Balancing Redox Reactions by the Oxidation Number Method (11.1.2)

Question 1

• Formal charge and oxidation state (or oxidation number) are both models of what is occurring with electrondistribution
in a molecule or ion.

Answer 1

• Formal charge and oxidation state (or oxidation number) are both models of what is occurring with electrondistribution
in a molecule or ion.

Question 2

• A reaction is balanced by assigning oxidation states, balancing the transferred electrons, and balancing the redox atoms.

Answer 2

• A reaction is balanced by assigning oxidation states, balancing the transferred electrons, and balancing the redox atoms.

Question 3

• Balancing atoms of oxygen and hydrogen depends on the solution.

Answer 3

• Balancing atoms of oxygen and hydrogen depends on the solution.

Question 4

Formal charge and oxidation state (or oxidation number) are
both models of what is occurring with electron distribution
in a molecule or ion.
Formal charge reflects the average of the charge on an atom
in a molecule. For oxidation number, the electrons are
assigned entirely to the more electronegative atom. Oxidation
number is a formalism that facilitates the balancing of
oxidation-reduction reactions.
A reaction is balanced by assigning oxidation states, balancing
the transferred electrons, and balancing the redox atoms.
In the example, first oxidation numbers are assigned to each
redox atom. The values shown are per atom and the values
for oxygen atoms are not shown. Next, the redox atoms (Cr
and Cl) are balanced, and then the electrons transferred are
identified based on the change in oxidation number of each
atom.
Electrons are balanced by multiplying each half reaction by
the lowest common factor (4 and 3), and these coefficients
are applied to the redox molecules (by multiplication) to
complete the balancing as shown in the box. Only the oxygen
atoms remain unbalanced.
Balancing atoms of oxygen and hydrogen depends on the
solution.
In an acidic solution, the oxygen atoms are balanced with
H2O molecules and the hydrogen atoms with
H+ ions. In the example, the left side of the equation has 40
oxygen atoms and the right side has 21 oxygen atoms.
Adding 19 H2O molecules (making up the difference) to the
right side to balance oxygen atoms. Adding 38 H+ ions to
the left side accounts for the added H+ ions.
In a basic solution, the oxygen atoms are balanced with 2OH–
ions and the hydrogen atoms with H2O molecules. In the
example, add 38 OH– ions to the right side to balance the
oxygen atoms and 19 H2O molecules to the left side.

Answer 4

Formal charge and oxidation state (or oxidation number) are
both models of what is occurring with electron distribution
in a molecule or ion.
Formal charge reflects the average of the charge on an atom
in a molecule. For oxidation number, the electrons are
assigned entirely to the more electronegative atom. Oxidation
number is a formalism that facilitates the balancing of
oxidation-reduction reactions.
A reaction is balanced by assigning oxidation states, balancing
the transferred electrons, and balancing the redox atoms.
In the example, first oxidation numbers are assigned to each
redox atom. The values shown are per atom and the values
for oxygen atoms are not shown. Next, the redox atoms (Cr
and Cl) are balanced, and then the electrons transferred are
identified based on the change in oxidation number of each
atom.
Electrons are balanced by multiplying each half reaction by
the lowest common factor (4 and 3), and these coefficients
are applied to the redox molecules (by multiplication) to
complete the balancing as shown in the box. Only the oxygen
atoms remain unbalanced.
Balancing atoms of oxygen and hydrogen depends on the
solution.
In an acidic solution, the oxygen atoms are balanced with
H2O molecules and the hydrogen atoms with
H+ ions. In the example, the left side of the equation has 40
oxygen atoms and the right side has 21 oxygen atoms.
Adding 19 H2O molecules (making up the difference) to the
right side to balance oxygen atoms. Adding 38 H+ ions to
the left side accounts for the added H+ ions.
In a basic solution, the oxygen atoms are balanced with 2OH–
ions and the hydrogen atoms with H2O molecules. In the
example, add 38 OH– ions to the right side to balance the
oxygen atoms and 19 H2O molecules to the left side.

Question 5

Which of the following correctly summarizes the electron transfer in the given reaction?

10H +(aq) + 8I −(aq) + SO42−(aq) → H2S(aq) + 4I2(aq) + 4H2O(l)

Answer 5

I − → I 0; S6+ → S2− (D)

I − goes to I 0 (in the product species I2 ). S6+ → S2− in the product species H2S by the transfer of eight electrons (to the iodine atoms).

Question 6

Which of the following correctly identifies the oxidizing agent and the reducing agent?

3H2S(aq) + 2NO3−(aq) + 2H +(aq) → 3S(s) + 2NO(g) + 4H2O(l)

Answer 6

H2S is the reducing agent. NO3− is the oxidizing agent. (C)

In this reaction, the oxidation numbers for oxygen and hydrogen do not change. Sulfur’s oxidation number increases (from −2 to 0) and nitrogen’s oxidation number decreases (from +5 to +2). Therefore, H2S is the reducing agent and NO3 is the oxidizing agent.

Question 7

Which of the following correctly shows the oxidation numbers for each of the elements in H2PO4−, I2, and Al(OH)4−?

Answer 7

H2PO4−: H = +1, P = +5, O = −2; I2 : I = 0; Al(OH)4− : Al = +3; O = −2; H = +1 (A)

Any neutral compound’s component species have oxidation states that sum to zero. Oxygen has an oxidation number of −2 in these compounds. Hydrogen has an oxidation number of +1 in these compounds. With this in mind, we can easily calculate the oxidation number of any other component in the compound. In H2PO4−, the overall charge is −1.

Therefore, 2 (+1) + (the oxidation number of P) + 4 (−2) = −1. (The “+1” refers to the oxidation number of hydrogen. The “-2” refers to the oxidation number of oxygen.) The oxidation number of P = +5. In Al(OH)4− the overall charge is −1.

Therefore, (the oxidation number of Al) + 4 (−2) + 4 (+1) = −1. (The “+1” refers to the oxidation number of hydrogen. The “-2” refers to the oxidation number of oxygen.) The oxidation number of Al = +3.

Question 8

Which of the following best describes what a redox reaction is?

Answer 8

It is a reaction in which the oxidation state of one element increases and the oxidation state of another element decreases. (B)

The term redox stands for reduction-oxidation. A redox reaction is a reaction in which one element is oxidized (causing an increase in its oxidation state) and another element is reduced (causing a decrease in its oxidation state).

Question 9

Which statement about the following redox reaction in acidic solution is not correct?

H2S(aq) + Cr2O72− (aq) → S(s) + Cr3+ (aq)

Answer 9

The final step involves adding oxygen atoms to the left side of the reaction and hydrogen atoms to the right side of the reaction. (D)

This statement is not correct. You can see from the given reaction that there are no oxygen on the right side. Therefore, you will have to add H2O molecules, not oxygen atoms, to the right side and H + atoms to the left side to balance the reaction.

Question 10

Which of the following shows the correct oxidation states for iodine, sulfur, and oxygen in the products of this reaction?

10H +(aq) + 8I −(aq) + SO42−(aq) → H2S(aq) + 4I2(aq) + 4H2O(l)

Answer 10

I is 0; S is −2; O is −2 (B)

Iodine exists as I2(aq) in the product state. Its oxidation number is 0. Sulfur, in H2S(aq), has an oxidation state of −2 to counterbalance the +2 from the two hydrogen atoms. Oxygen has a constant oxidation state of −2 in this reaction.

Question 11

Which of the following is not a correct rule for determining oxidation states?

Answer 11

Any compound’s component species have oxidation states that add up to zero. (C)

In ionic compounds, the sum of the oxidation states is equal to the charge on the overall ionic molecule (i.e., not equal to 0). For example, SO42− has one sulfur atom (oxidation state of +6) and four oxygen atoms (each with an oxidation state of −2). The overall charge of the compound is equal to the sum of the individual oxidation states, in this case-2.

Question 12

Which of the following correctly identifies the oxidizing agent, the reducing agent, the oxidized substance, and the reduced substance in the reaction?

6Fe2+(aq) + Cr2O72−(aq) + 14H +(aq) → 6Fe3+(aq) + 2Cr3+(aq) + 7H2O(l)

Answer 12

Cr2O72− is the oxidizing agent. Fe2+ is the reducing agent. Fe2+ is the oxidized substance. Cr2O72− is the reduced substance. (D)

Fe2+ is oxidized to become Fe3+. It is also the reducing agent. Chromium’s oxidation number goes from +6 in Cr2O72− to +3 in Cr3+. Therefore, Cr2O72− is reduced and is also the oxidizing agent.

Question 13

(Un-copy-able question)

Answer 13

(B)

Question 14

Which of the following correctly shows the oxidation numbers for each of the elements in H2PO4−, I2, and Al(OH)4−?

Answer 14

H2PO4−: H = +1, P = +5, O = −2; I2 : I = 0; Al(OH)4− : Al = +3; O = −2; H = +1 (A)