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.
Isotopes
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.
Isolectronic Structures
Isoelectronic species are elements or ions that have the same, or equal number of electrons.
Examples of Isoelectronic Species
Sc+3, Ca+2, K+1, Ar, Cl-1, S-2, P-3
All of the ions or elements above have got 18 electrons and therefore these would have the same electronic configuration.
Orbital Shapes
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.
Atomic Spectra
When atoms are excited they emit light of certain wavelengths which correspond to different colors. When looking at these wavelengths through a spectrometer these are seen as a series of colored lines with dark spaces in between; this series of colored lines is called a line or atomic spectra.
Each element produces a unique set of spectral lines. Since no two elements emit the same spectral lines, elements can be identified by their line spectrum.
Electronic Configuration
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.
An s-block element is a substance that has an electron in an s orbital.
An p-block element is a substance that has an electron in an p orbital.
An d-block element is a substance that has an electron in an d orbital.
Sodium is an s-block element while Carbon is a p-block element and iron is a d-block element. Can you show that this statement is correct?
The electronic configurations along period 2 are as follow:
B (Z=5) configuration: 1s2 2s2 2p1
C (Z=6) configuration:1s2 2s2 2p2
N (Z=7) configuration:1s2 2s2 2p3
O (Z=8) configuration:1s2 2s2 2p4
F (Z=9) configuration:1s2 2s2 2p5
Ne (Z=10) configuration:1s2 2s2 2p6
For two more complicated examples please find the electronic configurations of Yttrium and Bismuth:
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
Exceptions when writing the electronic configuration
When filling the d-orbitals both half full and full orbitals gain a little bit of stability and therefore these are favoured over orbitals where these are neither half full nor full. In the case of Chromium and Copper there will be an excitation of the electron from the to 4s the 3d. The electronic configurations of both Chromium and Copper are:
Chromium: 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Copper: 1s2 2s2 2p6 3s2 3p6 4s1 3d10
The period number is indicated by the position of the outermost electron in the atom. This would make chromium and copper as being two atoms in period 4.
The group number is indicated by the number of electrons in the outer orbital. This makes sodium and potassium as two atoms in group 1.
Ionic Radii
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 above. Anions become bigger while cations become smaller.
Changes across a period
Atomic radius decreases
Going across a period the number of protons increases while the shielding 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 from Magnesium to Aluminium there is a full s orbital, resulting in a little bit of extra shielding and thus less pull from the nucleus.
When it comes to Phosphorus and Sulfur there is a change from a half-full p orbital to paired electrons in the p orbital. These paired electrons repel, making it easier 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.
Successive Ionisation Energies
Ionisation energy is highly dependent on the size of the atom/ion and as the ion decreases in radius, the ionisation increases.
- The first electron in sodium is lost from the 3s1. This electron is found in the shell furthest away from the nucleus and therefore it would require the least amount of energy to be removed.
- The second electron in sodium is lost from the 2p6. This electron is found in the second shell and therefore it is much closer to the nucleus. This would need a higher amount of energy to be removed.
- The 3rd till the 9th electrons are then all lost from the 2nd shell. This would mean that the shielding remain similar, same as the nuclear pull. Since the number of electrons is decreasing the effective nuclear pull on each electron will be higher, making the radius smaller and resulting in a slight increase in the ionisation energy.
- The 10th electron will be removed from the 1s2 orbital, which is found in the first shell and thus closest to the nucleus. This would then result in another big increase in the ionisation energy.
The most imporant thing to note is that when an electron is removed from a shell that is closer to the nucleus there will be a big increase in the ionisation energy.