalline nitrides that are hydrolysed in water;
(3) addition of hydroxide ion to solutions of the salts gives an amphoteric hydroxide, which is soluble in excess hydroxide giving beryllate or aluminate ions [Be(OH)4]2− and [Al(OH)4]−;
(4) both form covalent halides and covalent alkyl compounds that display bridging structures;
(5) both metals dissolve in alkalis.
Boron and silicon:
(1) both display semiconductor properties;
(2) both form hydrides that are unstable in air and chlorides that hydrolyse in moist air;