3.1.1a-c - Periodic table, ionisation energy and electron configuration
The periodic table is more than a list of element symbols. It is a map built from atomic number, electron configuration and repeated chemical behaviour. In this lesson you will connect the layout of the table to Period 2 and Period 3 electron configurations, then use first and successive ionisation energy data as evidence for shells and sub-shells.
The periodic table as an evidence-based arrangement
The modern periodic table arranges elements in order of increasing atomic number. Atomic number means proton number, so the sequence is based on the number of protons in each atom's nucleus.
Period
A period is a horizontal row in the periodic table. Elements in the same period show a repeating pattern in physical and chemical properties as atomic number increases.
Group
A group is a vertical column in the periodic table. Elements in the same group usually have similar chemical properties because their atoms have similar outer-shell electron configurations.
Periodicity means a repeating pattern in properties at regular intervals. For example, Period 2 and Period 3 both begin with a reactive Group 1 metal and end with a noble gas. The details differ, but the same kind of pattern returns because the outer-shell electron patterns repeat.
The periodic table was accepted because it worked as a scientific model. Early versions arranged elements by patterns in properties and left gaps where unknown elements should fit. As new elements were discovered and their properties matched predictions, the Periodic Law became stronger evidence rather than just a neat arrangement. The table is still extended when new elements are discovered, tested and confirmed by the scientific community.
Worked reasoning: correcting a common statement
A student says, "The position of an element is decided by relative atomic mass."
That is not the modern rule. The modern periodic table is ordered by atomic number. Argon and potassium are a useful warning from earlier chemistry: argon has a greater relative atomic mass than potassium, but argon has atomic number 18 and potassium has atomic number 19, so argon comes before potassium.
Electron configuration patterns in Periods 2 and 3
Across a period, electrons are added one at a time to the same main shell. Period 2 fills the second shell; Period 3 fills the third shell. The repeating pattern is easiest to see if you write the outer part of each configuration.
| Element | Atomic number | Electron configuration | Outer-shell pattern |
|---|---|---|---|
| Li | 3 | 1s^2 2s^1 | 2s^1 |
| Be | 4 | 1s^2 2s^2 | 2s^2 |
| B | 5 | 1s^2 2s^2 2p^1 | 2s^2 2p^1 |
| C | 6 | 1s^2 2s^2 2p^2 | 2s^2 2p^2 |
| N | 7 | 1s^2 2s^2 2p^3 | 2s^2 2p^3 |
| O | 8 | 1s^2 2s^2 2p^4 | 2s^2 2p^4 |
| F | 9 | 1s^2 2s^2 2p^5 | 2s^2 2p^5 |
| Ne | 10 | 1s^2 2s^2 2p^6 | 2s^2 2p^6 |
Period 3 repeats the same outer-shell pattern, but with 3s and 3p sub-shells:
| Element | Atomic number | Electron configuration | Outer-shell pattern |
|---|---|---|---|
| Na | 11 | 1s^2 2s^2 2p^6 3s^1 | 3s^1 |
| Mg | 12 | 1s^2 2s^2 2p^6 3s^2 | 3s^2 |
| Al | 13 | 1s^2 2s^2 2p^6 3s^2 3p^1 | 3s^2 3p^1 |
| Si | 14 | 1s^2 2s^2 2p^6 3s^2 3p^2 | 3s^2 3p^2 |
| P | 15 | 1s^2 2s^2 2p^6 3s^2 3p^3 | 3s^2 3p^3 |
| S | 16 | 1s^2 2s^2 2p^6 3s^2 3p^4 | 3s^2 3p^4 |
| Cl | 17 | 1s^2 2s^2 2p^6 3s^2 3p^5 | 3s^2 3p^5 |
| Ar | 18 | 1s^2 2s^2 2p^6 3s^2 3p^6 | 3s^2 3p^6 |
This pattern explains why groups have similar chemical properties. Lithium and sodium both have one outer s electron. Fluorine and chlorine both have outer-shell pattern ns^2 np^5, so each is one electron short of a full outer shell.
The block of an element is named from the type of sub-shell being filled by its highest-energy electron in the atom:
| Block | Where it appears in the table | Typical outer configuration pattern |
|---|---|---|
| s-block | Groups 1 and 2, plus helium by configuration | ns^1 or ns^2 |
| p-block | Main groups on the right, excluding helium | ns^2 np^1 to ns^2 np^6 |
| d-block | Central block beginning in Period 4 | d sub-shell being filled |
For this lesson, the most important link is Periods 2 and 3: the first two elements are s-block, then the next six are p-block.
Worked example: identify the block
Element X has electron configuration:
The highest-energy electron is in a 3p sub-shell, so X is in the p-block. Its outer-shell pattern has five electrons, 3s^2 3p^3, so it is in the same main-group pattern as nitrogen.
First ionisation energy
Ionisation energy measures how much energy is needed to remove an electron from an atom or ion. In this lesson the formal definition you need is for first ionisation energy.
First ionisation energy
The first ionisation energy is the energy required to remove 1 mol of electrons from 1 mol of gaseous atoms to form 1 mol of gaseous 1+ ions.
For a general element X:
The state symbols matter. The atom must be gaseous, and the ion formed is gaseous. First ionisation energy values are usually given in kJ mol^-1.
For sodium:
The electron removed is the outer 3s electron. The size of the first ionisation energy depends on the attraction between that outer electron and the nucleus.
The main factors are:
- nuclear charge: more protons give a stronger attraction;
- atomic radius: an outer electron farther from the nucleus is less strongly attracted;
- attraction after inner-shell effects: inner electrons reduce the attraction felt by the outer electron.
High first ionisation energy means the outer electron is strongly attracted to the nucleus. Low first ionisation energy means it is easier to remove the outer electron.
Worked example: energy for one ion
The first ionisation energy of lithium is about 520 kJ mol^-1. This means 520 kJ is needed to form 1 mol of Li^+(g) ions from 1 mol of Li(g) atoms.
To estimate the energy for one Li(g) atom:
To two significant figures:
This is a data-translation calculation: the chemistry idea is molar ionisation energy, and the mathematical move is dividing by the Avogadro constant.
Explaining first-ionisation-energy trends
First ionisation energy is periodic: it shows repeating patterns as atomic number increases.
Across Period 2 from Li to Ne, and across Period 3 from Na to Ar, first ionisation energy generally increases. The outer electron is being removed from the same main shell across a period, so there is no new outer shell. At the same time, nuclear charge increases because each element has one more proton. The atomic radius also tends to decrease. The outer electron is therefore more strongly attracted to the nucleus, so more energy is needed to remove it.
That general trend has two important small decreases in each period.
Decrease from Be to B, and from Mg to Al
Beryllium has outer configuration 2s^2. Boron has outer configuration 2s^2 2p^1. The electron removed from boron is in a 2p sub-shell, which is higher in energy and slightly easier to remove than a 2s electron. So boron's first ionisation energy is lower than beryllium's, even though boron has a higher nuclear charge.
The same idea explains Mg to Al:
- Mg: 3s^2
- Al: 3s^2 3p^1
The first electron removed from Al is a 3p electron rather than a 3s electron, so there is a small decrease from Mg to Al.
Decrease from N to O, and from P to S
Nitrogen has outer configuration 2s^2 2p^3. The three 2p electrons occupy separate p orbitals. Oxygen has outer configuration 2s^2 2p^4, so one 2p orbital contains a pair of electrons. Repulsion between paired electrons makes one of those paired p electrons easier to remove. So oxygen's first ionisation energy is lower than nitrogen's.
The same paired-electron repulsion explains P to S:
- P: 3s^2 3p^3
- S: 3s^2 3p^4
Sulfur has a pair of electrons in one 3p orbital, so removing one paired p electron is slightly easier.
Down a group, first ionisation energy generally decreases. The outer electron is in a shell with a higher principal quantum number, farther from the nucleus. Even though nuclear charge increases down the group, the increased atomic radius and reduced attraction felt by the outer electron make removal easier.
Worked explanation: why Be is higher than B
Start from the trend, then explain the exception:
- B has a higher nuclear charge than Be, so a simple trend would suggest B should be higher.
- However, Be loses a 2s electron, while B loses a 2p electron.
- The 2p electron is in a higher-energy sub-shell and is easier to remove.
- Therefore the first ionisation energy of B is lower than that of Be.
Successive ionisation energy as shell evidence
Successive ionisation energies are the energies needed to remove electrons one after another from the same atom. Each removal needs more energy than the previous one because the electron is being removed from an increasingly positive ion.
The powerful exam skill is not memorising exact numbers. It is spotting the large jump.
When a large jump appears, the next electron is being removed from an inner shell closer to the nucleus. Inner-shell electrons experience a much stronger attraction, so their ionisation energies are much larger.
Method for successive-ionisation data
- Read the values in order.
- Find the first very large jump.
- Count the number of electrons removed before that jump.
- That count is the number of outer-shell electrons.
- Use the total number of electrons, if known, to build the shell pattern and identify the main group.
For main-group elements in the school-course grouping, one outer-shell electron means Group 1, two means Group 2, five means Group 5, seven means Group 7, and a full outer shell means Group 0. Modern IUPAC group numbers are also used in chemistry, but the shell-count logic is the same.
Worked example: identify a Period 3 element from successive ionisation energies
An element in Period 3 has these first eight successive ionisation energies, in kJ mol^-1:
| Electron removed | 1st | 2nd | 3rd | 4th | 5th | 6th | 7th | 8th |
|---|---|---|---|---|---|---|---|---|
| Ionisation energy | 1012 | 1903 | 2912 | 4957 | 6274 | 21269 | 25398 | 29855 |
The first large jump is between the 5th and 6th ionisation energies:
So five electrons are removed before the atom reaches an inner shell. The atom has five outer-shell electrons.
Because the element is in Period 3, its shell pattern is:
That corresponds to phosphorus, P. In main-group course language it is Group 5; in IUPAC numbering it is Group 15.
Explain It Back
Use this as a self-explanation check after the section above. It is for diagnosing what you can already explain, not for learning new material from scratch.
Successive ionisation energy data support a shell model of the atom. If all electrons were equally easy to remove, the graph would rise smoothly. Instead, very large jumps appear when removal moves into a shell closer to the nucleus. Smaller irregularities within a shell also support the idea of sub-shells and paired-electron repulsion.