Lesson 1.4: Ionisation Energies as Evidence for Structure

Introduction

Welcome to Lesson 1.4 of Foundation Chemistry! In this lesson, we will explore ionisation energies and how they serve as evidence for the structure of atoms. By the end of this lesson, you will be able to:

Define first and successive ionisation energies with correct state symbols and equations.
Identify the factors affecting ionisation energy, including nuclear charge, atomic radius, and electron shielding/repulsion.
Understand the trends in first ionisation energy across a period and down a group, including the notable "dips" at Groups 13 and 16 and the reasoning behind them.
Use successive ionisation energy data to deduce the number of electrons in each shell and confirm an element's group and position in the periodic table.

Let's dive in! 🚀

1. Definition of Ionisation Energies

1.1 First and Successive Ionisation Energies

Ionisation energy refers to the energy required to remove an electron from an atom in the gas phase. We categorize this energy into first ionisation energy and successive ionisation energies.

The first ionisation energy is the energy needed to remove the outermost electron:

$$ \text{X(g)} \rightarrow \text{X}^+(g) + e^- $$

where X represents the atom being ionised and e^- is the electron that is removed.

The second ionisation energy can be expressed similarly:

$$ \text{X}^+(g) \rightarrow \text{X}^{2+}(g) + e^- $$

This indicates that a second electron is removed from the already positively charged ion.

The process continues for successive ionisation energies where electrons are removed one after another. Each subsequent ionisation energy is usually higher than the last one due to increased positive charge, making it harder to remove additional electrons.

1.2 State Symbols

When writing the equations, it's essential to include state symbols:

(g) for gas
(l) for liquid
(s) for solid
(aq) for aqueous solution

In our examples above, we are dealing with gaseous atoms (g).

2. Factors Affecting Ionisation Energy

2.1 Nuclear Charge

Nuclear charge is the total charge of the nucleus, due to protons. As the number of protons increases, the nuclear charge increases, which generally increases ionisation energy because the electrons are held more tightly by the nucleus.

2.2 Atomic Radius

The atomic radius is the distance from the nucleus to the outermost electron. A larger atomic radius means the outer electrons are further from the nucleus, decreasing the force of attraction and lowering the ionisation energy. As you move down a group in the periodic table, the atomic radius increases, which usually leads to decreased ionisation energy.

2.3 Electron Shielding/Repulsion

Electrons in inner shells can shield the outer electrons from the full effect of the nucleus’ positive charge. This electron shielding reduces the effective nuclear charge experienced by outer electrons, thus decreasing ionisation energy. Additionally, repulsion between electrons can also affect ionisation energy; as electrons repel each other, it becomes easier to remove an outer electron.

3. Trends in Ionisation Energy

3.1 Across a Period

As we move from left to right across a period in the periodic table, the ionisation energy generally increases. This is due to:

An increase in nuclear charge (more protons) without a significant increase in shielding.
A decrease in atomic radius, making it harder to remove outer electrons.

3.2 Down a Group

As we go down a group, the first ionisation energy generally decreases because:

The atomic radius increases, making the outer electrons further from the nucleus.
Increased electron shielding means that outer electrons feel less attraction from the nucleus.

3.3 Dips in Ionisation Energy at Groups 13 and 16

You might notice that there are dips in ionisation energy at Group 13 (such as aluminium) and Group 16 (like oxygen):

Group 13 Dip: As we move from magnesium (Mg) to aluminium (Al), the first ionisation energy decreases due to the introduction of a new p-electron that experiences greater shielding.
Group 16 Dip: The dip between nitrogen (N) and oxygen (O) is due to the repulsion between paired electrons in the same orbital (the 2p subshell), making it easier to remove an electron from oxygen.

4. Successive Ionisation Energies

4.1 Deduce Electron Configuration

By examining the successive ionisation energies of an element, we can infer how many electrons are in each shell. For example, if there is a large jump in ionisation energy between the 5th and 6th ionisation, it indicates that the first five electrons are removed from the same shell, and the sixth is removed from a significantly closer shell to the nucleus.

4.2 Confirm Group/Position

Using the data from successive ionisation energies, we can confirm an element's group and position in the periodic table. When the energy requirements show a significant increase, this suggests that an electron from a more stable or lower energy level shell is being removed, thus indicating the outer electron count.

Conclusion

In summary, ionisation energies are critical for understanding atomic structure and behavior. They vary based on several factors including nuclear charge, atomic radius, and electron shielding. The trends observed across periods and groups highlight the organization within the periodic table, which helps us deduce important information about the elements. 🌟

Study Notes

Ionisation Energy is the energy required to remove an electron from an atom.
First Ionisation Energy involves removing the first outermost electron.
Successive Ionisation Energies increase due to increased nuclear charge.
Factors affecting ionisation energy: nuclear charge, atomic radius, and electron shielding/repulsion.
Ionisation energy increases across a period and decreases down a group.
Dips in ionisation energy occur at Group 13 and 16 due to electron configurations and repulsion forces.
Successive ionisation data helps deduce electron configurations and confirm the element's group in the periodic table.