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Chapter 8: Problem 80

Rank the following oxoanions in order of decreasing oxidation number of chlorine: (a) \(\mathrm{ClO}^{-} ;\) (b) \(\mathrm{ClO}_{2}^{-}\) (c) \(\mathrm{ClO}_{3}^{-} ;\) (d) \(\mathrm{ClO}_{4}^{-}\)

Short Answer

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Question: Rank the given oxoanions in order of decreasing oxidation number of chlorine: (a) \(\mathrm{ClO}^{-}\), (b) \(\mathrm{ClO}_{2}^{-}\), (c) \(\mathrm{ClO}_{3}^{-}\), and (d) \(\mathrm{ClO}_{4}^{-}\). Answer: The order of oxoanions in decreasing oxidation state of chlorine is: (d) \(\mathrm{ClO}_{4}^{-}\) > (c) \(\mathrm{ClO}_{3}^{-}\) > (b) \(\mathrm{ClO}_{2}^{-}\) > (a) \(\mathrm{ClO}^{-}\).

Step by step solution

01

Identify the Oxoanions

The given oxoanions are: a) \(\mathrm{ClO}^{-}\) - named as hypochlorite. b) \(\mathrm{ClO}_{2}^{-}\) - named as chlorite. c) \(\mathrm{ClO}_{3}^{-}\) - named as chlorate. d) \(\mathrm{ClO}_{4}^{-}\) - named as perchlorate.
02

Find the Oxidation State of Chlorine in Each Oxoanion

As the oxidation state of oxygen is always -2 (except in peroxides and other special cases), we can create equations to find the oxidation state of chlorine in each oxoanion: a) \(x + (-2) = -1\) for hypochlorite (ClO-), solve for x. b) \(x + 2(-2) = -1\) for chlorite (ClO2-), solve for x. c) \(x + 3(-2) = -1\) for chlorate (ClO3-), solve for x. d) \(x + 4(-2) = -1\) for perchlorate (ClO4-), solve for x.
03

Solve the Equations for the Oxidation State of Chlorine

Now, we will solve the equations for x (oxidation state of chlorine): a) \(x + (-2) = -1 \Rightarrow x = +1\) for hypochlorite (ClO-) b) \(x + 2(-2) = -1 \Rightarrow x = +3\) for chlorite (ClO2-) c) \(x + 3(-2) = -1 \Rightarrow x = +5\) for chlorate (ClO3-) d) \(x + 4(-2) = -1 \Rightarrow x = +7\) for perchlorate (ClO4-)
04

Rank the Oxoanions in Order of Decreasing Oxidation State of Chlorine

Now that we have the oxidation state of chlorine in each oxoanion, we can rank them in order of decreasing oxidation state: 1) Perchlorate (ClO4-) with +7 oxidation state. 2) Chlorate (ClO3-) with +5 oxidation state. 3) Chlorite (ClO2-) with +3 oxidation state. 4) Hypochlorite (ClO-) with +1 oxidation state. So, the order of oxoanions in decreasing oxidation state of chlorine is: (d) \(\mathrm{ClO}_{4}^{-}\) > (c) \(\mathrm{ClO}_{3}^{-}\) > (b) \(\mathrm{ClO}_{2}^{-}\) > (a) \(\mathrm{ClO}^{-}\).

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Oxidation State
The oxidation state, also known as oxidation number, is a concept that helps to keep track of electrons in a chemical compound. In oxoanions, this state indicates the effective charge of an atom. It's important to understand that the oxidation state can change during chemical reactions.
For chlorine in our given oxoanions, the oxidation states can be determined using some simple calculations. Oxygen typically has an oxidation state of -2. Knowing the overall charge of the anion allows us to set up an equation and solve for the oxidation state of chlorine:
  • For hypochlorite (\( \text{ClO}^- \)), if we assume chlorine has an oxidation state \( x \), the equation becomes: \( x - 2 = -1 \). Solving this gives \( x = +1 \).
  • For chlorite (\( \text{ClO}_2^- \)), the equation is \( x + 2(-2) = -1 \), leading to \( x = +3 \).
  • For chlorate (\( \text{ClO}_3^- \)), we have \( x + 3(-2) = -1 \), resulting in \( x = +5 \).
  • For perchlorate (\( \text{ClO}_4^- \)), the equation is \( x + 4(-2) = -1 \), which gives \( x = +7 \).
Oxidation states provide valuable insight into the nature of the compound and can predict reactivity and stability.
Chlorine
Chlorine is an element from group 17 of the periodic table, known as the halogens. It is highly electronegative, meaning it has a strong tendency to attract electrons. This characteristic is important in determining oxidation states in chlorine oxoanions.
Chlorine can exist in multiple oxidation states, ranging from -1 to +7. The different oxidation states are due to chlorine's ability to bond with oxygen in various ratios, as seen in oxoanions like chlorate and perchlorate.
  • In \( \text{ClO}^- \) (hypochlorite), chlorine has an oxidation state of +1.
  • In \( \text{ClO}_2^- \) (chlorite), it has an oxidation state of +3.
  • In \( \text{ClO}_3^- \) (chlorate), the oxidation state is +5.
  • In \( \text{ClO}_4^- \) (perchlorate), chlorine achieves its highest common oxidation state of +7.
These various states give chlorine-containing compounds a wide range of chemical properties and applications.
Chemical Nomenclature
Chemical nomenclature is the formal method of naming chemical compounds. In the case of oxoanions, which are negatively charged ions that contain oxygen, there's a specific system in place to name them. This system is crucial in reducing ambiguity in chemistry.
The names of chlorine-containing oxoanions reflect the number of oxygen atoms and the oxidation state of chlorine:
  • "Hypo-" is used when there is the least amount of oxygen, such as in \( \text{ClO}^- \) (hypochlorite).
  • "Chlorite" refers to \( \text{ClO}_2^- \), indicating a higher number of oxygen atoms compared to hypochlorite.
  • "Chlorate" (\( \text{ClO}_3^- \)) suggests an even higher oxygen content.
  • "Per-" is prefixed to "chlorate" in \( \text{ClO}_4^- \) to denote the maximum number of oxygen atoms, hence naming it perchlorate.
This nomenclature provides clarity and consistency, helping chemists and students identify the compound's composition, especially for those involving variable oxidation states.

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Most popular questions from this chapter

Iron is oxidized in a number of chemical weathering processes. How many moles of \(\mathrm{O}_{2}\) are consumed when one mole of magnetite (Fe \(_{3} \mathrm{O}_{4}\) ) is converted into hematite \(\left(\mathrm{Fe}_{2} \mathrm{O}_{3}\right) ?\)

What is the difference between a saturated solution and a supersaturated solution?

Some people who prefer natural foods make their own apple cider vinegar. They start with freshly squeezed apple juice that contains about \(6 \%\) natural sugars. These sugars, which all have nearly the same empirical formula, \(\mathrm{CH}_{2} \mathrm{O},\) are fermented with yeast in a chemical reaction that produces equal numbers of moles of ethanol \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\right)\) and carbon dioxide. The product of fermentation, called hard cider, undergoes an acid fermentation step in which ethanol and dissolved oxygen gas react together to form acetic acid (CH \(_{3} \mathrm{COOH}\) ) and water. This acetic acid is the principal solute in vinegar. a. Write a balanced chemical equation for the fermentation of natural sugars to ethanol and carbon dioxide. You may use in the equation the empirical formula given in the preceding paragraph. b. Write a balanced chemical equation for the acid fermentation of ethanol to acetic acid. c. What are the oxidation states of carbon in the reactants and products of the two fermentation reactions? d. If a sample of apple juice contains \(1.00 \times 10^{2} \mathrm{g}\) of natural sugar, what is the maximum quantity of acetic acid that could be produced by the two fermentation reactions?

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