Valence electrons are outer shell electrons of an atom. They participate in the formation of a chemical bond, if the outer shell of the element is not full.
Valency electrons of the first 20 elements
Electrons of the first 20 elements are arranged around the nucleus in different energy levels as follows:
- first shell holds a maximum of 2 electrons
- second shell holds a maximum of 8 electrons
- third shell holds a maximum of 8 electrons
| Element | Number of electrons | Electronic Configuration | Number of valence electrons | Group number |
|---|---|---|---|---|
| Hydrogen | 1 | 1 | 1 | 1 |
| Helium | 2 | 2 | 2 | 8 |
| Lithium | 3 | 2.1 | 1 | 1 |
| Beryllium | 4 | 2.2 | 2 | 2 |
| Boron | 5 | 2.3 | 3 | 3 |
| Carbon | 6 | 2.4 | 4 | 4 |
| Nitrogen | 7 | 2.5 | 5 | 5 |
| Oxygen | 8 | 2.6 | 6 | 6 |
| Fluorine | 9 | 2.7 | 7 | 7 |
| Neon | 10 | 2.8 | 8 | 8 |
| Sodium | 11 | 2.8.1 | 1 | 1 |
| Magnesium | 12 | 2.8.2 | 2 | 2 |
| Aluminium | 13 | 2.8.3 | 3 | 3 |
| Silicon | 14 | 2.8.4 | 4 | 4 |
| Phosphorus | 15 | 2.8.5 | 5 | 5 |
| Sulphur | 16 | 2.8.6 | 6 | 6 |
| Chlorine | 17 | 2.8.7 | 7 | 7 |
| Argon | 18 | 2.8.8 | 8 | 8 |
| Potassium | 19 | 2.8.8.1 | 1 | 1 |
| Calcium | 20 | 2.8.8.2 | 2 | 2 |
By arranging the first 20 electrons in as in the table above we can see that the number of valence electrons is equal to the group number of the element on the periodic table. That is true except for helium which has 2 valency electrons but is in group 8 of the periodic table. This is due to the fact that helium has 1 shell which holds a maximum of 2 electrons. So its outer shell is full just like all other group 8 elements. Group 8 is also known as group 0.
Valence electrons in relation to periodic table blocks
Valence electrons of s-block and p-block elements
The electronic configuration of an atom can be denoted according to the number of electrons in the subshells. Using the sub-shell notation, the electron fill up in the following sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.
We are not going to go into much detail about how to fill sub-shells because that is the subject of another chapter which I assume you went through before reading this chapter.
| Element | Number of electrons | Electronic Configuration (showing sub-shells) | Periodic Table Block | Valence |
|---|---|---|---|---|
| Hydrogen | 1 | 1s1 | s-block | 1 |
| Helium | 2 | 1s2 | s-block | 2 |
| Lithium | 3 | 1s2 2s1 | s-block | 1 |
| Beryllium | 4 | 1s2 2s2 | s-block | 2 |
| Boron | 5 | 1s2 2s2 2p1 | p-block | 3 |
| Carbon | 6 | 1s2 2s2 2p2 | p-block | 4 |
| Nitrogen | 7 | 1s2 2s2 2p3 | p-block | 5 |
| Oxygen | 8 | 1s2 2s2 2p4 | p-block | 6 |
| Fluorine | 9 | 1s2 2s2 2p5 | p-block | 7 |
| Neon | 10 | 1s2 2s2 2p6 | p-block | 8 |
| Sodium | 11 | 1s2 2s2 2p6 3s1 | s-block | 1 |
| Magnesium | 12 | 1s2 2s2 2p6 3s2 | s-block | 2 |
| Aluminium | 13 | 1s2 2s2 2p6 3s2 3p1 | p-block | 3 |
| Silicon | 14 | 1s2 2s2 2p6 3s2 3p2 | p-block | 4 |
| Phosphorus | 15 | 1s2 2s2 2p6 3s2 3p3 | p-block | 5 |
| Sulphur | 16 | 1s2 2s2 2p6 3s2 3p4 | p-block | 6 |
| Chlorine | 17 | 1s2 2s2 2p6 3s2 3p5 | p-block | 7 |
| Argon | 18 | 1s2 2s2 2p6 3s2 3p6 | p-block | 8 |
| Potassium | 19 | 1s2 2s2 2p6 3s2 3p6 4s1 | s-block | 1 |
| Calcium | 20 | 1s2 2s2 2p6 3s2 3p6 4s2 | s-block | 2 |
As we can see from the table above:
- p-block elements have less than 3 valency electrons
- s-block elements have 3 or more valency electrons
For s-block elements, the outer sub-shell is the s orbital which is made up of 1 sub-orbital. Since a single sub-orbital holds a maximum of 2 electrons, s-block elements are limited to a maximum of only 2 electrons.
However for p-block elements, the outer sub-shell is the p orbital which is made up of 3 sub-orbitals. 3 sub-orbitals can hold a maximum of 6 electrons. Since every p orbital has a corresponding s orbital in the same shell with it, the number of the valency electrons of a p-block element is the sum of the p orbital and corresponding s orbital electrons. For example, Oxygen has 6 valence electrons from 2 electrons in the s orbital (2s2) and 4 electrons in the p orbital (2p4). All of them in the second shell.
Valence electrons of d-block elements
For s-block and p-block elements valence electrons exist only in the outermost electron shell but however for d-block elements (transition metals), valence electrons can also exist in an inner shell.
Lets start by looking at the electronic configurations of the first 10 d-block elements.
| Element | Number of electrons | Electronic Configuration |
|---|---|---|
| Sc | 21 | [Ar] 4s2 3d1 |
| Ti | 22 | [Ar] 4s2 3d2 |
| V | 23 | [Ar] 4s2 3d3 |
| Cr | 24 | [Ar] 4s1 3d5 |
| Mn | 25 | [Ar] 4s2 3d5 |
| Fe | 26 | [Ar] 4s2 3d6 |
| Co | 27 | [Ar] 4s2 3d7 |
| Ni | 28 | [Ar] 4s2 3d8 |
| Cu | 29 | [Ar] 4s1 3d10 |
| Zn | 30 | [Ar] 4s2 3d10 |
Transition elements have a partially filled (n − 1)d energy level (shell) adjacent to the ns energy level. As you can see on the table above, those elements have a partially filled 3d energy level next to a full 4s energy level (except in the cases of Cu and Cr which also have partially filled 4s energy level).
On the table, we used the electronic configuration of Argon for the shorthand notation. For example, manganese (Mn) which has an electronic configuration of 1s2 2s2 2p6 3s2 3p6 4s2 3d5 is abbreviated as [Ar] 4s2 3d5, where [Ar] represents the configuration of argon (1s2 2s2 2p6 3s2 3p6).
Another important thing to note about the short hand we used in the table, all the electrons outside the “Argon configuration” can behave as valence electrons.
Going back again to the configuration of manganese (Mn) → [Ar] 4s2 3d5, each 3d electron has the same energy level as a 4s electron, and a much higher energy level than that of 3s or 3p electrons. This leads to manganese having up to seven valence electrons (4s2 3d5), hence its oxidation state being as high as +7 (such as in the permanganate ion, MnO−4).
The effect of valence electrons in chemical reactions
The number of valence electrons that an atom has determines its chemical properties.
For example:
- alkali metals (group 1) are generally the most reactive metals because each atom has only 1 valence electron. During the formation of an ionic bond 1 valence electron is easier to lose when forming a positive ion.
- alkaline earth metals (group 2) are a bit less reactive, because each atom must lose 2 valence electrons (which requires a bit more energy) to form a positive ion.
- in both group 1 and group 2 the reactivity of the elements increases down the group. The more the electron shells an element has the further the valence electrons are from the attraction of the nucleus and the easier they are to remove.
- a nonmetal atom reacts reacts in two ways to attain a full valence shell. It can either share electrons with another atom (to form a covalent bond), or it can pull electrons from another atom (to form an ionic bond). The closer the the valence electrons are to the nucleus the greater the sharing or pulling power and more reactive the atom is.
- halogens (group 7) are the most reactive nonmetal elements because each atom only needs to pull or share one electron to form a full valence shell.
- however in non metal groups the reactivity decreases down the group because the valency shells are further away from nucleus and therefore the “pulling power” of the nucleus is weakened.
- oxygen (group 6), though it requires 2 electron to complete its valence shell, is the most reactive nonmetal after fluorine. This is because its valence shell is closer to the nucleus than most halogens.