Ionisation enthalpy, also known as ionisation energy, is the energy required to remove an electron from a gaseous atom or ion. Here's how to approach its calculation and understand the concepts:
Understanding Ionisation Enthalpy
Ionisation enthalpy is a crucial concept in chemistry that helps us understand how atoms bond and form compounds. It's essentially the amount of energy it takes to overcome the attraction between an electron and the positively charged nucleus of an atom.
- First Ionisation Enthalpy: The energy needed to remove the first electron from a neutral gaseous atom.
- Second Ionisation Enthalpy: The energy needed to remove the second electron, which is taken from a unipositive ion (an atom that has lost one electron). It is always higher than the first ionisation enthalpy.
- Subsequent Ionisation Enthalpies: Similarly, there is a third, fourth, and so on, each progressively larger, as it becomes harder to remove electrons from an increasingly positively charged ion.
Calculating Ionisation Enthalpy: Key Principles
Although we don't directly "calculate" ionisation enthalpy using a single formula applicable to all elements, the energy can be determined through experiments, and we can understand how it is defined. The provided reference states the ionization energy = + 2.18 × 10–18 J/atom (or + 1312.3 KJ/mole). This is the amount of energy needed to remove the first electron from a Hydrogen atom and shows us the magnitude of energy. However, understanding how this number is found and defined is crucial.
Experimental Measurement
Ionisation enthalpies are typically measured experimentally through methods such as:
- Photoelectron Spectroscopy (PES): This technique uses photons to eject electrons from an atom and measures the kinetic energy of the ejected electrons. From this, we calculate the energy required to remove the electron. The energies can be assigned to individual electrons in their respective subshells.
Factors Influencing Ionisation Enthalpy
The following factors affect ionisation enthalpy values:
- Nuclear Charge: A higher nuclear charge (more protons) increases the attraction between the nucleus and the electrons, resulting in higher ionisation enthalpy.
- Atomic Radius: As atomic radius increases, electrons are further from the nucleus and less tightly held, leading to lower ionisation enthalpy.
- Shielding Effect: Core electrons shield outer electrons from the full nuclear charge, reducing the attraction and thus lowering the ionisation enthalpy.
- Subshell Occupancy: A completely or half-filled subshell usually results in an increased stability and consequently higher ionisation enthalpies.
Trend in Ionisation Enthalpy
- Across a Period (Left to Right): Ionisation enthalpy generally increases across a period due to increasing nuclear charge and decreasing atomic radius.
- Down a Group: Ionisation enthalpy generally decreases down a group due to an increase in atomic radius and shielding.
Practical Example
Let's consider the case of Magnesium (Mg). It has two valence electrons and the following would happen:
- First Ionisation Enthalpy: Energy is needed to remove the first valence electron, forming Mg+.
- Second Ionisation Enthalpy: Energy required to remove the second valence electron from Mg+, forming Mg2+. The second ionisation enthalpy is always greater than the first ionisation enthalpy for any atom, because we are removing an electron from a positively charged ion.
The provided reference only provides the specific ionisation energy for the removal of an electron from a hydrogen atom (+ 2.18 × 10–18 J/atom or + 1312.3 KJ/mole). Specific ionisation energies for different elements will be different, which have been experimentally determined.
Summary
In conclusion, ionisation enthalpy is not calculated with a single formula but is determined experimentally. The energies are affected by factors such as nuclear charge, atomic radius, and electron shielding. Understanding trends across the periodic table and the distinction between the first and subsequent ionisation enthalpies gives a comprehensive grasp of this essential chemical property.
