The oxidizer and reducer with oxidized and reduced products are to be stated. The oxidation and reduction half-reaction equations are to be stated. The balanced redox equation is to be stated. Concept introduction: The oxidizer is the species whose oxidation state falls during the course of reaction and reducer is the species whose oxidation number increases. Oxidized product is the oxidation product of the reducer and reduced product is the reduction product of the oxidizer.

Question

Want to see more full solutions like this?

Answer 1

Question

Chapter 19, Problem 49E

Interpretation Introduction

Interpretation:

The oxidizer and reducer with oxidized and reduced products are to be stated. The oxidation and reduction half-reaction equations are to be stated. The balanced redox equation is to be stated.

Concept introduction:

The oxidizer is the species whose oxidation state falls during the course of reaction and reducer is the species whose oxidation number increases. Oxidized product is the oxidation product of the reducer and reduced product is the reduction product of the oxidizer.

Expert Solution & Answer

Answer to Problem 49E

The oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation is shown below.

$2NH4+→N2+8H++6e−$

The reduction half-reaction equation is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

The balanced redox equation is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Explanation of Solution

The given redox reaction equation to be balanced is shown below.

$Cr2O72−+NH4+→Cr2O3+N2$

The oxidation state of the central metal atom is calculated by knowing the standard oxidation states of few elements.

The oxidation state of chromium in $Cr2O72−$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O7n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O72n 7(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O72n+7(−2)=−2$

Calculate the value of n by simplifying the equation as shown below.

$2n+7(−2)=−22n+(−14)=−22n=−2+142n=+12$

Divide by two on both sides and simplify as shown below.

$2n2=+122n=+6$

The oxidation state of chromium in $Cr2O72−$ is $+6$ .

The oxidation state of chromium in $Cr2O3$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O3n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O32n 3(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O32n+3(−2)=0$

Calculate the value of n by simplifying the equation as shown below.

$2n+3(−2)=02n+(−6)=02n=0+62n=+6$

Divide by two on both sides and simplify as shown below.

$2n2=+62n=+3$

The oxidation state of chromium in $Cr2O3$ is $+3$ .

The oxidation state of the nitrogen in $NH4+$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$N H4n +1$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$N H4n 4(+1)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$N H4n+4(+1)=+1$

Calculate the value of n by simplifying the equation as shown below.

$n+4(+1)=+1n+4=+1n=+1−4n=−3$

The oxidation state of nitrogen is $−3$ in $NH4+$ .

The oxidation number of nitrogen in $N2$ is zero as this is the elemental form of the nitrogen.

The chromium in $Cr2O72−$ $(Cr=+6)$ is getting reduced because the oxidation state of chromium in the product $Cr2O3$ $(Cr=+3)$ is less than in the reactant. Similarly, the nitrogen is getting oxidized from the reactant $NH4+$ $(N=−3)$ to the product $N2$ $(N=0)$ .

Therefore, the oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation for the above equation is shown below.

$NH4+→N2$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The nitrogen is getting oxidized and the number of atoms of that is not balanced on both sides of the equation. Multiply ammonium ion by two to balance the nitrogen.

$2NH4+→N2$

Step-2: Balance elements other than oxygen and hydrogen if any.

$2NH4+→N2$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

$2NH4+→N2$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding the eight $H+$ on the right-hand side of the equation.

$2NH4+→N2+8H+$

Step-5: Balance the charge by adding electrons to the appropriate side.

Six electrons are added to the right-hand side in order to balance the charge.

$2NH4+→N2+8H++6e−$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$2NH4+→N2+8H++6e−$ …(1)

The reduction half-reaction for the above reaction is shown below.

$Cr2O72−→Cr2O3$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The chromium is getting reduced and its number of atoms is balanced on both sides.

$Cr2O72−→Cr2O3$

Step-2: Balance elements other than oxygen and hydrogen if any.

$Cr2O72−→Cr2O3$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

The number of oxygen atoms is balanced by adding four water molecules on the right-hand side of the equation.

$Cr2O72−→Cr2O3+4H2O$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding eight $H+$ ions on the left-hand side of the equation.

$Cr2O72−+8H+→Cr2O3+4H2O$

Step-5: Balance the charge by adding electrons to the appropriate side.

The charge is balanced by adding six electrons on the left-hand side of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$ …(2)

The balanced redox equation is obtained by adding equation (1) and (2) in such a way that electrons are canceled out.

Add equation (1) and (2) and cancel out the common things on both sides of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O+2NH4+→N2+8H++6e−$

The balance redox equation after adding these equations is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Conclusion

The oxidizer and reducer with oxidized and reduced products, oxidation and reduction half-reaction equations, and balanced redox equation are rightfully stated above.

Want to see more full solutions like this?

Subscribe now to access step-by-step solutions to millions of textbook problems written by subject matter experts!

Students have asked these similar questions

21.50 Determine the combinations of haloalkane(s) and alkoxide(s) that could be used to synthesize the following ethers through Williamson ether synthesis. (a) (c) (d) (e) (f) H₂CO

1. Arrange the following in order of increasing bond energy (lowest bond energy first, highest bond energy last). Provide your rationale. C=C, C-F, C=C, C-N, C-C List the bond order for each example.

What is the major enolate formed when treated with LDA? And why that one?

Answer 2

Question

Chapter 19, Problem 49E

Interpretation Introduction

Interpretation:

The oxidizer and reducer with oxidized and reduced products are to be stated. The oxidation and reduction half-reaction equations are to be stated. The balanced redox equation is to be stated.

Concept introduction:

The oxidizer is the species whose oxidation state falls during the course of reaction and reducer is the species whose oxidation number increases. Oxidized product is the oxidation product of the reducer and reduced product is the reduction product of the oxidizer.

Expert Solution & Answer

Answer to Problem 49E

The oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation is shown below.

$2NH4+→N2+8H++6e−$

The reduction half-reaction equation is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

The balanced redox equation is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Explanation of Solution

The given redox reaction equation to be balanced is shown below.

$Cr2O72−+NH4+→Cr2O3+N2$

The oxidation state of the central metal atom is calculated by knowing the standard oxidation states of few elements.

The oxidation state of chromium in $Cr2O72−$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O7n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O72n 7(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O72n+7(−2)=−2$

Calculate the value of n by simplifying the equation as shown below.

$2n+7(−2)=−22n+(−14)=−22n=−2+142n=+12$

Divide by two on both sides and simplify as shown below.

$2n2=+122n=+6$

The oxidation state of chromium in $Cr2O72−$ is $+6$ .

The oxidation state of chromium in $Cr2O3$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O3n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O32n 3(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O32n+3(−2)=0$

Calculate the value of n by simplifying the equation as shown below.

$2n+3(−2)=02n+(−6)=02n=0+62n=+6$

Divide by two on both sides and simplify as shown below.

$2n2=+62n=+3$

The oxidation state of chromium in $Cr2O3$ is $+3$ .

The oxidation state of the nitrogen in $NH4+$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$N H4n +1$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$N H4n 4(+1)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$N H4n+4(+1)=+1$

Calculate the value of n by simplifying the equation as shown below.

$n+4(+1)=+1n+4=+1n=+1−4n=−3$

The oxidation state of nitrogen is $−3$ in $NH4+$ .

The oxidation number of nitrogen in $N2$ is zero as this is the elemental form of the nitrogen.

The chromium in $Cr2O72−$ $(Cr=+6)$ is getting reduced because the oxidation state of chromium in the product $Cr2O3$ $(Cr=+3)$ is less than in the reactant. Similarly, the nitrogen is getting oxidized from the reactant $NH4+$ $(N=−3)$ to the product $N2$ $(N=0)$ .

Therefore, the oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation for the above equation is shown below.

$NH4+→N2$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The nitrogen is getting oxidized and the number of atoms of that is not balanced on both sides of the equation. Multiply ammonium ion by two to balance the nitrogen.

$2NH4+→N2$

Step-2: Balance elements other than oxygen and hydrogen if any.

$2NH4+→N2$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

$2NH4+→N2$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding the eight $H+$ on the right-hand side of the equation.

$2NH4+→N2+8H+$

Step-5: Balance the charge by adding electrons to the appropriate side.

Six electrons are added to the right-hand side in order to balance the charge.

$2NH4+→N2+8H++6e−$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$2NH4+→N2+8H++6e−$ …(1)

The reduction half-reaction for the above reaction is shown below.

$Cr2O72−→Cr2O3$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The chromium is getting reduced and its number of atoms is balanced on both sides.

$Cr2O72−→Cr2O3$

Step-2: Balance elements other than oxygen and hydrogen if any.

$Cr2O72−→Cr2O3$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

The number of oxygen atoms is balanced by adding four water molecules on the right-hand side of the equation.

$Cr2O72−→Cr2O3+4H2O$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding eight $H+$ ions on the left-hand side of the equation.

$Cr2O72−+8H+→Cr2O3+4H2O$

Step-5: Balance the charge by adding electrons to the appropriate side.

The charge is balanced by adding six electrons on the left-hand side of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$ …(2)

The balanced redox equation is obtained by adding equation (1) and (2) in such a way that electrons are canceled out.

Add equation (1) and (2) and cancel out the common things on both sides of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O+2NH4+→N2+8H++6e−$

The balance redox equation after adding these equations is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Conclusion

The oxidizer and reducer with oxidized and reduced products, oxidation and reduction half-reaction equations, and balanced redox equation are rightfully stated above.

Want to see more full solutions like this?

Subscribe now to access step-by-step solutions to millions of textbook problems written by subject matter experts!

Answer 3

Question

Chapter 19, Problem 49E

Interpretation Introduction

Interpretation:

The oxidizer and reducer with oxidized and reduced products are to be stated. The oxidation and reduction half-reaction equations are to be stated. The balanced redox equation is to be stated.

Concept introduction:

The oxidizer is the species whose oxidation state falls during the course of reaction and reducer is the species whose oxidation number increases. Oxidized product is the oxidation product of the reducer and reduced product is the reduction product of the oxidizer.

Expert Solution & Answer

Answer to Problem 49E

The oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation is shown below.

$2NH4+→N2+8H++6e−$

The reduction half-reaction equation is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

The balanced redox equation is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Explanation of Solution

The given redox reaction equation to be balanced is shown below.

$Cr2O72−+NH4+→Cr2O3+N2$

The oxidation state of the central metal atom is calculated by knowing the standard oxidation states of few elements.

The oxidation state of chromium in $Cr2O72−$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O7n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O72n 7(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O72n+7(−2)=−2$

Calculate the value of n by simplifying the equation as shown below.

$2n+7(−2)=−22n+(−14)=−22n=−2+142n=+12$

Divide by two on both sides and simplify as shown below.

$2n2=+122n=+6$

The oxidation state of chromium in $Cr2O72−$ is $+6$ .

The oxidation state of chromium in $Cr2O3$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O3n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O32n 3(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O32n+3(−2)=0$

Calculate the value of n by simplifying the equation as shown below.

$2n+3(−2)=02n+(−6)=02n=0+62n=+6$

Divide by two on both sides and simplify as shown below.

$2n2=+62n=+3$

The oxidation state of chromium in $Cr2O3$ is $+3$ .

The oxidation state of the nitrogen in $NH4+$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$N H4n +1$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$N H4n 4(+1)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$N H4n+4(+1)=+1$

Calculate the value of n by simplifying the equation as shown below.

$n+4(+1)=+1n+4=+1n=+1−4n=−3$

The oxidation state of nitrogen is $−3$ in $NH4+$ .

The oxidation number of nitrogen in $N2$ is zero as this is the elemental form of the nitrogen.

The chromium in $Cr2O72−$ $(Cr=+6)$ is getting reduced because the oxidation state of chromium in the product $Cr2O3$ $(Cr=+3)$ is less than in the reactant. Similarly, the nitrogen is getting oxidized from the reactant $NH4+$ $(N=−3)$ to the product $N2$ $(N=0)$ .

Therefore, the oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation for the above equation is shown below.

$NH4+→N2$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The nitrogen is getting oxidized and the number of atoms of that is not balanced on both sides of the equation. Multiply ammonium ion by two to balance the nitrogen.

$2NH4+→N2$

Step-2: Balance elements other than oxygen and hydrogen if any.

$2NH4+→N2$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

$2NH4+→N2$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding the eight $H+$ on the right-hand side of the equation.

$2NH4+→N2+8H+$

Step-5: Balance the charge by adding electrons to the appropriate side.

Six electrons are added to the right-hand side in order to balance the charge.

$2NH4+→N2+8H++6e−$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$2NH4+→N2+8H++6e−$ …(1)

The reduction half-reaction for the above reaction is shown below.

$Cr2O72−→Cr2O3$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The chromium is getting reduced and its number of atoms is balanced on both sides.

$Cr2O72−→Cr2O3$

Step-2: Balance elements other than oxygen and hydrogen if any.

$Cr2O72−→Cr2O3$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

The number of oxygen atoms is balanced by adding four water molecules on the right-hand side of the equation.

$Cr2O72−→Cr2O3+4H2O$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding eight $H+$ ions on the left-hand side of the equation.

$Cr2O72−+8H+→Cr2O3+4H2O$

Step-5: Balance the charge by adding electrons to the appropriate side.

The charge is balanced by adding six electrons on the left-hand side of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$ …(2)

The balanced redox equation is obtained by adding equation (1) and (2) in such a way that electrons are canceled out.

Add equation (1) and (2) and cancel out the common things on both sides of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O+2NH4+→N2+8H++6e−$

The balance redox equation after adding these equations is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Conclusion

The oxidizer and reducer with oxidized and reduced products, oxidation and reduction half-reaction equations, and balanced redox equation are rightfully stated above.

Want to see more full solutions like this?

Subscribe now to access step-by-step solutions to millions of textbook problems written by subject matter experts!

Answer 4

Question

Chapter 19, Problem 49E

Interpretation Introduction

Interpretation:

The oxidizer and reducer with oxidized and reduced products are to be stated. The oxidation and reduction half-reaction equations are to be stated. The balanced redox equation is to be stated.

Concept introduction:

The oxidizer is the species whose oxidation state falls during the course of reaction and reducer is the species whose oxidation number increases. Oxidized product is the oxidation product of the reducer and reduced product is the reduction product of the oxidizer.

Expert Solution & Answer

Answer to Problem 49E

The oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation is shown below.

$2NH4+→N2+8H++6e−$

The reduction half-reaction equation is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

The balanced redox equation is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Explanation of Solution

The given redox reaction equation to be balanced is shown below.

$Cr2O72−+NH4+→Cr2O3+N2$

The oxidation state of the central metal atom is calculated by knowing the standard oxidation states of few elements.

The oxidation state of chromium in $Cr2O72−$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O7n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O72n 7(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O72n+7(−2)=−2$

Calculate the value of n by simplifying the equation as shown below.

$2n+7(−2)=−22n+(−14)=−22n=−2+142n=+12$

Divide by two on both sides and simplify as shown below.

$2n2=+122n=+6$

The oxidation state of chromium in $Cr2O72−$ is $+6$ .

The oxidation state of chromium in $Cr2O3$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O3n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O32n 3(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O32n+3(−2)=0$

Calculate the value of n by simplifying the equation as shown below.

$2n+3(−2)=02n+(−6)=02n=0+62n=+6$

Divide by two on both sides and simplify as shown below.

$2n2=+62n=+3$

The oxidation state of chromium in $Cr2O3$ is $+3$ .

The oxidation state of the nitrogen in $NH4+$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$N H4n +1$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$N H4n 4(+1)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$N H4n+4(+1)=+1$

Calculate the value of n by simplifying the equation as shown below.

$n+4(+1)=+1n+4=+1n=+1−4n=−3$

The oxidation state of nitrogen is $−3$ in $NH4+$ .

The oxidation number of nitrogen in $N2$ is zero as this is the elemental form of the nitrogen.

The chromium in $Cr2O72−$ $(Cr=+6)$ is getting reduced because the oxidation state of chromium in the product $Cr2O3$ $(Cr=+3)$ is less than in the reactant. Similarly, the nitrogen is getting oxidized from the reactant $NH4+$ $(N=−3)$ to the product $N2$ $(N=0)$ .

Therefore, the oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation for the above equation is shown below.

$NH4+→N2$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The nitrogen is getting oxidized and the number of atoms of that is not balanced on both sides of the equation. Multiply ammonium ion by two to balance the nitrogen.

$2NH4+→N2$

Step-2: Balance elements other than oxygen and hydrogen if any.

$2NH4+→N2$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

$2NH4+→N2$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding the eight $H+$ on the right-hand side of the equation.

$2NH4+→N2+8H+$

Step-5: Balance the charge by adding electrons to the appropriate side.

Six electrons are added to the right-hand side in order to balance the charge.

$2NH4+→N2+8H++6e−$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$2NH4+→N2+8H++6e−$ …(1)

The reduction half-reaction for the above reaction is shown below.

$Cr2O72−→Cr2O3$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The chromium is getting reduced and its number of atoms is balanced on both sides.

$Cr2O72−→Cr2O3$

Step-2: Balance elements other than oxygen and hydrogen if any.

$Cr2O72−→Cr2O3$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

The number of oxygen atoms is balanced by adding four water molecules on the right-hand side of the equation.

$Cr2O72−→Cr2O3+4H2O$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding eight $H+$ ions on the left-hand side of the equation.

$Cr2O72−+8H+→Cr2O3+4H2O$

Step-5: Balance the charge by adding electrons to the appropriate side.

The charge is balanced by adding six electrons on the left-hand side of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$ …(2)

The balanced redox equation is obtained by adding equation (1) and (2) in such a way that electrons are canceled out.

Add equation (1) and (2) and cancel out the common things on both sides of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O+2NH4+→N2+8H++6e−$

The balance redox equation after adding these equations is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Conclusion

The oxidizer and reducer with oxidized and reduced products, oxidation and reduction half-reaction equations, and balanced redox equation are rightfully stated above.

Want to see more full solutions like this?

Subscribe now to access step-by-step solutions to millions of textbook problems written by subject matter experts!

Answer 5

Chapter 19, Problem 49E

Interpretation Introduction

Interpretation:

The oxidizer and reducer with oxidized and reduced products are to be stated. The oxidation and reduction half-reaction equations are to be stated. The balanced redox equation is to be stated.

Concept introduction:

The oxidizer is the species whose oxidation state falls during the course of reaction and reducer is the species whose oxidation number increases. Oxidized product is the oxidation product of the reducer and reduced product is the reduction product of the oxidizer.

Expert Solution & Answer

Answer to Problem 49E

The oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation is shown below.

$2NH4+→N2+8H++6e−$

The reduction half-reaction equation is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

The balanced redox equation is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Explanation of Solution

The given redox reaction equation to be balanced is shown below.

$Cr2O72−+NH4+→Cr2O3+N2$

The oxidation state of the central metal atom is calculated by knowing the standard oxidation states of few elements.

The oxidation state of chromium in $Cr2O72−$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O7n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O72n 7(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O72n+7(−2)=−2$

Calculate the value of n by simplifying the equation as shown below.

$2n+7(−2)=−22n+(−14)=−22n=−2+142n=+12$

Divide by two on both sides and simplify as shown below.

$2n2=+122n=+6$

The oxidation state of chromium in $Cr2O72−$ is $+6$ .

The oxidation state of chromium in $Cr2O3$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O3n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O32n 3(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O32n+3(−2)=0$

Calculate the value of n by simplifying the equation as shown below.

$2n+3(−2)=02n+(−6)=02n=0+62n=+6$

Divide by two on both sides and simplify as shown below.

$2n2=+62n=+3$

The oxidation state of chromium in $Cr2O3$ is $+3$ .

The oxidation state of the nitrogen in $NH4+$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$N H4n +1$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$N H4n 4(+1)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$N H4n+4(+1)=+1$

Calculate the value of n by simplifying the equation as shown below.

$n+4(+1)=+1n+4=+1n=+1−4n=−3$

The oxidation state of nitrogen is $−3$ in $NH4+$ .

The oxidation number of nitrogen in $N2$ is zero as this is the elemental form of the nitrogen.

The chromium in $Cr2O72−$ $(Cr=+6)$ is getting reduced because the oxidation state of chromium in the product $Cr2O3$ $(Cr=+3)$ is less than in the reactant. Similarly, the nitrogen is getting oxidized from the reactant $NH4+$ $(N=−3)$ to the product $N2$ $(N=0)$ .

Therefore, the oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation for the above equation is shown below.

$NH4+→N2$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The nitrogen is getting oxidized and the number of atoms of that is not balanced on both sides of the equation. Multiply ammonium ion by two to balance the nitrogen.

$2NH4+→N2$

Step-2: Balance elements other than oxygen and hydrogen if any.

$2NH4+→N2$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

$2NH4+→N2$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding the eight $H+$ on the right-hand side of the equation.

$2NH4+→N2+8H+$

Step-5: Balance the charge by adding electrons to the appropriate side.

Six electrons are added to the right-hand side in order to balance the charge.

$2NH4+→N2+8H++6e−$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$2NH4+→N2+8H++6e−$ …(1)

The reduction half-reaction for the above reaction is shown below.

$Cr2O72−→Cr2O3$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The chromium is getting reduced and its number of atoms is balanced on both sides.

$Cr2O72−→Cr2O3$

Step-2: Balance elements other than oxygen and hydrogen if any.

$Cr2O72−→Cr2O3$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

The number of oxygen atoms is balanced by adding four water molecules on the right-hand side of the equation.

$Cr2O72−→Cr2O3+4H2O$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding eight $H+$ ions on the left-hand side of the equation.

$Cr2O72−+8H+→Cr2O3+4H2O$

Step-5: Balance the charge by adding electrons to the appropriate side.

The charge is balanced by adding six electrons on the left-hand side of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$ …(2)

The balanced redox equation is obtained by adding equation (1) and (2) in such a way that electrons are canceled out.

Add equation (1) and (2) and cancel out the common things on both sides of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O+2NH4+→N2+8H++6e−$

The balance redox equation after adding these equations is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Conclusion

The oxidizer and reducer with oxidized and reduced products, oxidation and reduction half-reaction equations, and balanced redox equation are rightfully stated above.

Answer 6

Chapter 19, Problem 49E

Interpretation Introduction

Interpretation:

The oxidizer and reducer with oxidized and reduced products are to be stated. The oxidation and reduction half-reaction equations are to be stated. The balanced redox equation is to be stated.

Concept introduction:

The oxidizer is the species whose oxidation state falls during the course of reaction and reducer is the species whose oxidation number increases. Oxidized product is the oxidation product of the reducer and reduced product is the reduction product of the oxidizer.

Expert Solution & Answer

Answer to Problem 49E

The oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation is shown below.

$2NH4+→N2+8H++6e−$

The reduction half-reaction equation is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

The balanced redox equation is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Explanation of Solution

The given redox reaction equation to be balanced is shown below.

$Cr2O72−+NH4+→Cr2O3+N2$

The oxidation state of the central metal atom is calculated by knowing the standard oxidation states of few elements.

The oxidation state of chromium in $Cr2O72−$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O7n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O72n 7(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O72n+7(−2)=−2$

Calculate the value of n by simplifying the equation as shown below.

$2n+7(−2)=−22n+(−14)=−22n=−2+142n=+12$

Divide by two on both sides and simplify as shown below.

$2n2=+122n=+6$

The oxidation state of chromium in $Cr2O72−$ is $+6$ .

The oxidation state of chromium in $Cr2O3$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O3n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O32n 3(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O32n+3(−2)=0$

Calculate the value of n by simplifying the equation as shown below.

$2n+3(−2)=02n+(−6)=02n=0+62n=+6$

Divide by two on both sides and simplify as shown below.

$2n2=+62n=+3$

The oxidation state of chromium in $Cr2O3$ is $+3$ .

The oxidation state of the nitrogen in $NH4+$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$N H4n +1$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$N H4n 4(+1)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$N H4n+4(+1)=+1$

Calculate the value of n by simplifying the equation as shown below.

$n+4(+1)=+1n+4=+1n=+1−4n=−3$

The oxidation state of nitrogen is $−3$ in $NH4+$ .

The oxidation number of nitrogen in $N2$ is zero as this is the elemental form of the nitrogen.

The chromium in $Cr2O72−$ $(Cr=+6)$ is getting reduced because the oxidation state of chromium in the product $Cr2O3$ $(Cr=+3)$ is less than in the reactant. Similarly, the nitrogen is getting oxidized from the reactant $NH4+$ $(N=−3)$ to the product $N2$ $(N=0)$ .

Therefore, the oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation for the above equation is shown below.

$NH4+→N2$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The nitrogen is getting oxidized and the number of atoms of that is not balanced on both sides of the equation. Multiply ammonium ion by two to balance the nitrogen.

$2NH4+→N2$

Step-2: Balance elements other than oxygen and hydrogen if any.

$2NH4+→N2$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

$2NH4+→N2$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding the eight $H+$ on the right-hand side of the equation.

$2NH4+→N2+8H+$

Step-5: Balance the charge by adding electrons to the appropriate side.

Six electrons are added to the right-hand side in order to balance the charge.

$2NH4+→N2+8H++6e−$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$2NH4+→N2+8H++6e−$ …(1)

The reduction half-reaction for the above reaction is shown below.

$Cr2O72−→Cr2O3$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The chromium is getting reduced and its number of atoms is balanced on both sides.

$Cr2O72−→Cr2O3$

Step-2: Balance elements other than oxygen and hydrogen if any.

$Cr2O72−→Cr2O3$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

The number of oxygen atoms is balanced by adding four water molecules on the right-hand side of the equation.

$Cr2O72−→Cr2O3+4H2O$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding eight $H+$ ions on the left-hand side of the equation.

$Cr2O72−+8H+→Cr2O3+4H2O$

Step-5: Balance the charge by adding electrons to the appropriate side.

The charge is balanced by adding six electrons on the left-hand side of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$ …(2)

The balanced redox equation is obtained by adding equation (1) and (2) in such a way that electrons are canceled out.

Add equation (1) and (2) and cancel out the common things on both sides of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O+2NH4+→N2+8H++6e−$

The balance redox equation after adding these equations is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Conclusion

The oxidizer and reducer with oxidized and reduced products, oxidation and reduction half-reaction equations, and balanced redox equation are rightfully stated above.

Answer 7

Chapter 19, Problem 49E

Interpretation Introduction

Interpretation:

The oxidizer and reducer with oxidized and reduced products are to be stated. The oxidation and reduction half-reaction equations are to be stated. The balanced redox equation is to be stated.

Concept introduction:

The oxidizer is the species whose oxidation state falls during the course of reaction and reducer is the species whose oxidation number increases. Oxidized product is the oxidation product of the reducer and reduced product is the reduction product of the oxidizer.

Answer 8

Expert Solution & Answer

Answer to Problem 49E

The oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation is shown below.

$2NH4+→N2+8H++6e−$

The reduction half-reaction equation is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

The balanced redox equation is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Answer 9

Expert Solution & Answer

Answer 10

Expert Solution & Answer

Answer 11

Explanation of Solution

The given redox reaction equation to be balanced is shown below.

$Cr2O72−+NH4+→Cr2O3+N2$

The oxidation state of the central metal atom is calculated by knowing the standard oxidation states of few elements.

The oxidation state of chromium in $Cr2O72−$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O7n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O72n 7(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O72n+7(−2)=−2$

Calculate the value of n by simplifying the equation as shown below.

$2n+7(−2)=−22n+(−14)=−22n=−2+142n=+12$

Divide by two on both sides and simplify as shown below.

$2n2=+122n=+6$

The oxidation state of chromium in $Cr2O72−$ is $+6$ .

The oxidation state of chromium in $Cr2O3$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$Cr2 O3n −2$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$Cr2 O32n 3(−2)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$Cr2 O32n+3(−2)=0$

Calculate the value of n by simplifying the equation as shown below.

$2n+3(−2)=02n+(−6)=02n=0+62n=+6$

Divide by two on both sides and simplify as shown below.

$2n2=+62n=+3$

The oxidation state of chromium in $Cr2O3$ is $+3$ .

The oxidation state of the nitrogen in $NH4+$ is calculated below.

Step-1: Write down the oxidation number of every element and for unknown take “n”.

$N H4n +1$

Step-2: Multiply the oxidation state with their number of atoms of an element.

$N H4n 4(+1)$

Step-3: Add the oxidation numbers and set them equal to the charge of the species.

$N H4n+4(+1)=+1$

Calculate the value of n by simplifying the equation as shown below.

$n+4(+1)=+1n+4=+1n=+1−4n=−3$

The oxidation state of nitrogen is $−3$ in $NH4+$ .

The oxidation number of nitrogen in $N2$ is zero as this is the elemental form of the nitrogen.

The chromium in $Cr2O72−$ $(Cr=+6)$ is getting reduced because the oxidation state of chromium in the product $Cr2O3$ $(Cr=+3)$ is less than in the reactant. Similarly, the nitrogen is getting oxidized from the reactant $NH4+$ $(N=−3)$ to the product $N2$ $(N=0)$ .

Therefore, the oxidizer is $Cr2O72−$ and reducer is $NH4+$ and their corresponding reduced and oxidized product are $Cr2O3$ and $N2$ .

The oxidation half-reaction equation for the above equation is shown below.

$NH4+→N2$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The nitrogen is getting oxidized and the number of atoms of that is not balanced on both sides of the equation. Multiply ammonium ion by two to balance the nitrogen.

$2NH4+→N2$

Step-2: Balance elements other than oxygen and hydrogen if any.

$2NH4+→N2$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

$2NH4+→N2$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding the eight $H+$ on the right-hand side of the equation.

$2NH4+→N2+8H+$

Step-5: Balance the charge by adding electrons to the appropriate side.

Six electrons are added to the right-hand side in order to balance the charge.

$2NH4+→N2+8H++6e−$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$2NH4+→N2+8H++6e−$ …(1)

The reduction half-reaction for the above reaction is shown below.

$Cr2O72−→Cr2O3$

The balancing of the half-reactions is done by the following the steps shown below.

Step-1: Identify and balance the element getting oxidized or reduced.

The chromium is getting reduced and its number of atoms is balanced on both sides.

$Cr2O72−→Cr2O3$

Step-2: Balance elements other than oxygen and hydrogen if any.

$Cr2O72−→Cr2O3$

Step-3: Balance oxygen atoms by adding water on the appropriate side.

The number of oxygen atoms is balanced by adding four water molecules on the right-hand side of the equation.

$Cr2O72−→Cr2O3+4H2O$

Step-4: Balance the hydrogen atoms by adding $H+$ to the relevant side when necessary.

The number of hydrogen atoms is balanced by adding eight $H+$ ions on the left-hand side of the equation.

$Cr2O72−+8H+→Cr2O3+4H2O$

Step-5: Balance the charge by adding electrons to the appropriate side.

The charge is balanced by adding six electrons on the left-hand side of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$

Step-6: Recheck the equation to be sure that it is perfectly balanced.

The equation is completely balanced and is shown below.

$Cr2O72−+8H++6e−→Cr2O3+4H2O$ …(2)

The balanced redox equation is obtained by adding equation (1) and (2) in such a way that electrons are canceled out.

Add equation (1) and (2) and cancel out the common things on both sides of the equation.

$Cr2O72−+8H++6e−→Cr2O3+4H2O+2NH4+→N2+8H++6e−$

The balance redox equation after adding these equations is shown below.

$Cr2O72−+2NH4+→Cr2O3+4H2O+N2$

Answer 12

Conclusion

The oxidizer and reducer with oxidized and reduced products, oxidation and reduction half-reaction equations, and balanced redox equation are rightfully stated above.

Answer 13

Ch. 19 - Prob. 1E Ch. 19 - Prob. 2E Ch. 19 - Classify each of the following half-reaction...Ch. 19 - Prob. 4E Ch. 19 - Prob. 5E Ch. 19 - Prob. 6E Ch. 19 - Prob. 7E Ch. 19 - Prob. 8E Ch. 19 - Prob. 9E Ch. 19 - Prob. 10E

Answer to Problem 49E

Explanation of Solution

Want to see more full solutions like this?

Chapter 19 Solutions