Question

Give the oxidation number of bromine in each of the following: (a) \(\mathrm{KBr}\) (b) \(\mathrm{BrF}_{3}\) (c) \(\mathrm{HBrO}_{3}\) (d) \(\mathrm{CBr}_{4}\)

Short answer

(a) -1 (b) +3 (c) +5 (d) -1

Step-by-step solution

Determine oxidation number in \(\textbf{KBr}\)

Potassium (K) has an oxidation number of +1. Since the compound is neutral, the sum of oxidation numbers for K and Br must be zero. Therefore, the oxidation number of bromine (Br) in \(\text{KBr}\) is -1.

Determine oxidation number in \(\textbf{BrF}_{3}\)

Fluorine (F) always has an oxidation number of -1. With 3 fluorine atoms, the total oxidation number contributed by fluorine atoms is \(-1 \times 3 = -3\). Since the molecule is neutral, bromine (Br) must balance this out. Therefore, the oxidation number of bromine in \(\text{BrF}_{3}\) is +3.

Determine oxidation number in \(\textbf{HBrO}_{3}\)

Hydrogen (H) has an oxidation number of +1, and oxygen (O) has an oxidation number of -2. With 3 oxygen atoms, the total oxidation number from oxygen is \(-2 \times 3 = -6\). Including hydrogen, the total so far is \(+1 - 6 = -5\). For the molecule to be neutral, bromine (Br) must have an oxidation number of +5 in \(\text{HBrO}_{3}\).

Determine oxidation number in \(\textbf{CBr}_{4}\)

Assume carbon's (C) oxidation number is +4 in many organic compounds. Given the compound is neutral, the total oxidation number of four bromine atoms should balance carbon's oxidation number of +4. Therefore, each bromine atom must have an oxidation number of -1 in \(\text{CBr}_{4}\).

Key concepts

Oxidation States

Oxidation states, also known as oxidation numbers, are crucial in understanding chemical reactions. They provide insight into the electron transfer between atoms in a compound. Each element in a compound is assigned an oxidation number, which is a hypothetical charge that an atom would have if all bonds were purely ionic. This helps in identifying elements that are gaining or losing electrons during chemical reactions.

There are simple rules to follow:

• The oxidation state of an element in its elemental form is always zero. For example, the oxidation number of O₂ or Br₂ is 0.
• For ions, the oxidation state equals the charge of the ion.
• In compounds, specific elements often have consistent oxidation states. For example, fluorine is always -1, and oxygen is usually -2, except in peroxides.
• The sum of oxidation states in a neutral compound is zero, while in a polyatomic ion, it equals the ion's charge.

Following these rules will help you deduce the oxidation state of any element in a compound reliably.

Chemical Compounds Analysis

Analyzing chemical compounds is key to determining how elements interact. This involves determining the oxidation states of the atoms in the compound. Let's break down the process with the given examples:

• KBr: Potassium (K) typically has an oxidation number of +1. The compound is neutral, so the oxidation state of bromine (Br) must be -1 to balance it out.
• BrF₃: Each fluorine atom has an oxidation number of -1, and there are three fluorine atoms contributing a total of -3. Bromine must be +3 to balance the total oxidation number to zero in this neutral molecule.
• HBrO₃: Hydrogen (H) is +1 and oxygen (O) is typically -2. With three oxygens, their total contribution is -6. Adding hydrogen, we have +1 - 6 = -5. Thus, to balance the molecule, bromine must be +5.
• CBr₄: Carbon (C) often has an oxidation state of +4 in its compounds. Since the compound is neutral, each bromine must have an oxidation number of -1 to balance the +4 of carbon, giving us -4 from bromines.

This kind of analysis allows chemists to understand the properties and reactivities of different compounds.

Redox Reactions

Redox (reduction-oxidation) reactions are processes where the oxidation states of atoms change. These reactions involve the transfer of electrons between molecules.

In a redox reaction:

• Oxidation: Loss of electrons by a molecule, atom, or ion. This usually results in an increase in its oxidation number.
• Reduction: Gain of electrons by a molecule, atom, or ion. This leads to a decrease in its oxidation number.

It is essential to identify which substances are oxidized and reduced in a reaction. For example, in the reaction between hydrogen and fluorine to form HF, hydrogen is oxidized (loses electrons, oxidation state goes from 0 to +1), and fluorine is reduced (gains electrons, oxidation state goes from 0 to -1).

Understanding redox reactions is fundamental in fields such as energy production, metabolism in biology, and industrial chemical processes. They help explain phenomena ranging from rusting iron to the functioning of batteries.