Question

The first ionisation enthalpies of the alkaline earth metals are higher than that of alkali metals but second ionisation enthalpies are smaller, why? (a) In alkali metals, second ionisation enthalpy involves removal of electron from noble gas electronic configuration while in alkaline earth metals, second electron is removed from \(n s^{2}\) configuration. (b) Alkaline earth metals have very high melting point as compared to alkali metals. (c) Electrons in s-orbital are more closely packed in alkaline earth metals than alkali metals. (d) Due to smaller size alkaline earth metals do not form divalent ions very easily.

Short answer

The answer is (a) because the second ionisation enthalpy of alkali metals involves removing an electron from a noble gas configuration which requires more energy, while in alkaline earth metals, the second electron is removed from a less stable ns^2 configuration, thus requiring less energy.

Step-by-step solution

Understanding Ionisation Enthalpy

Ionisation enthalpy refers to the energy required to remove an electron from a gaseous atom or ion. The first ionisation enthalpy is the energy needed to remove the first electron, while the second ionisation enthalpy is the energy required to remove the second electron.

Comprehending the Statement Given in (a)

For alkali metals, the second electron is being removed from a noble gas configuration, which is highly stable. Thus, the second ionisation enthalpy is very high. In contrast, for alkaline earth metals, the second electron is removed from an ns^2 configuration, which is less stable than a noble gas configuration, resulting in a lower second ionisation enthalpy.

Considering Melting Point from Statement (b)

Melting point is a bulk property and doesn't directly affect ionisation enthalpy of individual atoms. Thus, although alkaline earth metals have a higher melting point than alkali metals, it is not related to the trends in first and second ionisation enthalpies.

Evaluating Electron Configuration from Statement (c)

The fact that electrons are more closely packed in s-orbital of alkaline earth metals than in alkali metals does influence ionisation enthalpy. In alkaline earth metals, there are two electrons in the s-orbital which slightly shields each other making it easier to remove one, while in alkali metals only one electron occupies the s-orbital, making it more strongly held and increasing the first ionisation enthalpy.

Analyzing Atomic Size and Valency from Statement (d)

The smaller size of alkaline earth metals compared to alkali metals means that the nucleus holds onto the (relatively more) electrons more tightly, leading to a higher first ionisation enthalpy. However, the formation of divalent ions is not directly related to the comparison between the first and second ionisation enthalpies.

Deducing the Correct Explanation

From the given statements, (a) explains the observed trend most effectively. The removal of the second electron from a noble gas configuration (in the case of alkali metals) requires much more energy than removing an electron from an ns^2 configuration (in the case of alkaline earth metals) due to the extremely stable nature of the noble gas electronic configuration.

Key concepts

Alkaline Earth Metals

Alkaline earth metals comprise the second group of the periodic table including beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). Characterized by their ns² valence shell configuration, these metals possess two electrons in their outermost s orbital. An intriguing property of alkaline earth metals is their higher first ionisation enthalpy relative to alkali metals. This phenomenon is attributable to the full s orbital, which, although allows slight shielding effects between the two electrons, still retains a strong attraction to the nucleus due to the effective nuclear charge.

• The full s orbital contributes to a higher first ionisation energy because the two electrons provide some mutual electron-electron repulsion, slightly reducing the ionisation energy required to remove one electron.
• However, their second ionisation enthalpy is smaller because the removal of the first electron leaves an electron in the s orbital that is easier to remove than if the configuration were a noble gas.

It is important to note that the atomic size of alkaline earth metals is generally smaller than alkali metals, leading to a stronger nuclear pull on the valence electrons, hence, requiring more energy to remove the first electron.

Alkali Metals

Belonging to the first group of the periodic table, alkali metals include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). These elements are defined by their single electron in the outermost s orbital, with an electronic configuration of ns¹. This lone electron experiences a relatively low effective nuclear charge, which facilitates its removal and consequently results in the lower first ionisation enthalpy compared to alkaline earth metals.

• Upon the removal of this single valence electron, alkali metals attain the noble gas configuration, making the second ionisation enthalpy exceptionally high.
• The larger atomic radius of alkali metals also factors into the reduced ionisation enthalpy, as the valence electron is further away from the nucleus and less tightly bound.

Their chemical reactivity and tendency to form +1 cations are characteristic properties that arise from this ease of electron loss.

Noble Gas Electronic Configuration

The noble gases, located at the far right of the periodic table, are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). Their outer electron shells are fully occupied, which confers chemical inertness and exceptional stability, a state often referred to as the noble gas electronic configuration. This configuration is represented as ns²np⁶ for noble gases starting from neon.

• The fully filled outer shell renders these gases highly unreactive, as they have no tendency to lose or gain electrons.
• For elements other than noble gases, attaining this configuration through ionisation comes at a substantial energetic cost, which explains the high second ionisation energies in alkali metals.
• Additionally, the octet rule in chemistry is based on the stability of this noble gas configuration, with many elements reacting to achieve a similar electron arrangement.

Understanding the noble gas configuration is essential when evaluating the ionisation enthalpies of other elements that desire to reach this state.

ns^2 Configuration

The ns² configuration refers to the presence of two electrons in the outermost s orbital of an atom. This configuration is particularly noteworthy for the alkaline earth metals, which have such an electronic arrangement in their ground state.

• The ns² configuration of alkaline earth metals is less stable than the noble gas configuration, making the second ionisation enthalpy lower than that of the alkali metals, which have a noble gas configuration after the first ionisation.
• The two electrons in the ns² configuration exert repulsive forces on each other, which means that after one is removed, it is relatively easier to remove the second electron.

In summary, the ns² configuration plays a pivotal role in determining the reactivity and ionisation properties of a group of elements in the periodic table, particularly influencing their first and second ionisation enthalpies.