Ionization Enthalpy: Definition, Principle, Factors & its Trends across the Periodic Table

The energy required to remove an electron from the gaseous atom is known as Ionization enthalpy. The energy required to remove the 1st electron from an isolated atom is called first ionization enthalpy. Thus, it is a quantitative measure of how much a gaseous atom is capable of losing its electrons. In this article, we will learn about the factors affecting ionization enthalpy along with its trends in the modern periodic table of Chemistry Class 11 syllabus.

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What is Ionization Enthalpy?

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In chemistry, ionization enthalpy is the minimum amount of energy required for any gaseous atom (X) to lose electrons from its ground state. The tendency of any element to lose electrons can be assessed by quantitative analysis of its ionization enthalpy. In other words, the term Ionization enthalpy is defined as “the amount of energy necessary to extract one electron from a gaseous atom”. Ionization enthalpies are always positive since it takes energy to remove electrons from an atom.

Measurement of ionization enthalpy is based on a principle, “In any electrical discharge tube, when a fast moving electron produced by the flow of current collides with an isolated gaseous atom it falls out”. In chemistry class 11, ionization enthalpy is represented by I.E and units of I.E are: Electron volts (per atom), Kilo-calorie per mole, or Kilo-joules per mole. 

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Factors Affecting Ionization Enthalpy

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Ionization enthalpy depends mainly on three factors, namely, penetration effect, shielding effect and electronic configuration. These three factors are mentioned in detail below:

1. Penetration Effect

The closeness or affinity of electrons to nucleus is called its penetration power. The shell and sub-shells of electron are determined by the penetration power of electron. In general, the penetration effect is shown as: 

2s>2p>3s>3p>4s>3d

2. Electronic Configuration

Removing electrons from more stable orbitals changes them to less stable orbitals. For example, half-filled and completely filled are the two most stable configurations of an element. Every element in chemistry becomes stable by acquiring one of these two electronic configurations. Furthermore, there is a direct relationship between stability and ionization enthalpy. Thus, a more stable element will have a higher ionization enthalpy and vice versa.

3. Shielding Effect

The repelling influence of the inner electrons shields the outermost electron from the nucleus' attraction. As shielding improves, the attraction of the positive nucleus to the negative electron decreases, requiring less energy to remove the electron. As a result the ionization enthalpy decreases. Due to this, the effective nuclear charge of the outer electrons becomes comparatively less. In chemistry, the formula for effective nuclear charge is given by:

Z_eff= Z-S

Where, Zeff= Effective Nuclear Charge,

Z= Orignal Nuclear Charge,

S= Screening Constant

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Trends in the Periodic Table

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From the above explanation of ionization enthalpy, it should now be clear that the element with higher ionization enthalpy is stronger than its counterpart. Now, let us understand the trends in ionization enthalpies across both groups and periods in the modern periodic table.

Across Groups

As we proceed down the group from one element to the next, the ionisation enthalpy continues to decrease. In other words, with increase in the number of shells down the group, the ionization enthalpy decreases. Due to this the distance between the outermost electron and the nucleus increases and the effective nuclear charge also decreases. In groups (moving from top to bottom), as the number of shells increases, the shielding effect also increases. This also results in increased ionization enthalpy across the groups.

Across periods

Ionization enthalpy increases across the periods. As we move from left to right in the periodic table, we will see that the atomic radii are gradually decreasing. Thus, with decrease in atomic radius, the force of attraction between the nucleus and outer electrons increases resulting in increase in ionization enthalpy. Since chemistry is full of exceptions, there is an exception in this case too. If we look at the ionization enthalpy from boron to beryllium, ideally the I.E of boron should be higher than beryllium, but in real scenario this is not true. Let us now know the reason behind this!

There are two main reasons: (a) the fully-filled configuration of beryllium, and (b) penetration effects.

Due to the presence of 2s and 2p orbitals in boron, it becomes easier to remove electrons from the 2p subshell compared to beryllium, which has only 2s orbitals or 2s subshell. Furthermore, the penetrating power of the 2s orbital is greater than that of the 2p orbital. Thus, the I.E of beryllium becomes higher than that of boron.

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Things to Remember

• Ionization enthalpy is the minimum amount of energy required for an electron to escape from its orbital. 
• The factors affecting ionization enthalpy are (i) penetration effect, (ii) shielding effect, and (iii) electronic configuration. 
• In chemistry, it is denoted by I.E and the units are Electron volts (per atom), Kilo-calorie per mole, or Kilo-joules per mole. 
• There is an indirect relationship between number of shells and ionization enthalpy across the group (moving from top to bottom).
• The value of I.E increases across the periods (moving from left to right). 

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Previous Year Questions

1. Which ion has the largest radius?… [DUET 2011]
2. It is believed that atoms combine with each other such that the outermost shell acquires a stable configuration of… [DUET 2009]
3. Which one of the following ions has the highest value of ionic radius ?… [DUET 2009]
4. Which of the following has the highest ionisation energy?… [DUET 2008]
5. An element with configuration… [DUET 2008]
6. The value of electronegativity of /atoms… [DUET 2009]
7. Chloride ion and potassium ion are isoelectronic. 
8. Beryllium shows diagonal relationship with aluminium. Which of the following similarity is incorrect ?
9. The smallest among the following ion is:
10. Total number of rare earth elements is… [ BHU UET 2008]
11. The first ionisation potential (in eV) of Be and B, respectively are… [NEET 1998]
12. Identify the wrong statement in the following… [NEET 2012]
13. Among the elements Ca,Mg,P and Cl, the order of increasing atomic radii is… [NEET 2010]
14. Among the following which one has the highest cation to anion size ratio?… [NEET 2010]
15. Which one of the following ions will be smallest in size?… [NEET 1996]
16. The decreasing order of basic character of K₂O,BaO,CaO and MgO is… [JIPMER 2015]
17. Which one of the following arrangements represents the correct order of electron gain enthalpy… [NEET 2005]
18. One of the characteristic properties of non-metals is that they… [NEET 1993]
19. Which electronic configuration of an element has abnormally high difference between second and third ionisation energy?… [NEET 1993]