Problem 61 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following sequence is correct with reference to the oxidation number of iodine? (a) \(\mathrm{HI}<\mathrm{I}_{2}<\mathrm{ICl}<\mathrm{HIO}_{4}\) (b) \(\mathrm{I}_{2}<\mathrm{HI}<\mathrm{ICl}<\mathrm{HIO}_{4}\) (c) \(\mathrm{HIO}_{4}<\mathrm{ICl}<\mathrm{I}_{2}<\mathrm{HI}\) (d) \(\mathrm{I}_{2}<\mathrm{ICl}<\mathrm{HI}<\mathrm{HIO}_{4}\)

Step-by-Step Solution

Answer

The correct sequence is option (a) HI < I_2 < ICl < HIO_4.

Step 1: Identify Oxidation Numbers

First, determine the oxidation number of iodine in each compound: - In HI, iodine has an oxidation number of -1. - In I_2, iodine has an oxidation number of 0 as it is in its elemental form. - In ICl, iodine is assigned an oxidation number of +1 because chlorine is more electronegative. - In HIO_4, the oxidation number of iodine is +7 due to the presence of multiple oxygens pulling electrons away from iodine.

Step 2: Order the Oxidation Numbers

Now that we have the oxidation numbers of iodine in each compound, we can order them from lowest to highest: - HI: -1 - I_2: 0 - ICl: +1 - HIO_4: +7 Thus, the correct order of the sequences by increasing oxidation number is HI < I_2 < ICl < HIO_4.

Step 3: Select the Correct Sequence

With the ordered list, compare against the given options: a) HI < I_2 < ICl < HIO_4 b) I_2 < HI < ICl < HIO_4 c) HIO_4 < ICl < I_2 < HI d) I_2 < ICl < HI < HIO_4 The correct sequence according to our ordering is option (a) HI < I_2 < ICl < HIO_4.

Key Concepts

Iodine Compounds

Iodine, a versatile halogen, forms a variety of compounds by reacting with different elements. Each iodine compound has unique properties and behaves differently in chemical reactions. A common compound of iodine is hydrogen iodide (HI), where iodine bonds with hydrogen. Here, iodine attains an oxidation number of -1, gaining an electron from hydrogen, which is less electronegative.

Another important compound is iodine monochloride (ICl), where iodine reacts with chlorine. Due to chlorine's higher electronegativity, chlorine pulls electrons away from iodine, resulting in an oxidation number of +1 for iodine. Iodine can also form compounds with oxygen, such as periodic acid ( HIO_4), where iodine is in a very high oxidation state +7 due to oxygen's strong ability to attract electrons.

Understanding the different oxidation states in these compounds is crucial for analyzing iodine reactions and their chemical behavior.

Electronegativity

Electronegativity is a vital concept to understand when discussing oxidation numbers, especially in iodine compounds. It refers to an atom's ability to attract and hold onto electrons when forming chemical bonds. The Pauling scale is commonly used to measure electronegativity, with fluorine being the most electronegative element.

In hydrogen iodide (HI), hydrogen has an electronegativity of 2.1, while iodine has 2.5. Consequently, iodine is more electronegative than hydrogen and tends to attract the shared electrons toward itself, gaining an oxidation number of -1.

When iodine bonds with chlorine to form iodine monochloride (ICl), chlorine's higher electronegativity (3.0) compared to iodine leads to electrons being pulled from iodine to chlorine. This shift results in an oxidation number of +1 for iodine in ICl. By understanding electronegativity, it becomes easier to determine why an element in a compound has a specific oxidation number.

Elemental Form

When elements exist in their natural, uncombined form, they are said to be in their elemental state. Iodine, like other elements in the periodic table, has a distinct oxidation number when it is elemental. This number is always zero.

In its elemental form, represented as I_2, iodine atoms are bonded together but have no net charge from gaining or losing electrons. No other elements influence the oxidation state, so the electrons are equally shared between the two iodine atoms. Consequently, the oxidation number is zero.

The elemental form of iodine is unique for its clear zero oxidation state, providing a neutral reference point. Understanding elemental form allows students to distinguish between elemental and compound states, guiding them to correctly assign oxidation numbers in chemical reactions.