Question 15.6: Classifying Chemical Reactions as Redox or Nonredox Classify......

The oxidation numbers are calculated by the methods illustrated in Example 15.4.

a. Ni + F _2  \longrightarrow NiF _2
\quad 0     0         +2 -1
rules 1  rule 1        rules 4, 8

This is a redox reaction; the oxidation numbers of both Ni and F change. Since one substance is produced from two substances, it is also a synthesis reaction. We thus have a redox synthesis reaction.

b. Fe _2 O _3 + 3 C \longrightarrow 2 Fe + 3 CO

\quad +3  -2    0              0    +2  -2
rules 5, 8 rule 1    rule 1 rules 5, 8

This is a redox reaction; carbon is oxidized, iron is reduced. Having an element and a compound as reactants and an element and compound as products is a characteristic of a single-replacement reaction. That is the type of reaction we have here: Iron and carbon are exchanging places. We thus have a redox single-replacement reaction.

c. C _4 H _8 + 6 O _2  \longrightarrow 4 CO _2 + 4 H _2 O

\quad -2  +1    0            +4  -2     +1  -2
rules 6, 8 rule 1       rules 5, 7 rules 5, 8

This is a redox reaction; the oxidation numbers of both carbon and oxygen change. This reaction is also a combustion reaction. We thus have a redox combustion reaction.

d. H _2 SO _4 + 2 NaOH \longrightarrow Na _2 SO _4 + 2 H _2 O

    + 1 +6 -2    +1 -2 +1           +1  +6  -2    +1  -2
rules 5,6,8 rules 3,5,6     rules 3,5,8  rules 5,6

This is a nonredox reaction; there are no oxidation number changes. The reaction is also a double-replacement reaction; hydrogen and sodium are changing places, that is, “swapping partners.” Thus we have a nonredox double-replacement reaction.

Rule 1: The oxidation number of any free element (an element not combined chemically with another element) is zero.

For example, O in O _2 , P in P _4 , and S in S _8 all have an oxidation number of zero. This rule is independent of the molecular complexity of the element.

Rule 3: The oxidation numbers of groups IA and IIA elements in compounds are always + 1 and +2, respectively.

Rule 4: The oxidation number of fluorine in compounds is always -1 and that of the other group VIIA elements (Cl, Br, and I) is usually -1.

The exception for these latter elements is when they are bonded to more electronegative elements. In this case they are assigned positive oxidation numbers.

Rule 5: The usual oxidation number for oxygen in compounds is -2.

The exceptions occur when oxygen is bonded to the more electronegative fluorine (O then is assigned a positive oxidation number) or found in compounds containing oxygen-oxygen bonds (peroxides). In peroxides the oxidation number -1 is assigned to oxygen. Peroxides exist for hydrogen (H _2 O _2 ), group IA elements (Na _2 O _2 , etc.), and group IIA elements (BaO _2 , etc.).

Rule 6: The usual oxidation number for hydrogen in compounds is +1.

The exception occurs in hydrides, compounds where hydrogen is bonded to a metal of lower electronegativity. In such compounds hydrogen is assigned an oxidation number of -1. Examples of hydrides are NaH, CaH _2 , and LiH.

Rule 7: In binary compounds, the element with the greater electronegativity is as­ signed a negative oxidation number equal to its charge as an anion in its ionic compounds.

For example, in the compound AlN, N (the more electronegative element) is assigned an oxidation number of -3, the charge on a nitride ion (N ^{3-} ).

Rule 8: The algebraic sum of the oxidation numbers of all atoms in a neutral molecule must be zero.