What is the difference between the first and second ionisation energy of an element?

The first ionisation energy is the energy to remove an electron from a neutral gaseous atom, while the second is for removing an electron from a $1+$ gaseous ion. The second IE is always higher because the remaining electrons are more strongly attracted to the nucleus in a positive ion.

How does the trend in ionisation energy differ from the trend in atomic radius across a period?

They are inversely related: as you move across a period, the atomic radius decreases while the ionisation energy increases. This is because the increasing nuclear charge pulls electrons closer and holds them more tightly, making them harder to remove.

Why does Boron have a lower first ionisation energy than Beryllium?

Beryllium's outer electron is in a $2s$ subshell, while Boron's is in a $2p$ subshell. The $2p$ subshell is higher in energy and further from the nucleus, and it is also shielded by the $2s$ electrons, making the $2p$ electron easier to remove.

What is a common mistake when writing the equation for the second ionisation energy?

A common error is starting with a neutral atom ($X$) instead of a $1+$ ion ($X^+$). The correct equation must be $X^+_{(g)} \rightarrow X^{2+}_{(g)} + e^-$ to represent the removal of the second electron specifically.

Why is it incorrect to define ionisation energy using solid or liquid states?

In solids or liquids, energy would be required to overcome intermolecular forces or lattice energies in addition to removing the electron. Using the gaseous state ensures the measurement reflects only the atomic property of electron attraction.

What error do students often make when interpreting successive ionisation energy graphs?

Students often read the graph from right to left or miscount the 'jumps'. It is essential to count the number of electrons removed starting from the first (lowest energy) to identify which shell the electrons are coming from.

Define the First Ionisation Energy.

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous $1+$ ions. It is measured in units of kJ/mol.

What is the 'Shielding Effect'?

The reduction in the effective nuclear charge on the electron cloud, caused by the inner shells of electrons repelling the outer valence electrons. This weakens the nucleus's hold on the outer electrons.

What does a 'big jump' in a successive ionisation energy plot represent?

A big jump represents the removal of an electron from a new principal energy level (shell) that is closer to the nucleus. This indicates that all valence electrons from the outer shell have already been removed.

Why does ionisation energy decrease down a group despite an increasing nuclear charge?

The increase in atomic radius and the additional shielding from inner electron shells outweigh the increase in nuclear charge. This results in a weaker net attraction between the nucleus and the valence electrons.

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3. Periodic Table & Energy

Ionisation Energy

Summary

Ionisation energy is a fundamental periodic property that measures the energy required to remove electrons from gaseous atoms or ions. It provides direct evidence for the existence of electronic shells and subshells, as the energy required varies predictably based on nuclear attraction, atomic radius, and electron shielding.

1. Definition & Core Concepts

Graph showing the trend of first ionisation energy across Period 2, highlighting the general increase and the specific dips at Boron and Oxygen.

2. Underlying Principles

3. Factors Influencing Ionisation Energy

Atomic Radius: As the distance between the nucleus and the outer electron increases, the electrostatic force of attraction weakens significantly, making the electron easier to remove and lowering the ionisation energy.
Nuclear Charge: A higher number of protons in the nucleus creates a stronger positive pull on the electrons. If all other factors are constant, an increase in nuclear charge leads to a higher ionisation energy.
Electron-Electron Repulsion: Within the same subshell, paired electrons in an orbital experience extra repulsion compared to unpaired electrons. This repulsion makes it slightly easier to remove one of the paired electrons.

4. Periodic Trends

Across a Period: Ionisation energy generally increases. This occurs because the nuclear charge increases while the shielding remains relatively constant (electrons are added to the same shell), leading to a smaller atomic radius and stronger attraction.
Down a Group: Ionisation energy decreases. Although the nuclear charge increases, the addition of new principal energy shells increases the atomic radius and the shielding effect, which outweighs the increase in proton number.
Period Transitions: There is a sharp drop in ionisation energy when moving from a noble gas at the end of one period to an alkali metal at the start of the next, due to the electron being placed in a new shell further from the nucleus.

5. Successive Ionisation Energies

6. Key Distinctions & Anomalies

7. Exam Strategy & Tips