Periodic Table & Periodicity

To determine the decreasing order of ionic size for the ions Mg²⁺, O^2-, Na⁺, and F^-, we need to consider several factors such as nuclear charge, electrons’ configurations, and their positions in the periodic table.

Firstly, when ions have the same electronic configuration (which we can think of as isoelectronic), the nuclear charge and its effective pull on the electrons will determine their sizes. A higher positive charge (more protons in the nucleus) results in a smaller ionic radius because the protons pull the electron cloud more tightly towards the nucleus.

The ions Mg²⁺, O^2-, Na⁺, and F^- are all approximately isoelectronic, having the same number of electrons as the noble gas Ne (10 electrons). Thus, the order of their ionic sizes will mainly depend on their nuclear charge:

• O^2- (8 protons)
• F^- (9 protons)
• Na⁺ (11 protons)
• Mg²⁺ (12 protons)

Higher proton count correlates with greater effective nuclear charge, leading to a more strongly pulled-in electron cloud, and thus a smaller ionic radius.

Therefore, the decreasing order of ionic size from largest to smallest is:

$ \text{O}^{2-} > \text{F}^{-} > \text{Na}^{+} > \text{Mg}^{2+} $

O^2-, having the fewest protons among these ions, exhibits the largest ionic size due to a lesser effective nuclear pull on its electron cloud, compared to the other ions with more protons. Conversely, Mg²⁺ is the smallest as it has the highest number of protons exerting the strongest pull on the same number of electrons.

When evaluating the acidic properties of compounds, we must consider the nature of the oxides formed by different elements. The acidity of an oxide depends on the position of the respective element in the periodic table, particularly their electronegativity and the period to which they belong.

Oxides can be broadly categorized into four types based on the nature of the elements forming them:

• Metallic oxides (usually basic)
• Non-metallic oxides (usually acidic)
• Amphoteric oxides (show both acidic and basic behavior)
• Neutral oxides (do not show acidic or basic behavior)

Let's analyze the given compounds:

• ZnO: Zinc oxide is amphoteric. It reacts with both acids and bases, but tends more towards basic than acidic properties.
• Na₂O₂: Sodium peroxide is generally a basic oxide, but it can also react with water to form hydroxide ions, indicating it is less acidic than ZnO.
• MgO: Magnesium oxide is a basic oxide. It reacts with acids to form salts and water but does not react with bases.
• P₂O₅: Diphosphorus pentoxide is strongly acidic. It reacts readily with water to form phosphoric acid.

Based on the typical behavior of their chemical classification, here is the order of increasing acidic property:

1. Na₂O₂ (mostly basic)
2. MgO (basic)
3. ZnO (amphoteric, weakly acidic)
4. P₂O₅ (strongly acidic)

The increasing acidic behavior aligns with moving from basic, through amphoteric, to strongly acidic oxides.

On the Mulliken scale, the average of ionization potential and electron affinity is known as electronegativity. In essence, the Mulliken electronegativity scale is defined by the formula:

$ \chi_{Mulliken} = \frac{I + E}{2} $

Where:

• $ \chi_{Mulliken} $ is the Mulliken electronegativity,
• $ I $ is the ionization potential (the energy required to remove an electron from an atom in the gaseous state),
• $ E $ is the electron affinity (the energy released when an electron is added to a neutral atom in the gaseous state).

This formula essentially provides a measure of how strongly an atom in a molecule attracts electrons towards itself. The concept of electronegativity is crucial in understanding chemical bonding, particularly the polarity of bonds and molecular structures. By using the average of ionization potential and electron affinity, Mulliken's approach provides a quantitative way of assessing an atom's ability to attract shared electrons in a chemical bond.

The energy released when an electron is added to a neutral gaseous atom is called the electron affinity of the atom. Electron affinity is a measure of the energy change that occurs when an electron is added to a neutral atom in the gas phase to form a negative ion. This property is important in understanding the chemical behavior of atoms. It can either release energy, indicated by a negative value, meaning the process is exothermic, or absorb energy, indicated by a positive value, indicating an endothermic process.