When \(\mathrm{I}_{2}\) is passed through \(\mathrm{KCl}, \mathrm{KF}\) and \(\mathrm{KBr}\) : (a) \(\mathrm{Cl}_{2}, \mathrm{~F}_{2}\) and \(\mathrm{Br}\) are evolved (b) \(\mathrm{Cl}_{2}\) and \(\mathrm{Br}_{2}\) are evolved (c) \(\mathrm{Cl}_{2}\) is evolved (d) None of the gas is formed

Understanding the chemical reactivity series is crucial for predicting the outcomes of reactions, particularly displacement reactions. The chemical reactivity series, also known as the activity series, is a list classifying elements, especially metals and halogens, based on their reactivity. Beginning with the most reactive, the elements are arranged in descending order, indicating which elements can displace others from compounds.
For halogens, the series follows the order: F2 (Fluorine) > Cl2 (Chlorine) > Br2 (Bromine) > I2 (Iodine). In displacement reactions, a more reactive halogen can displace less reactive ones from their compounds – for example, chlorine can displace bromine from a bromide. Conversely, iodine, being lower than chlorine and bromine in the reactivity series, is not reactive enough to displace them from their salts, leading to the correct answer in our exercise, where no gases are evolved.

Halogens are a group of non-metal elements located in Group 17 (previously Group VIIA) of the periodic table. This group includes fluorine, chlorine, bromine, iodine, and astatine. Halogens are known for their high reactivity, which diminishes as we move down the group. This trend is due to the increasing atomic size, resulting in a reduced attraction for electrons.
Here are some key properties of halogens:

• Physical State: They exist in different physical states at room temperature – fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid.
• Diatomic Molecules: Halogens naturally occur as diatomic molecules (F2, Cl2, etc).
• Electronegativity: They have high electronegativity, with fluorine being the most electronegative element.
• Displacement Reactions: A more reactive halogen will displace a less reactive halogen from its compounds, consistent with the chemical reactivity series.

These distinctive properties explain the reactivity behavior seen in halogen displacement reactions, including why iodine does not displace other halogens in the given exercise.

The periodic table displays elements in an orderly arrangement, enabling the prediction of chemical behavior based on observed trends. As you move across a period (left to right), electronegativity typically increases, and atoms generally become less metallic. Moving down a group, the elements become more metallic, and the reactivity of metals increases, while the reactivity of non-metals, like halogens, decreases.

Understanding Trends in Halogens

For halogens, reactivity is influenced by:

• Electronegativity: Higher up in the group, halogens are more electronegative and more reactive.
• Atomic Radius: A smaller atomic radius closer to the top of the group makes these elements more reactive.
• Ionization Energy: Lower ionization energy in the upper part of the group correlates with higher reactivity.

These periodic trends explain why iodine, found lower down in the group, is less reactive than other halogens and cannot displace them from their salts, as exemplified in the step-by-step solution.