Question

Which compound among the following has lowest oxidation number of chlorine? (a) \(\mathrm{HClO}_{4}\) (b) \(\mathrm{HClO}_{3}\) (c) \(\mathrm{HCl}\) (d) \(\mathrm{HOCl}\)

Short answer

\(\mathrm{HCl}\) has the lowest oxidation number of chlorine, which is -1.

Step-by-step solution

Understand the concept of oxidation numbers

The oxidation number or oxidation state is the charge an atom would have if all bonds to atoms of different elements were fully ionic. Each compound will have a set oxidation number for chlorine that we can calculate based on the known rules.

Calculate oxidation numbers in each compound

Use the rules for oxidation numbers to find the oxidation state of chlorine in each compound. Oxygen usually has an oxidation number of -2, and hydrogen is +1. In all compounds, the sum of oxidation numbers should equal the overall charge of the compound (which is 0 for all the given molecules).

Calculate the oxidation number of chlorine in \(\mathrm{HClO}_{4}\)

Let the oxidation number of Cl in \(\mathrm{HClO}_{4}\) be x. Considering that the compound is neutral and the oxidation numbers of H is +1 and O is -2, the equation becomes: +1 + x + 4(-2) = 0. Solving for x gives x = +7.

Calculate the oxidation number of chlorine in \(\mathrm{HClO}_{3}\)

Similarly, for \(\mathrm{HClO}_{3}\) with chlorine's oxidation number as x, the equation is: +1 + x + 3(-2) = 0. This gives x = +5.

Calculate the oxidation number of chlorine in \(\mathrm{HCl}\)

For \(\mathrm{HCl}\), chlorine's oxidation number (x) is obtained by +1 + x = 0 because there are no oxygen atoms. This gives x = -1.

Calculate the oxidation number of chlorine in \(\mathrm{HOCl}\)

For \(\mathrm{HOCl}\), assuming chlorine's oxidation number as x, the equation is: +1 + x + (-2) = 0. This implies x = +1.

Identify the compound with the lowest oxidation number of chlorine

Compare the calculated oxidation numbers of chlorine in each compound, and select the one with the lowest value.

Key concepts

Oxidation State

The concept of oxidation state, also known as oxidation number, plays a critical role in understanding redox reactions and balancing redox equations. An oxidation state is a theoretical charge that an atom would have if all its bonds to atoms of different elements were completely ionic.

To determine the oxidation state, we must apply a set of rules. For example, the oxidation state of a pure element is always zero. For monoatomic ions, the oxidation state is equal to the ion's charge. Hydrogen generally has an oxidation state of +1, while oxygen usually has an oxidation state of -2, except in peroxides or when bonded to fluorine.

In a compound, the sum of the oxidation states must equal the overall charge. For neutral molecules, this sum is zero. Knowing how to assign these values is essential because the changes in oxidation states indicate which atoms are oxidized and which are reduced in redox reactions. Accurately assigning oxidation states also aids in balancing redox equations, ensuring that charge and mass are conserved.

Chemical Compounds

Chemical compounds consist of atoms of different elements bonded together. The properties of compounds, such as reactivity and stability, are largely determined by the nature of these bonds and the oxidation states of the atoms involved. In the context of oxidation states, each element in a compound has its own specific value that can be calculated based on established rules.

When looking at compounds with elements that can have multiple oxidation states (like chlorine in the given problem), it's important to consider the whole compound's structure. For instance, in oxyacids (compounds that contain hydrogen, oxygen, and another element), hydrogen is assigned a +1 oxidation state, and oxygen is -2. The oxidation state of the central element, such as chlorine in our case, varies and is what determines the compound's reactivity and strength as an acid.

Understanding how to determine the oxidation state of an element in various compounds is crucial not only for balancing equations but also for predicting the behavior of substances during a chemical reaction.

Balancing Redox Equations

Balancing redox equations is about ensuring that the same number of electrons are lost and gained during a reaction, maintaining the principle of conservation of charge. To balance redox reactions, it helps to separate the two half-reactions – the reduction half and the oxidation half. Each half-reaction is balanced separately for mass and charge.

The steps typically involve: assigning oxidation states to all atoms to determine which are oxidized and reduced; writing the oxidation half-reaction and reduction half-reaction separately; balancing the atoms in each half-reaction that are neither hydrogen nor oxygen; balancing oxygen atoms by adding water molecules; balancing hydrogen atoms by adding hydrogen ions (in acidic solution) or hydroxide ions (in basic solution); and balancing the charge by adding electrons.

Once both half-reactions are balanced, they can be combined by multiplying each half-reaction by a suitable integer to have an equal number of electrons in both. This ensures that the electrons lost in the oxidation reaction are gained in the reduction reaction. The final step is to combine the half-reactions, cancel anything that is the same on both sides, and ensure the total number of atoms and charges are equal. This rigorous process is vital in understanding the stoichiometry of the reaction as well as the energy changes that occur.