Periods 1 to 3 Ions Shell Electron Configuration
Key Concepts
- An ion is formed when an atom gains or loses 1 or more electrons.
- A cation is a positively charged ion.
- An anion is a negatively charged ion.
- A cation is formed when an atom loses 1 or more electrons.
- An anion is formed when an atom gains 1 or more electrons.
- The electron configuration (electronic configuration) of Group 18 (Noble gas) atoms is stable, that is, Group 18 elements do not readily form ions.
- An atom will gain or lose electrons in order to achieve the same stable electron configuration as a Group 18 (Noble gas) atom.1
- A cation is formed by an atom losing electrons from the outer shell of electrons (valence shell).
- An anion is formed by adding electrons to the outer shell of electrons (valence shell).
Concepts: Ions and Electronic Configuration
The number of electrons (negative charges) in an atom of an element is equal to the number of protons (positive charges) in its nucleus.
The number of protons in the nucleus of an atom is given by its atomic number (Z) which can be found in the Periodic Table.
The electron configuration (electronic configuration) of an atom is written such that the number of electrons in the lowest energy level or shell is written first, followed by the number of electrons in the next lowest energy level or shell.
The electronic configuration of an atom of a Group 18 (Noble Gas) element is stable, that is, Group 18 elements do not readily form compounds:
| |
|
No ions formed |
|---|
| Group 1 |
Group 2 |
|
Group 13 |
Group 14 |
Group 15 |
Group 16 |
Group 17 |
Group 18
(Noble Gas) |
| Period 1 |
1
H |
|
2
He |
|---|
| Period 2 |
3
Li |
4
Be |
5
B |
6
C |
7
N |
8
O |
9
F |
10
Ne |
|---|
| Period 3 |
11
Na |
12
Mg |
13
Al |
14
Si |
15
P |
16
S |
17
Cl |
18
Ar |
|---|
The electronic configuration for the first three Noble gases is shown below with the highest energy level, the valence shell, highlighted in light blue:
Group 18 atom
(Noble Gas atom) |
Electronic Configuration |
|---|
| 1st energy level K shell
(2 electrons maximum) |
2nd energy level L shell
(8 electrons maximum) |
3rd energy level M shell |
|---|
| He |
2 |
|
|
| Ne |
2, |
8 |
|
| Ar |
2, |
8, |
8 |
The valence shell of atoms of Group 18 (Noble Gas) elements is full therefore there is no need for a Group 18 atom to form ions by gaining or losing or electrons, the atom has already achieved the most stable electronic configuration possible.
When an atom with only a small number of electrons in its valence shell forms an ion, it loses the electrons from its valence shell to form a positively charged ion, a cation, with the same electronic configuration as the preceding Noble gas (Group 18 element) in the Periodic Table.
Elements on the left-hand side of the Periodic Table, metallic elements (Groups 1 and 2, as well as aluminium in Group 3), lose electrons to form positively charged ions (cations).
| |
form cations |
|
|---|
| Group 1 |
Group 2 |
|
Group 13 |
Group 14 |
Group 15 |
Group 16 |
Group 17 |
Group 18 |
| Period 2 |
3
Li |
4
Be |
5
B |
6
C |
7
N |
8
O |
9
F |
10
Ne |
|---|
| Period 3 |
11
Na |
12
Mg |
13
Al |
14
Si |
15
P |
16
S |
17
Cl |
18
Ar |
|---|
When the valence shell of an atom is almost full, it forms an ion by adding electrons to its valence shell to form a negatively charged ion with the same electronic configuration as the next Noble gas (Group 18 element) in the Periodic Table.
Atoms of the non-metallic elements on the right-hand side of the Periodic Table, (Groups 15, 16 and 17 but NOT Group 18), gain electrons to form negatvely charged ions (anions).
| |
|
form anions |
|
|---|
| Group 1 |
Group 2 |
|
Group 13 |
Group 14 |
Group 15 |
Group 16 |
Group 17 |
Group 18 |
| Period 2 |
3
Li |
4
Be |
5
B |
6
C |
7
N |
8
O |
9
F |
10
Ne |
|---|
| Period 3 |
11
Na |
12
Mg |
13
Al |
14
Si |
15
P |
16
S |
17
Cl |
18
Ar |
|---|
If the electronic configuration of an atom results in the valence shell being half-filled, the atom may either gain or lose electrons in order to form an ion, as in the case of hydrogen, or it may not readily form ions and prefer to form covalent compounds as in the case of carbon and silicon (note that hydrogen also forms covalent bonds!).
Period 1: Electron Configuration of Hydrogen and its Ions
The electron configuration for an atom of hydrogen is 1
An atom of hydrogen has only 1 electron and this electron occupies the first energy level (K shell).
| ↑ ↑ ↑ ↑ | 3rd energy level
(M shell) | ____ | |
| ↑ ↑ | | |
| ↑ ↑ ↑ | 2nd energy level
(L shell) | ____ | (maximum of 8 electrons) |
| ↑ | | |
| ↑ | 1st energy level
(K shell) |
↑
____ | (maximum of 2 electrons) |
↑
energy | levels
(shells) | | |
The maximum number of electrons that can be accomodated in the K shell is 2, so the electronic configuration of hydrogen shows us that this energy level (shell) is only half-filled.
This means that hydrogen can either:
- lose 1 electron to form a positively charged ion (cation)
- gain 1 electron to form a negatively charged ion (anion)
If an atom of hydrogen loses its valence electron it forms the H+ ion, and the H+ ion has no electrons at all!
| ↑ ↑ ↑ ↑ | 3rd energy level
(M shell) | ____ | |
| ↑ ↑ | | |
| ↑ ↑ ↑ | 2nd energy level
(L shell) | ____ | (maximum of 8 electrons) |
| ↑ | | |
| ↑ | 1st energy level
(K shell) |
____ | (maximum of 2 electrons) |
↑
energy | levels
(shells) | | |
But, the nucleus of the H+ still contains 1 positively charged proton, hence the overall charge on the ion is +1.
The most common naturally occurring hydrogen atoms have a nucleus containing 1 proton and 0 neutrons, so when the hydrogen atom loses its electron to form H+ the species is just a proton!
For this reason Chemists often refer to "protons" being involved in chemical reactions2 when they mean H+.
And, if you are thinking that a bare proton would be a very reactive species, you'd be right!
If an atom of hydrogen gains 1 electron, this electron enters the valence shell (K shell) so that the ion has 2 electrons in the K shell:
| ↑ ↑ ↑ ↑ | 3rd energy level
(M shell) | ____ | |
| ↑ ↑ | | |
| ↑ ↑ ↑ | 2nd energy level
(L shell) | ____ | (maximum of 8 electrons) |
| ↑ | | |
| ↑ | 1st energy level
(K shell) |
↑↓
____ | (maximum of 2 electrons) |
↑
energy | levels
(shells) | | |
This ion still has a nucleus containing 1 positively charged proton, but now it has 2 negatively charged electrons surrounding the nucleus, so the overall charge is -1, and the ion is H-.
The electronic configuration of H- is 2
That is, by gaining 1 electron to form H-, the hydrogen atom has filled the K shell resulting in a stable electronic configuration.
The electronic configuration of helium (He) is also 2
By gaining 1 electron to form H- hydrogen achieves the stable electronic configuration of an atom of helium.
Group 1: Electronic Configuration of Atoms and Ions
Group 1 (alkali metal) elements have only 1 electron in their valence shell (highest energy level). This valence electron is highlighted in light blue in the table below:
Group 1 atom
(alkali metal atom) |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| Li |
2, |
1 |
|
| Na |
2, |
8, |
1 |
An atom of a Group 1 element will easily lose its valence electron in order to form a positively charged ion (cation) in which the shell(s) are filled with electrons, that is, with the same stable electronic configuration as an atom of the preceding Noble gas (Group 18) element.
Since the ion will have 1 more positively charged proton in its nucleus than it has negatively charged electrons surrounding the nucleus, Group 1 atoms form cations with a charge of +1:
Group 1 cation
(alkali metal cation) |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| Li+ |
2, |
|
|
| Na+ |
2, |
8, |
|
Note that Li+ has the electronic configuration 2 which is the same as the electronic configuration of an atom of the preceding Group 18 element helium.
Na+ has the electronic configuration 2,8 which is the same as the electronic configuration of an atom of the preceding Group 18 element neon.
Group 2: Electronic Configuration of Atoms and Ions
An atom of a Group 2 element (alkali earth metal or alkaline earth metal) has 2 electrons in its valence shell (highest energy level):
Group 2 atom
(alkaline-earth metal atom) |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| Be |
2, |
2 |
|
| Mg |
2, |
8, |
2 |
Atoms of Group 2 elements can lose these 2 valence electrons to form cations in which the shells containing electrons are full, in other words, these ions have the same stable electronic configuration as an atom of a Group 18 (Noble Gas) element.
The ion of a Group 2 element will therefore have 2 more positively charged protons in its nucleus than it has negatively charged electrons surrounding its nucleus, so the overall charge on the ion will be +2:
Group 2 cation
(alkaline-earth metal cation) |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| Be2+ |
2, |
|
|
| Mg2+ |
2, |
8, |
|
The electronic configuration of Be2+ is 2 which is the same as the electronic configuration of an atom of the preceding Noble gas (Group 18) element, helium.
The electronic configuration of Mg2+ is 2,8 which is the same as the electronic configuration of an atom of the preceding Noble gas (Group 18) element, neon.
Group 13: Electronic Configuration of Atoms and Ions
An atom of a Group 13 element has 3 electrons in its valence shell (highest energy level):
| Group 13 atom |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| B |
2, |
3 |
|
| Al |
2, |
8, |
3 |
Atoms of Group 13 elements can lose these 3 valence electrons to form cations in which the shells containing electrons are full, that is, they have the same stable electronic configuration as an atom of a Group 18 (Noble Gas) element.
The ion of a Group 13 element will therefore have 3 more positively charged protons in its nucleus than it has negatively charged electrons surrounding its nucleus, so the overall charge on the ion will be +3:
| Group 13 cation |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| B3+ |
2, |
|
|
| Al3+ |
2, |
8, |
|
The electronic configuration of B3+ is 2 which is the same as the electronic configuration of an atom of the preceding Noble gas (Group 18) element, helium.
The electronic configuration of Al3+ is 2,8 which is the same as the electronic configuration of an atom of the preceding Noble gas (Group 18) element, neon.
Group 15: Electron Configuration of Atoms and Ions
An atom of a Group 15 non-metallic element has 5 electrons in its valence shell (highest energy level):
| Group 15 atom |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| N |
2, |
5 |
|
| P |
2, |
8, |
5 |
Atoms of Group 15 elements can gain 3 electrons to form anions which will have a valence shell filled with electrons, that is, the same stable electronic configuration as an atom of a Group 18 (Noble Gas) element.
These 3 electrons are added to the valence shell (the highest energy level available).
The ion of a Group 15 element will therefore have 3 more negatively charged electrons surrounding its nucleus than it has positively charged protons in its nucleus, so the overall charge on the ion will be 3-:
| Group 15 anion |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| N3- |
2, |
8 |
|
| P3- |
2, |
8, |
8 |
The electronic configuration of N3- is 2,8 which is the same as the electronic configuration of an atom of the next Noble gas (Group 18) element in the Periodic Table, neon.
The electronic configuration of P3- is 2,8,8 which is the same as the electronic configuration of an atom of the next Noble gas (Group 18) element in the Periodic Table, argon.
Group 16: Electronic Configuration of Atoms and Ions
An atom of a Group 16 non-metallic element has 6 electrons in its valence shell (highest energy level):
| Group 16 atom |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| O |
2, |
6 |
|
| S |
2, |
8, |
6 |
Atoms of Group 16 elements can gain 2 electrons to form anions which will have a valence shell filled with electrons, that is, the same stable electronic configuration as an atom of a Group 18 (Noble Gas) element.
These 2 electrons are added to the valence shell.
The ion of a Group 16 element will therefore have 2 more negatively charged electrons surrounding its nucleus than it has positively charged protons in its nucleus, so the overall charge on the ion will be 2-:
| Group 16 anion |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| O2- |
2, |
8 |
|
| S2- |
2, |
8, |
8 |
The electronic configuration of O2- is 2,8 which is the same as the electronic configuration of an atom of the next Noble gas (Group 18) element in the Periodic Table, neon.
The electronic configuration of S2- is 2,8,8 which is the same as the electronic configuration of an atom of the next Noble gas (Group 18) element in the Periodic Table, argon.
Group 17: Electron Configuration of Atoms and Ions
An atom of a Group 17 (halogen) non-metallic element has 7 electrons in its valence shell (highest energy level):
| Group 16 atom |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| F |
2, |
7 |
|
| Cl |
2, |
8, |
7 |
Atoms of Group 17 elements can gain 1 electron to form anions which will have a full valence shell of electrons, that is, the same stable electronic configuration as an atom of a Group 18 (Noble Gas) element..
This electron is added to the valence shell.
The ion of a Group 17 element will therefore have 1 more negatively charged electron surrounding its nucleus than it has positively charged protons in its nucleus, so the overall charge on the ion will be 1-:
| Group 17 anion |
Electronic Configuration |
|---|
| 1st energy level K shell |
2nd energy level L shell |
3rd energy level M shell |
|---|
| F- |
2, |
8 |
|
| Cl- |
2, |
8, |
8 |
The electronic configuration of F- is 2,8 which is the same as the electronic configuration of an atom of the next Noble gas (Group 18) element in the Periodic Table, neon.
The electronic configuration of Cl- is 2,8,8 which is the same as the electronic configuration of an atom of the next Noble gas (Group 18) element in the Periodic Table, argon.
Summary: Electronic Configuration of Atoms and Ions Using Shell Notation
Group 18 (Noble gas) elements have a full valence shell of electrons and are stable.
Atoms of other elements do not have a full valence shell of electrons and may form ions in order to achieve a stable Noble gas electronic configuration.
An atom that has few electrons in its valence shell will lose those valence electrons to form a cation.
An atom that has an almost full valence shell will gain the number of electrons required to complete this valence shell which results in an anion.
The table below summarises the electronic configurations of Periods 1 to 3 atoms and their corresponding ions.
| |
Atom |
Electronic
Configuration |
|
Ion |
Electronic
Configuration |
|---|
| Period 1 |
H |
1 |
|
H+ H- |
2 |
| Period 2 |
Li |
2,1 |
|
Li+ |
2 |
| Be |
2,2 |
|
Be2+ |
2 |
| B |
2,3 |
|
B3+ |
2 |
| N |
2,5 |
|
N3- |
2,8 |
| O |
2,6 |
|
O2- |
2,8 |
| F |
2,7 |
|
F- |
2,8 |
| Period 3 |
Na |
2,8,1 |
|
Na+ |
2,8 |
| Mg |
2,8,2 |
|
Mg2+ |
2,8 |
| Al |
2,8,3 |
|
Al3+ |
2,8 |
| P |
2,8,5 |
|
P3- |
2,8,8 |
| S |
2,8,6 |
|
S2- |
2,8,8 |
| Cl |
2,8,7 |
|
Cl- |
2,8,8 |
Worked Example
Question: The ion of element X is known to have the ground state electronic configuration 2,8 and a charge of +3.
What is the ground state electronic configuration of an atom of element X?
Solution (Based on the StoPGoPS method for problem solving)
| STOP |
STOP! State the Question.
|
| |
What is the question asking you to do?
Write the ground state electronic configuration for an atom of element X
|
| PAUSE |
PAUSE to Prepare a Game Plan
|
| |
- What data have you been given?
(1) Electronic configuration of the ion of X : 2,8
(2) Charge on an ion of element X: 3+
- What is the relationship between what you have been given and what you need to find?
(1) Since the ion is positively charged, the atom must have lost electrons.
(2) When forming a cation, an atom loses electrons from its valence shell.
(3) The atom must have had more electrons in a higher energy level (valence shell) than the ion.
|
| GO |
GO with the Game Plan |
| |
Ion has a charge of 3+ so the atom must have lost 3 electrons.
Electronic configuration of X3+ is 2,8 (highest energy level is full)
The atom has 3 more electrons than the ion.
These 3 extra electrons must have occupied the valence shell (highest energy level) of the atom.
The 3 electrons lost from the atom must have occupied the 3rd energy level (M shell)
Electron configuration of an atom of X is 2,8,3
|
| PAUSE |
PAUSE to Ponder Plausibility |
| |
Is your answer plausible?
Work backwards: If the electronic configuration of an atom of X is 2,8,3 then if the atom loses 3 electrons from its valence shell to form an ion with a +3 charge (X3+) then the electronic configuration of X3+ will be 2,8
Since the electronic configuration for X3+ is the same as that given in the question, we are confident our answer is correct.
|
| STOP |
STOP! State the Solution |
| |
State your solution to the problem.
The electronic configuration for an atom of element X in the ground state is 2,8,3
|
1. Species with the same electronic configuration are said to be isoelectronic.
2. IUPAC prefers the term "hydron" for the H+ species and "hydride" for the H- species. These names better reflect the fact that the nuclei of some hydrogen atoms (and therefore ions) do contain neutrons so that the H+ species may not necessarily be a bare proton. However, given the preponderance of hydrogen atoms with 0 neutrons that occur in nature, and the fact that the term "proton" has been in use for a very long time, most people will still refer to "protons" rather than "hydrons".
