Periodic Trends in Ionisation Enthalpy of Elements - Definition, Factors, Trends, FAQs

Edited By Team Careers360 | Updated on Jul 02, 2025 04:38 PM IST

Ionisation Enthalpy Definition: Ionization enthalpy meaning is the amount of energy required by an isolated gaseous atom to lose an electron in its ground state is known as ionization energy, or ionization enthalpy of elements. Cation formation is the outcome of electron loss. The energy required to detach one mole of electrons from one mole of isolated gaseous ions or atoms is known as the initial/first ionization energy of a molecule/atom.

This Story also Contains

Ionization Energy of First, Second, and Subsequent Ones
Factors Affecting the Ionization Energy
Ionization Energy Trend in Periodic Table
General Periodic Trends
Trends in Ionization Enthalpy in a Group
Trends in Ionization Enthalpy Across a Period
Valency
Valency and its Periodic Trends
Key Points

Ionization Energy of First, Second, and Subsequent Ones

The energy required to remove the next electron is the second ionization energy, and so on. At the same time, when compared to the first ionization energy, the second ionization energy is always higher. An alkali metal atom, for example, can be used. Because its loss gives the atom a stable electron shell, removing the first electron is comparatively simple. In addition, removing the second electron creates a new electron shell that is closer to the atomic nucleus and more closely linked.

The following equation can be used to express the first ionization energy of hydrogen:

H (g) → H+ (g) + e^-

ΔH° = -1312.0 kJ/mol

Also read -

Factors Affecting the Ionization Energy

Ionization energy is principally determined by two factors:

Electron repulsion force.
The nucleus and electrons are attracted to each other.

NEET Highest Scoring Chapters & Topics

This ebook serves as a valuable study guide for NEET exams, specifically designed to assist students in light of recent changes and the removal of certain topics from the NEET exam.

Download EBook

When compared to the actual nuclear charge, the effective nuclear charge, which is sensed by the outermost electrons, will be lower. As the inner electrons will obstruct the nuclear charge path, the outermost electrons will be shielded. The shielding effect is the name given to this phenomenon.

Ionization Energy Trend in Periodic Table

General Periodic Trends

It diminishes while moving from the top to the bottom of a group.

It gradually grows from left to right over time.

Trends in Ionization Enthalpy in a Group

As proceeded down in a group, the first ionization enthalpy of the elements drops. Moving down in a group, on the other hand, increases the atomic number and thus the number of shells. The outermost electrons are far distant from the nucleus and can thus be easily removed. The shielding effect as we proceed down a group due to an increased number of shells is the second or another factor that reduces the ionization energy.

NCERT Chemistry Notes:

Trends in Ionization Enthalpy Across a Period

The ionization energy of an element increases when moved from left to right over time. This occurs as the size of atoms decreases over time. Because of the higher nuclear charge, the valence electrons get closer to the atom's nucleus as we proceed from left to right. As the attraction force between electrons and the nucleus grows stronger, it takes more energy to remove an electron from the valence shell.

Two deviations to the pattern appear to be clearly obvious in the first ionization energies figure. Boron's initial ionization energy is lower than that of beryllium, and oxygen's first ionization energy is lower than that of nitrogen. The Hund's rule and the electron configuration of these atoms are to blame for the disparity. Although the boron ionization requires a 2p electron, the first ionization potential electron for beryllium comes from the 2s orbital. The electron comes from the 2p orbital in both oxygen and nitrogen.

Valency

The ability of an atom or a group of chemically connected atoms to create chemical connections with other atoms or groups of atoms is referred to as valence. The number of electrons in an element's outer shell determines its valence (valence). The valence of polyatomic ions is the charge on the particle (such as SO₄^2-).

Related Topics link,

Valency and its Periodic Trends

The periodic table divides elements into groups (columns) based on the number of valence electrons they have, therefore the item's position in the table will give an indication of its valence.

As they tend to give away one electron, both group 1 elements have one valence electron and so have a +1 valence.

The same holds true for group 2 losing two electrons and group 3 losing three electrons.

Group 5 elements, on the other hand, have 5 valence electrons and will tend to take 3 electrons, giving them a valence of -3.

The elements in Group 6 have 6 valence electrons, which tend to take 2 electrons and have a valence of -2.

Components in Group 7 contain 7 valence electrons, while those with a valence of -1 are more likely to take one electron.

Group 8 elements do not react, hence their value is zero

Also read :

Key Points

Let's have a look at the important critical points related to the periodic changes in ionization enthalpy:

Electron volts (eV) and kilojoules per mole (kJ/M) are the two most used unit of ionization energy.
In the gas phase, ionization energy is defined as the minimal energy required to remove an electron from an ion or an atom.

3. Over the course of an element period, the ionization energy tends to rise as it moves from left to right. The atomic radius, on the other hand, diminishes as one moves from left to right over time. As a result, electrons are drawn closer to the nucleus.

4. Ionization energy has a regular pattern on the periodic table.

While moving down a ionization energy periodic table group from top to bottom, the ionization energy tends to decrease. Moving down a group, however, a valence shell can be added. Furthermore, because the outermost electrons come from a positively charged nucleus, they are easier to remove.

Also check-

Frequently Asked Questions (FAQs)

1. What is ionization enthalpy?

Ionization enthalpy is defined as the energy required to remove an electron from a gaseous atom, where the smaller the atom, the higher the energy required to remove the electron because the outermost electron in the ionization energy periodic table has a high attraction force, and the size of the atom decreases from left to right. As a result, ionization enthalpy rises, and as group size reduces, ionization enthalpy falls.

2. What is ionization enthalpy?

Ionization enthalpy is the energy required to remove an electron from a neutral atom in its gaseous state. It's a measure of how strongly an atom holds onto its outermost electron.

3. Mention the element which holds the highest ionization enthalpy in the periodic table?

Helium has the highest ionization enthalpy in the periodic table. It translates to 24.5874 eV (2372.3 kJ/mol). This occurs because helium has two electrons in 1s orbit, which means the electron is very close to the nucleus and unscreened, necessitating a large amount of energy to release it from the nucleus.

4. What are the two factors of ionization energy?

.Ionization energy is principally determined by two factors:

Electron repulsion force.
The nucleus and electrons are attracted to each other.

5. Define the term enthalpy?

Enthalpy is a thermodynamic property that can be defined as the sum of a system's internal energy and the product of the system's pressure and volume. The letter 'H' can be used to represent it.

The following is a mathematical interpretation of the preceding statement:

H = PV + U.

6. What is valency?

7. How does electron configuration affect ionization enthalpy?

Electron configuration plays a crucial role in determining ionization enthalpy. Atoms with stable electron configurations (like full or half-filled subshells) tend to have higher ionization enthalpies because their electrons are more tightly bound.

8. How does the shielding effect influence ionization enthalpy?

The shielding effect occurs when inner electrons partially shield outer electrons from the full nuclear charge. As shielding increases, the effective nuclear charge decreases, resulting in lower ionization enthalpies.

9. How does the effective nuclear charge affect ionization enthalpy?

Effective nuclear charge is the net positive charge experienced by an electron in an atom. As the effective nuclear charge increases, the attraction between the nucleus and the outermost electrons becomes stronger, leading to higher ionization enthalpies.

10. How does the concept of penetration affect ionization enthalpy?

Penetration refers to the ability of electrons to approach the nucleus closely. Electrons in s-orbitals penetrate more than those in p-orbitals, experiencing a stronger nuclear attraction. This contributes to the generally higher ionization enthalpies of s-block elements compared to p-block elements in the same period.

11. How does Hund's rule influence ionization enthalpy?

Hund's rule states that electrons in an orbital will occupy separate sublevels with the same spin before pairing up. This configuration is more stable, resulting in higher ionization enthalpies for elements with half-filled subshells (like nitrogen in the second period).

12. Why is there a larger jump in ionization enthalpy between the second and third periods compared to subsequent periods?

The larger jump in ionization enthalpy between the second and third periods is due to the introduction of d-orbitals in the third period. These d-orbitals provide additional shielding, which results in a less dramatic increase in ionization enthalpy for subsequent periods.

13. How does the ionization enthalpy of metals compare to that of non-metals?

Generally, metals have lower ionization enthalpies compared to non-metals. This is because metals tend to lose electrons easily to form cations, while non-metals tend to gain or share electrons, making it harder to remove electrons from them.

14. What is the relationship between ionization enthalpy and atomic radius?

Ionization enthalpy and atomic radius have an inverse relationship. As atomic radius decreases, ionization enthalpy typically increases because the outermost electrons are closer to the nucleus and experience a stronger attractive force.

15. Why do transition elements have relatively small variations in ionization enthalpies across a period?

Transition elements have small variations in ionization enthalpies across a period because they add electrons to inner d-orbitals. These d-electrons provide increased shielding, which counteracts the increase in nuclear charge, resulting in only slight changes in ionization enthalpy.

16. How does the concept of effective nuclear charge explain the trend in ionization enthalpy across a period?

Effective nuclear charge increases across a period as protons are added to the nucleus while electrons are added to the same energy level. This results in a stronger attraction between the nucleus and the outermost electrons, leading to an increase in ionization enthalpy across the period.

17. What role does quantum mechanics play in understanding ionization enthalpy trends?

Quantum mechanics provides a framework for understanding electron behavior in atoms, including concepts like orbital shapes, electron probability distributions, and energy levels. These principles help explain the observed trends in ionization enthalpy based on electron configurations and nuclear interactions.

18. Why is there a slight decrease in ionization enthalpy from nitrogen to oxygen in the second period?

The slight decrease in ionization enthalpy from nitrogen to oxygen is due to electron pairing. Nitrogen has a half-filled p-subshell (2p3), which is relatively stable. Oxygen's additional electron must pair up in the same orbital, increasing electron-electron repulsion and making it slightly easier to remove an electron.

19. How does the ionization enthalpy of lanthanides compare to that of other elements?

Lanthanides generally have lower ionization enthalpies compared to other elements in the same period. This is due to the poor shielding effect of f-electrons, which results in a lower effective nuclear charge experienced by the outermost electrons.

20. How does the ionization enthalpy of an element relate to its reactivity?

Generally, elements with low ionization enthalpies are more reactive, especially as reducing agents. They can easily lose electrons to form cations, facilitating chemical reactions. Conversely, elements with high ionization enthalpies are often less reactive and may serve as oxidizing agents.

21. Why is there a general increase in ionization enthalpy from left to right across the d-block elements?

The general increase in ionization enthalpy across the d-block elements is due to the increasing nuclear charge. However, this increase is less pronounced than in s and p block elements because the added d-electrons provide some shielding, partially counteracting the increased nuclear attraction.

22. What is the difference between first and second ionization enthalpy?

First ionization enthalpy is the energy required to remove one electron from a neutral atom, while second ionization enthalpy is the energy needed to remove a second electron from the resulting +1 ion. Second ionization enthalpy is always higher than the first.

23. Why is the second ionization enthalpy always higher than the first?

The second ionization enthalpy is always higher because the second electron is being removed from a positively charged ion. The increased nuclear attraction due to the reduced electron shielding makes it harder to remove subsequent electrons.

24. What is the significance of successive ionization enthalpies?

Successive ionization enthalpies provide information about an element's electron configuration and the energy required to remove electrons from different subshells. The large jumps in successive ionization enthalpies often correspond to the removal of core electrons, indicating the transition between different electron shells.

25. Why is the first ionization enthalpy of boron lower than that of beryllium?

The first ionization enthalpy of boron is lower than that of beryllium because boron's outermost electron is in a 2p orbital, which is slightly higher in energy and experiences less nuclear attraction than the 2s orbital of beryllium. This makes it easier to remove an electron from boron.

26. How does ionization enthalpy relate to metallic character?

Ionization enthalpy is inversely related to metallic character. Elements with low ionization enthalpies tend to be more metallic because they can easily lose electrons to form cations, a characteristic property of metals.

27. Why does ionization enthalpy generally increase across a period in the periodic table?

Ionization enthalpy generally increases across a period because the nuclear charge increases while the shielding effect remains relatively constant. This results in a stronger attraction between the nucleus and the outermost electrons, making it harder to remove an electron.

28. How does atomic size affect ionization enthalpy?

Atomic size is inversely related to ionization enthalpy. As atomic size decreases, the outermost electron is closer to the nucleus, experiencing a stronger attractive force. This makes it harder to remove the electron, resulting in higher ionization enthalpy.

29. Why do noble gases have the highest ionization enthalpies in their respective periods?

Noble gases have completely filled outer electron shells, which are very stable configurations. This stability makes it extremely difficult to remove an electron, resulting in very high ionization enthalpies.

30. How does the octet rule relate to ionization enthalpy trends?

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons. Elements that are close to achieving an octet (like halogens) have higher ionization enthalpies, while those that can easily lose electrons to achieve an octet (like alkali metals) have lower ionization enthalpies.

31. What causes the slight dip in ionization enthalpy between Group 2 and Group 13 elements?

The slight dip in ionization enthalpy between Group 2 and Group 13 elements is due to the start of filling the p-subshell. The p-electrons are slightly less tightly bound than s-electrons, making them easier to remove and causing a small decrease in ionization enthalpy.

32. What is the trend in ionization enthalpy down a group in the periodic table?

Ionization enthalpy generally decreases down a group. This is because the atomic size increases, and the outermost electron is farther from the nucleus, experiencing less attraction and requiring less energy to remove.

33. Why do alkali metals have the lowest ionization enthalpies in their respective periods?

Alkali metals have the lowest ionization enthalpies in their periods because they have a single valence electron in their outermost s-orbital. This electron is relatively far from the nucleus and experiences significant shielding from inner electrons, making it easy to remove.

34. What is the relationship between electronegativity and ionization enthalpy?

Electronegativity and ionization enthalpy generally have a positive correlation. Elements with high electronegativity tend to have high ionization enthalpies because they strongly attract electrons, making it difficult to remove them.

35. What is the significance of ionization enthalpy in chemical reactions?

Ionization enthalpy is significant in chemical reactions because it indicates how easily an atom can form a cation. This information is crucial for predicting and understanding the behavior of elements in various chemical processes, such as bond formation and electron transfer reactions.

36. How does the concept of electron affinity relate to ionization enthalpy?

Electron affinity and ionization enthalpy are related but opposite processes. While ionization enthalpy measures the energy required to remove an electron, electron affinity measures the energy released when an atom gains an electron. Generally, elements with high ionization enthalpies also have high electron affinities.

37. Why is there a larger jump in ionization enthalpy when removing an electron from a noble gas configuration?

There is a larger jump in ionization enthalpy when removing an electron from a noble gas configuration because these configurations are exceptionally stable. Breaking this stable octet requires significantly more energy than removing electrons from less stable configurations.

38. How does the concept of electron configuration anomalies (e.g., Cr, Cu) affect ionization enthalpies?

Electron configuration anomalies, such as those in chromium and copper, can affect ionization enthalpies. These elements have slightly higher ionization enthalpies than expected due to the extra stability conferred by their half-filled or completely filled d-subshells.

39. What is the relationship between ionization enthalpy and the energy of emitted photons in atomic emission spectra?

The energy of photons emitted in atomic emission spectra is related to the energy differences between electron energy levels. While not directly equivalent to ionization enthalpy, these spectral lines provide information about electron configurations and energy levels, which are closely related to ionization enthalpies.

40. How does the concept of electronegativity difference between atoms in a molecule relate to ionization enthalpy?

While electronegativity and ionization enthalpy are distinct properties, they are related. Elements with high electronegativity tend to have high ionization enthalpies because they strongly attract electrons. In molecules, large electronegativity differences between atoms can lead to polar bonds, which can influence the energy required to remove electrons from the molecule.

41. How does the ionization enthalpy of metalloids compare to that of metals and non-metals?

Metalloids generally have ionization enthalpies intermediate between those of metals and non-metals. Their ionization enthalpies are typically higher than those of metals but lower than those of non-metals, reflecting their intermediate chemical properties.

42. How does the presence of a full valence shell affect ionization enthalpy?

Elements with a full valence shell (like noble gases) have very high ionization enthalpies. This is because full electron shells are very stable, and it requires a large amount of energy to disrupt this stable configuration by removing an electron.

43. Why do elements in the same group have similar chemical properties despite having different ionization enthalpies?

Elements in the same group have similar chemical properties because they have the same number of valence electrons, which largely determine an element's reactivity. While ionization enthalpies may differ due to factors like atomic size and shielding, the valence electron configuration remains similar within a group.

44. How does the ionization enthalpy of diatomic halogens compare to that of noble gases?

Diatomic halogens have high ionization enthalpies, but they are generally lower than those of noble gases in the same period. This is because halogens are one electron short of a full octet, while noble gases have a complete octet, making it even more difficult to remove an electron from noble gases.

45. What is the significance of the ionization enthalpy of hydrogen in relation to other elements?

Hydrogen's ionization enthalpy is unique because it has only one electron and no core electrons for shielding. Its ionization enthalpy is higher than those of alkali metals but lower than most other elements, reflecting its special position in the periodic table and its ability to behave like both a metal and a non-metal.

46. Why is there a larger difference in ionization enthalpies between Groups 1 and 2 compared to the difference between Groups 2 and 13?

The larger difference in ionization enthalpies between Groups 1 and 2 is due to the significant increase in effective nuclear charge when moving from a single valence electron (Group 1) to two valence electrons (Group 2). The difference between Groups 2 and 13 is smaller because the added electrons in Group 13 go into p-orbitals, which provide some shielding.

47. What is the relationship between ionization enthalpy and electron shell?

Generally, as the principal quantum number (n) increases, the ionization enthalpy decreases. This is because electrons in higher shells are farther from the nucleus and experience less attraction, making them easier to remove.

48. How does spin-orbit coupling affect the ionization enthalpies of heavier elements?

Spin-orbit coupling becomes more significant for heavier elements, particularly those with partially filled d or f subshells. This effect can lead to slight deviations from expected trends in ionization enthalpies, as it influences the energy levels of electrons in these orbitals.

49. How does the concept of diagonal relationships in the periodic table relate to ionization enthalppy?

Diagonal relationships occur between elements diagonally adjacent in the periodic table (e.g., Li and Mg, Be and Al). These elements often have similar ionization enthalpies due to a balance between increasing nuclear charge and increasing atomic size, resulting in comparable effective nuclear charges.

50. How does the relativistic effect influence the ionization enthalpies of heavy elements?

Relativistic effects become significant for very heavy elements, causing the innermost electrons to move at speeds approaching the speed of light. This leads to contraction of s and p orbitals and expansion of d and f orbitals, which can affect shielding and, consequently, ionization enthalpies.

51. How does the lanthanide contraction affect the ionization enthalpies of elements following the lanthanide series?

The lanthanide contraction results in a smaller-than-expected increase in atomic radius for elements following the lanthanide series. This leads to higher-than-expected effective nuclear charges and, consequently, slightly higher ionization enthalpies for these elements.

52. Why is there a general increase in ionization enthalpy from the s-block to the p-block elements within a period?

The general increase in ionization enthalpy from s-block to p-block elements within a period is due to the increasing effective nuclear charge. As more protons are added to the nucleus without adding a new electron shell, the attraction between the nucleus and the outermost electrons becomes stronger.

53. How does the concept of isoelectronic species relate to ionization enthalpy trends?

Isoelectronic species have the same number of electrons but different nuclear charges. Among isoelectronic species, ionization enthalpy increases with increasing nuclear charge because the electrons experience stronger attraction to the nucleus, making them harder to remove.

54. What is the significance of Slater's rules in understanding ionization enthalpy trends?

Slater's rules provide a method for estimating the effective nuclear charge experienced by electrons in multi-electron atoms. These rules help explain and predict trends in ionization enthalpies by accounting for electron shielding effects in different orbitals.

55. What role does ionization enthalpy play in understanding the periodic law and the organization of the periodic table?

Ionization enthalpy is a fundamental periodic property that reflects the electronic structure of atoms. Its periodic trends help validate and explain the organization of elements in the periodic table, supporting the periodic law by