Chapter 19, Problem 47E

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 to Problem 47E

The oxidizer is ${\text{MnO}}_4{}^{-}$ and reducer is ${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}$ and their corresponding reduced and oxidized product are ${\text{Mn}}^{2+}$ and ${\text{CO}}_2$.

The oxidation half-reaction equation is shown below.

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{CO}}_2+2{\text{e}}^{-}$

The reduction half-reaction equation is shown below.

${\text{MnO}}_4{}^{-}+{\text{8H}}^{+}+5{\text{e}}^{-}\to {\text{Mn}}^{2+}+{\text{4H}}_{\text{2}}\text{O}$

The balanced redox equation is shown below.

${\text{2MnO}}_4{}^{-}+16{\text{H}}^{+}+{\text{5C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{Mn}}^{2+}+8{\text{H}}_{\text{2}}\text{O}+10{\text{CO}}_2$

Explanation of Solution

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

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}+{\text{MnO}}_4{}^{-}\to {\text{CO}}_2+{\text{Mn}}^{2+}$

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

The oxidation state of manganese in ${\text{MnO}}_4{}^{-}$ is calculated below.

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

$\begin{array}{l}{\text{Mn O}}_4\\ \text{n }-\text{2}\end{array}$

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

$\begin{array}{l}{\text{Mn O}}_4\\ \text{n 4}\left(-\text{2}\right)\end{array}$

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

$\begin{array}{l}{\text{Mn O}}_4\\ \text{n}+\text{4}\left(-\text{2}\right)=-1\end{array}$

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

$\begin{array}{rll}\text{n}+\text{4}\left(-\text{2}\right)& =& -1\\ \text{n}+\text{(}-\text{8)}& =& -1\\ \text{n}& =& -1+8\\ \text{n}& =& +7\end{array}$

The oxidation state of ${\text{Mn}}^{2+}$ is $+2$ which comes from the charge on the manganese.

The oxidation state of the carbon in ${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}$ is calculated below.

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

$\begin{array}{l}{\text{C}}_2{\text{ O}}_4\\ \text{n }-\text{2}\end{array}$

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

$\begin{array}{l}{\text{C}}_2{\text{ O}}_4\\ \text{2n 4}\left(-\text{2}\right)\end{array}$

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

$\begin{array}{l}{\text{C}}_2{\text{ O}}_4\\ \text{2n}+\text{4}\left(-\text{2}\right)=-2\end{array}$

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

$\begin{array}{rll}\text{2n}+\text{4}\left(-\text{2}\right)& =& -2\\ \text{2n}+\text{(}-\text{8)}& =& -2\\ \text{2n}& =& -2+8\\ \text{2n}& =& +6\end{array}$

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

$\begin{array}{rll}\frac{\text{2n}}{\text{2}}& =& \frac{+6}{2}\\ \text{n}& =& +3\end{array}$

The oxidation state of carbon is $+3$ in ${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}$.

The oxidation number of carbon in ${\text{CO}}_2$ is calculated below.

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

$\begin{array}{l}{\text{C O}}_2\\ \text{n }-\text{2}\end{array}$

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

$\begin{array}{l}{\text{C O}}_2\\ \text{n 2}\left(-\text{2}\right)\end{array}$

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

$\begin{array}{l}{\text{C O}}_2\\ \text{n}+\text{2}\left(-\text{2}\right)=0\end{array}$

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

$\begin{array}{rll}\text{n}+\text{2}\left(-\text{2}\right)& =& 0\\ \text{n}-4& =& 0\\ \text{n}& =& +4\end{array}$

The oxidation state of carbon in ${\text{CO}}_2$ is $+4$.

The manganese in ${\text{MnO}}_4{}^{-}$$\left(\text{Mn}=+7\right)$ is getting reduced because the oxidation state of manganese in the product ${\text{Mn}}^{2+}$$\left(\text{Mn}=+2\right)$ is less than in the reactant. Similarly, the carbon is getting oxidized from the reactant ${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}$$\left(\text{C}=+3\right)$ to the product ${\text{CO}}_2$$\left(\text{C}=+4\right)$.

Therefore, the oxidizer is ${\text{MnO}}_4{}^{-}$ and reducer is ${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}$ and their corresponding reduced and oxidized product are ${\text{Mn}}^{2+}$ and ${\text{CO}}_2$.

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

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}\to {\text{CO}}_2$

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 carbon is getting oxidized and the number of atoms of that is not balanced on either side of reaction. Multiply ${\text{CO}}_2$ by two to balance carbon atoms as shown below.

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{CO}}_2$

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

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{CO}}_2$

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

Oxygen atoms are already balanced on both sides.

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{CO}}_2$

Step-4: Balance the hydrogen atoms by adding ${\text{H}}^{\text{+}}$ to the relevant side when necessary.

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{CO}}_2$

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

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{CO}}_2+2{\text{e}}^{-}$

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

The equation is completely balanced and is shown below.

${\text{C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{CO}}_2+2{\text{e}}^{-}$…(1)

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

${\text{MnO}}_4{}^{-}\to {\text{Mn}}^{2+}$

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 manganese is getting reduced and its number of atoms are balanced on both sides.

${\text{MnO}}_4{}^{-}\to {\text{Mn}}^{2+}$

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

${\text{MnO}}_4{}^{-}\to {\text{Mn}}^{2+}$

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.

${\text{MnO}}_4{}^{-}\to {\text{Mn}}^{2+}+{\text{4H}}_{\text{2}}\text{O}$

Step-4: Balance the hydrogen atoms by adding ${\text{H}}^{\text{+}}$ to the relevant side when necessary.

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

${\text{MnO}}_4{}^{-}+{\text{8H}}^{+}\to {\text{Mn}}^{2+}+{\text{4H}}_{\text{2}}\text{O}$

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

The charge is balanced by adding five electrons on the left-hand side as shown below.

${\text{MnO}}_4{}^{-}+{\text{8H}}^{+}+5{\text{e}}^{-}\to {\text{Mn}}^{2+}+{\text{4H}}_{\text{2}}\text{O}$

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

The equation is completely balanced and is shown below.

${\text{MnO}}_4{}^{-}+{\text{8H}}^{+}+5{\text{e}}^{-}\to {\text{Mn}}^{2+}+{\text{4H}}_{\text{2}}\text{O}$…(2)

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

Multiply equation (1) by five and equation (2) by two in order to cancel out the number of electrons on both sides as shown below.

$\begin{array}{l}{\text{5C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 10{\text{CO}}_2+10{\text{e}}^{-}\\ {\text{2MnO}}_4{}^{-}+16{\text{H}}^{+}+10{\text{e}}^{-}\to 2{\text{Mn}}^{2+}+8{\text{H}}_{\text{2}}\text{O}\end{array}$

Add the two balanced equations to get a balanced redox equation as shown below.

$\begin{array}{c}{\text{2MnO}}_4{}^{-}+16{\text{H}}^{+}+10{\text{e}}^{-}\to 2{\text{Mn}}^{2+}+8{\text{H}}_{\text{2}}\text{O}\\ +\\ {\text{5C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 10{\text{CO}}_2+10{\text{e}}^{-}\end{array}$

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

${\text{2MnO}}_4{}^{-}+16{\text{H}}^{+}+{\text{5C}}_{\text{2}}{\text{O}}_4{}^{2-}\to 2{\text{Mn}}^{2+}+8{\text{H}}_{\text{2}}\text{O}+10{\text{CO}}_2$

Conclusion

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