Question

The best reducing agent of the following substances is (a) \(\mathrm{H}_{2} \mathrm{S} ;\) (b) \(\mathrm{O}_{3} ;\) (c) \(\mathrm{H}_{2} \mathrm{SO}_{4} ;\) (d) \(\mathrm{NaF}\); (e) \(\mathrm{H}_{2} \mathrm{O}\).

Short answer

The best reducing agent among the given substances is \(\mathrm{H}_{2}\mathrm{S}\).

Step-by-step solution

Identify the oxidations states of the substances

The substances given are \(\mathrm{H}_{2}\mathrm{S}\), \(\mathrm{O}_{3}\), \(\mathrm{H}_{2} \mathrm{SO}_{4}\), \(\mathrm{NaF}\) and \(\mathrm{H}_{2} \mathrm{O}\). Their oxidation states are -2, 0, +6, -1, and -2 respectively. Remember that the oxidation state indicates the degree of oxidation of an atom in a chemical compound. The most electronegative element in a compound is assigned the negative oxidation number.

Analyze the standard reduction potentials of the substances

Knowing the oxidation states, the standard reduction potentials of the substances can be checked from a standard reduction potentials table. The substance with the most positive standard reduction potential is the strongest oxidizing agent, while the one with the most negative standard reduction potential is the strongest reducing agent.

Identify the strongest reducing agent

By comparing the standard reduction potentials, the substance with the highest ability to lose electrons (i.e., the one with the most negative or least positive standard reduction potential) will be the best reducing agent. Among the given substances, \( \mathrm{H}_{2}\mathrm{S} \) is the best reducing agent as it has the most negative standard reduction potential compared to the others.

Key concepts

Understanding Oxidation States

Oxidation states, also known as oxidation numbers, help us understand how electrons are distributed in a compound. They represent the hypothetical charge an atom would have if all bonds were ionic. For example, in

• a compound like \( \mathrm{H}_{2}\mathrm{S}\), the oxidation state of hydrogen is typically +1, while sulfur is -2.
• In \( \mathrm{O}_{3} \) or ozone, the oxygen atoms have an oxidation state of 0 as it is a molecule made up of the same element.
• In complex molecules like \( \mathrm{H}_{2} \mathrm{SO}_{4} \) (sulfuric acid), hydrogen is +1, sulfur is usually considered to be +6, and oxygen is -2.

Oxidation states help us determine how likely a substance is to donate or accept electrons during a chemical reaction. The substance with the lowest or negative oxidation state is more likely to act as a reducing agent.

Deciphering Standard Reduction Potential

Standard reduction potential is a measure of the tendency of a chemical species to acquire electrons and be reduced. These potentials are typically measured in volts.

• A substance with a high positive standard reduction potential readily gains electrons, making it a strong oxidizing agent.
• Conversely, a substance with a low or negative standard reduction potential tends to lose electrons easily, acting as a good reducing agent.

A standard table of reduction potentials can reveal potential differences. Comparing these values tells us which substances are better at donating electrons (stronger reducing agents) or accepting electrons (stronger oxidizing agents). For instance, \( \mathrm{H}_{2}\mathrm{S} \) with a low negative standard reduction potential is a strong candidate for a reducing agent. Understanding this concept allows chemists to predict reaction outcomes and manipulate chemical behavior efficiently.

Role of Electronegativity

Electronegativity is an atom's ability to attract and hold onto electrons in a chemical bond. The more electronegative an element, the stronger its pull on electrons. This concept plays a crucial role in determining oxidation states, as the more electronegative element often takes on a negative oxidation state, while the less electronegative one takes on a positive state.

• In \( \mathrm{NaF} \), fluorine is more electronegative than sodium, so it exhibits a -1 oxidation state.
• Electronegativity also helps explain why, in compounds like \( \mathrm{H}_{2}\mathrm{O} \), oxygen carries a -2 oxidation state because it's more electronegative than hydrogen.

Electronegativity not only influences oxidation states but also affects how molecules interact, the strength of bonds, and the nature of reactions. Recognizing the interplay between electronegativity and oxidation states is vital for predicting and understanding chemical reactivity.