How does ionization enthalpy vary in a group?

Ionization enthalpy generally decreases down a group due to the increase in atomic size and screening effect.

Why do alkali metals have low ionization enthalpy?

Alkali metals have the largest atomic radii in their respective periods and a single electron in their outermost shell, making it easy to remove.

What is the element with atomic number 87?

The element with atomic number 87 is Francium (Fr), which is an alkali metal in Period 7.

What is the element with atomic number 19?

The element with atomic number 19 is Potassium (K), an alkali metal in Period 4.

Between 19 and 87, which has lower ionization energy?

Francium (87) has a much larger atomic radius than Potassium (19), so its ionization enthalpy is lower.

What are the elements for atomic numbers 32 and 35?

Atomic number 32 is Germanium (Ge) and 35 is Bromine (Br).

Does electronic configuration affect ionization enthalpy?

Yes, atoms with stable half-filled or fully-filled orbitals often show higher ionization enthalpies than expected.

How does effective nuclear charge affect this property?

An increase in effective nuclear charge ($Z_{eff}$) increases the attraction between the nucleus and electrons, raising ionization enthalpy.

Which group in the periodic table has the lowest ionization enthalpy?

Group 1 (Alkali Metals) has the lowest ionization enthalpy in any given period.

The atomic number of the element from the following with lowest 1st ionisation enthalpy is

Q: What is first ionization enthalpy?

It is the energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.

The atomic number of the element from the following with lowest 1st ionisation enthalpy is : | JEE Main Chemistry

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QMCQ

The atomic number of the element from the following with lowest $1^{st}$ ionisation enthalpy is :

A) 32
B) 35
C) 19
D) 87

✅ Correct Answer

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Solution Steps

Identify the Elements

First, identify the elements corresponding to the given atomic numbers ($Z$):

* $Z = 32$ is Germanium (Ge), Group 14, Period 4.

* $Z = 35$ is Bromine (Br), Group 17, Period 4.

* $Z = 19$ is Potassium (K), Group 1, Period 4.

* $Z = 87$ is Francium (Fr), Group 1, Period 7.

Analyze Group Trends

In any given period, the Group 1 elements (Alkali Metals) have the lowest $1^{st}$ ionisation enthalpy because they have the largest atomic size and only one electron in their valence shell ($ns^1$).

Compare Period 4 Elements

Comparing K (19), Ge (32), and Br (35) which are all in Period 4: Ionisation enthalpy increases from left to right across a period. Therefore, K (19) has a lower ionisation enthalpy than Ge and Br.

Compare Group 1 Elements

Now compare Potassium (19) and Francium (87), both of which are Group 1 elements. Down a group, ionisation enthalpy decreases because:

1. Atomic radius increases significantly.

2. The outermost electron is further from the nucleus and more effectively shielded by inner electrons.

Determine the Lowest Value

Since Francium (87) is much further down Group 1 (Period 7) than Potassium (Period 4), its $1^{st}$ ionisation enthalpy is significantly lower than that of Potassium.

Final Conclusion

Among the given options, the alkali metal in the highest period (Period 7) will have the lowest ionisation energy.

Final Answer: 87 (Option D)

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Theory

1. Definition of Ionisation Enthalpy

Ionisation enthalpy ($\Delta_i H$) represents the minimum amount of energy required to remove the most loosely bound electron from an isolated neutral gaseous atom ($X$) to convert it into a gaseous cation ($X^+$). It is always an endothermic process, meaning the value is positive. The first ionisation enthalpy refers to the removal of the first electron. Factors affecting this value include nuclear charge, atomic size, screening effect of inner electrons, and the nature of the electronic configuration (such as half-filled or fully-filled subshells which provide extra stability). Generally, the smaller the atom and the higher the effective nuclear charge, the higher the energy required to remove an electron.

2. Periodic Trends in Groups

Down a group in the periodic table, the first ionisation enthalpy consistently decreases. This trend occurs because each successive element has an additional principal energy level (shell), which increases the distance between the valence electrons and the nucleus. Furthermore, the number of inner-shell electrons increases, leading to a stronger shielding or screening effect. This shielding offsets the increase in actual nuclear charge, resulting in a lower effective nuclear charge ($Z_{eff}$) acting on the valence electrons. Consequently, the attraction between the nucleus and the outermost electron weakens, making it easier (requiring less energy) to remove the electron as we move from top to bottom.

3. Periodic Trends in Periods

Across a period from left to right, the first ionisation enthalpy generally increases. This is primarily due to the increase in nuclear charge (number of protons) while electrons are added to the same valence shell. This results in a higher effective nuclear charge ($Z_{eff}$), which pulls the electrons closer to the nucleus, decreasing atomic size and increasing the electrostatic attraction. As a result, more energy is required to overcome this attraction to remove an electron. Alkali metals, located at the far left (Group 1), have the lowest ionisation enthalpies in their periods, while noble gases at the far right (Group 18) have the highest due to their stable octet configurations.

4. Properties of Alkali Metals (Group 1)

Alkali metals, including Lithium, Sodium, Potassium, Rubidium, Cesium, and Francium, are characterized by having a single valence electron in an $s$-orbital ($ns^1$). They possess the largest atomic radii in their respective periods. Because they only need to lose one electron to achieve a stable noble gas configuration, and because that electron is relatively far from the nucleus, Group 1 elements exhibit the lowest first ionisation enthalpies of all elements. Within Group 1, Francium (atomic number 87) is the largest atom in Period 7, meaning its valence electron is the most weakly held, giving it one of the lowest ionisation enthalpies in the entire periodic table.

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FAQs

Why is Francium’s ionisation energy so low?

Francium is a very large atom with its valence electron in the 7s orbital, far from the nucleus and heavily shielded by 86 inner electrons, making the electron easy to remove.

Does $Z=87$ have the lowest IE in the whole periodic table?

Cesium (Cs, $Z=55$) actually has the lowest measured stable ionisation enthalpy. Francium’s value is slightly higher than Cesium’s due to relativistic effects on its 7s electrons, but among the given options, 87 is the lowest.

Why do noble gases have high ionisation enthalpy?

Noble gases have a completely filled valence shell (stable octet) and high effective nuclear charge, making them very resistant to losing an electron.

What is the unit of ionisation enthalpy?

It is typically measured in kJ/mol or electron-volts (eV) per atom.

What is the screening effect?

The screening effect is the reduction in nuclear attraction on valence electrons caused by the repulsion from inner-shell electrons.

How does half-filled configuration affect IE?

Half-filled subshells (like $p^3$ or $d^5$) provide extra stability, leading to higher-than-expected ionisation enthalpies (e.g., Nitrogen vs. Oxygen).

Is the second ionisation enthalpy higher than the first?

Yes, the second ionisation enthalpy is always higher than the first because it’s harder to remove an electron from a positively charged ion.

What element is atomic number 35?

Atomic number 35 is Bromine, a halogen in Period 4. It has a relatively high ionisation enthalpy.

Why does atomic size increase down a group?

Because with each step down a group, a new principal energy level (electron shell) is added to the atom.

What is effective nuclear charge?

It is the net positive charge experienced by an electron in a multi-electron atom, calculated as $Z_{eff} = Z – S$, where $S$ is the shielding constant.

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