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Chapter 4: Problem 17

Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms. $$ 2 \mathrm{KrF}_{2}+2 \mathrm{H}_{2} \mathrm{O} \rightarrow 2 \mathrm{Kr}+4 \mathrm{HF}+\mathrm{O}_{2} $$

Short Answer

Expert verified
Kr is reduced; O in H鈧侽 is oxidized.

Step by step solution

01

Assign Oxidation Numbers to Elements in Reactants

Assign oxidation numbers to all elements in the reactants. \( \mathrm{KrF}_{2} \): \( \mathrm{Kr} \) is \(+4\), \( \mathrm{F} \) is \(-1\), \( \mathrm{H}_{2} \mathrm{O} \): \( \mathrm{H} \) is \(+1\), \( \mathrm{O} \) is \(-2\).
02

Assign Oxidation Numbers to Elements in Products

Assign oxidation numbers to all elements in the products. \( \mathrm{Kr} \) remains \(0\), \( \mathrm{HF} \): \( \mathrm{H} \) is \(+1\), \( \mathrm{F} \) is \(-1\), \( \mathrm{O}_{2} \): \( \mathrm{O} \) is \(0\).
03

Identify Changes in Oxidation Numbers

Compare the oxidation numbers from reactants to products. \( \mathrm{Kr} \) changes from \(+4\) to \(0\); \( \mathrm{O} \) changes from \(-2\) in \( \mathrm{H}_{2} \mathrm{O} \) to \(0\) in \( \mathrm{O}_{2} \).
04

Determine Oxidized and Reduced Elements

Identify which elements gain or lose electrons. \( \mathrm{Kr} \) is reduced because its oxidation number decreases from \(+4\) to \(0\). \( \mathrm{O} \) is oxidized because its oxidation number increases from \(-2\) to \(0\).

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Key Concepts

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

Oxidation Numbers
In the study of chemistry, oxidation numbers are a helpful tool for tracking the transfer of electrons in oxidation-reduction (redox) reactions. An oxidation number assigns a charge to an atom assuming electronegativity fully involves electrons, whether bonding or not.
To assign these numbers, remember these rules:
  • The oxidation number of an uncombined element is always 0.
  • For ions, the oxidation number equals the ion's charge.
  • In compounds, hydrogen typically has an oxidation number of +1, while oxygen often has -2.
  • Fluorine, being highly electronegative, always carries an oxidation number of -1 in its compounds.
  • In neutral substances, the sum of oxidation numbers of all atoms must equal 0. For polyatomic ions, it equals the ion鈥檚 charge.
In the given redox equation, applying these rules allows us to assign oxidation numbers to all involved elements, facilitating further analysis of oxidation and reduction processes.
Electron Transfer
Electron transfer is a fundamental concept in redox reactions, essential for understanding the movement of electrons between reactants and products. In a redox reaction, transfer of electrons from one substance to another defines the process.
Consider what really happens when an atom changes its oxidation state:
  • Oxidation occurs when an atom or molecule loses electrons. This process increases its oxidation number.
  • Reduction occurs when an atom or molecule gains electrons, decreasing its oxidation number.
For the given equation, we observe this clearly:
  • Krypton (Kr) gains electrons, moving its oxidation state from +4 in KrF extsubscript{2} to 0 as Kr. This represents the reduction of Kr.
  • Meanwhile, oxygen (O) loses electrons as it transitions from -2 in H extsubscript{2}O to 0 in O extsubscript{2}, signifying its oxidation.
Electron transfer drives the conversion, ensuring a balanced redox process.
Redox Equations
Redox equations offer a clear representation of redox reactions, where one species undergoes oxidation and another undergoes reduction. These reactions play a critical role in various chemical and biological processes.
In balancing redox equations:
  • Identify the species that get oxidized and reduced by determining changes in oxidation numbers.
  • Ensure that the electrons lost in the oxidation process are equal to those gained during reduction.
  • The equation often includes both half-reactions showing oxidation and reduction separately before combining them into the overall balanced equation.
The given equation, 2KrF extsubscript{2} + 2H extsubscript{2}O 鈫� 2Kr + 4HF + O extsubscript{2}, is a classic example of a balanced redox equation:
- **Krypton (Kr)** undergoes reduction, as seen by the decrease in oxidation number. - **Oxygen (O)** is oxidized, with its oxidation number increasing.
This equation illustrates the essential balance needed for accurate depiction of the electron flow in chemical reactions.

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Most popular questions from this chapter

From the statement "propane reacts with oxygen to produce carbon dioxide and water," identify the reactants and the products.

Use the periodic table or the activity series to predict if each single- replacement reaction will occur and, if so, write a balanced chemical equation. a) \(\mathrm{Pt}+\mathrm{H}_{3} \mathrm{PO}_{4} \rightarrow\) ? b) \(\mathrm{Li}+\mathrm{H}_{2} \mathrm{O} \rightarrow\) ? (Hint: treat \(\mathrm{H}_{2} \mathrm{O}\) as if it were composed of \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-}\) ions. \()\)

Assuming that each double-replacement reaction occurs, predict the products and write each balanced chemical equation. a) \(\mathrm{Sn}(\mathrm{OH})_{2}+\mathrm{FeBr}_{3} \rightarrow\) ? b) \(\mathrm{CsNO}_{3}+\mathrm{KCl} \rightarrow\) ?

Assign oxidation numbers to each atom in each substance. a) \(\mathrm{C}_{6} \mathrm{H}_{6}\) b) \(\mathrm{B}(\mathrm{OH}) 3\) c) \(\mathrm{Li}_{2} \mathrm{~S}\) d) \(\mathrm{Au}\)

Write a chemical equation that represents ( \(\left.\mathrm{NH}_{4}\right)_{3} \mathrm{PO}_{4}(\mathrm{~s})\) dissociating in water.
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