What is the difference between the first and second ionisation energy in terms of the starting species?

The first ionisation energy starts with a neutral gaseous atom ($X$), while the second ionisation energy starts with a gaseous $1+$ ion ($X^+$). Both processes remove exactly one mole of electrons to form a more positively charged gaseous ion.

How does the trend in ionisation energy across a period differ from the trend down a group?

Across a period, ionisation energy generally increases due to increasing nuclear charge and constant shielding. Down a group, it decreases because the increasing atomic radius and shielding from additional shells outweigh the increase in nuclear charge.

Why is the first ionisation energy of Oxygen lower than that of Nitrogen, despite Oxygen having more protons?

Oxygen has a pair of electrons in one of its $2p$ orbitals, whereas Nitrogen has only unpaired electrons in its $2p$ subshell. The spin-pair repulsion between the two electrons in Oxygen's filled orbital makes it easier to remove one of them.

What is the most common error when writing the chemical equation for the first ionisation energy of Magnesium?

The most common error is omitting the gaseous state symbols $(g)$. The definition of ionisation energy specifically requires the species to be in the gaseous state: $Mg(g) \rightarrow Mg^+(g) + e^-$.

Why is it incorrect to say that ionisation energy decreases down a group because the nucleus gets weaker?

The nucleus actually gets stronger (more protons) down a group. The decrease in ionisation energy is due to the increased distance and shielding provided by extra electron shells, which reduce the effective nuclear pull on the valence electrons.

What mistake is made when identifying an element's group from a jump between the 2nd and 3rd ionisation energies?

Students often miscount the valence electrons. If the jump is between the 2nd and 3rd IE, it means the first 2 electrons were easy to remove (valence), so the element belongs to Group 2, not Group 3.

What are the standard conditions used for measuring ionisation energies?

Standard conditions are typically a temperature of $298$ K ($25$ degrees Celsius) and a pressure of $101$ kPa ($1$ atmosphere).

What is the general formula for the $n^{th}$ ionisation energy?

The formula is $X^{(n-1)+}(g) \rightarrow X^{n+}(g) + e^-$. This shows the removal of one electron from an ion that has already lost $n-1$ electrons.

Why do successive ionisation energies always increase for a given element?

As electrons are removed, the remaining electrons experience less repulsion and are pulled closer to the nucleus by the same number of protons. Removing an electron from a positive ion is more difficult than from a neutral atom.

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1. Physical Chemistry

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 quantum shells and subshells within the atomic structure.

1. Definition & Core Concepts

2. Underlying Principles: Factors Affecting IE

Nuclear Charge: A higher number of protons in the nucleus creates a stronger positive charge, increasing the electrostatic attraction for the outer electrons and raising the ionisation energy.
Atomic Radius (Distance): As the distance between the nucleus and the outer electron increases, the electrostatic force of attraction weakens significantly according to Coulomb's Law, making the electron easier to remove.
Shielding Effect: Inner shell electrons repel outer electrons and 'block' the attractive force of the nucleus. More internal shells result in greater shielding and a lower ionisation energy.
Spin-Pair Repulsion: Within an orbital, two electrons with opposite spins experience a degree of magnetic and electrostatic repulsion. This repulsion makes it slightly easier to remove one of the paired electrons compared to an unpaired electron in the same subshell.

3. Periodic Trends

Across a Period

General Trend: Ionisation energy generally increases across a period. This occurs because the nuclear charge increases while the shielding remains relatively constant (as electrons are added to the same principal shell), pulling the outer electrons closer and tighter to the nucleus.

Down a Group

General Trend: Ionisation energy decreases down a group. Although the nuclear charge increases, the addition of new principal quantum shells significantly increases the atomic radius and the shielding effect, which outweighs the increase in proton number.

Graph showing the general increase of ionisation energy across a period with characteristic dips at Group 3 and Group 6.

4. Key Distinctions: Trend Anomalies

5. Exam Strategy & Tips

Ionisation Energy

Q: 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 under standard conditions.

Summary

1. Definition & Core Concepts

First Ionisation Energy ( $IE_1$ ) is defined as the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous $1+$ ions. This process is endothermic, as energy must be supplied to overcome the electrostatic attraction between the nucleus and the outer electron.

The measurement must occur under standard conditions (typically $298$ K and $101$ kPa) to ensure consistency across different elements. The standard units for ionisation energy are kilojoules per mole ( $kJ \, mol^{-1}$ ).

The general equation for the first ionisation of an element $X$ is represented as:

$X(g) \rightarrow X^+(g) + e^-$

Successive Ionisation Energies refer to the removal of subsequent electrons from the same atom (e.g., $IE_2$ , $IE_3$ ). These values always increase because removing a negative electron from an increasingly positive ion requires significantly more energy due to reduced electron-electron repulsion and a higher proton-to-electron ratio.