Question

In each of the following reactions, identify which element is being oxidized and which is being reduced by assigning oxidation numbers. a. \(\mathrm{Fe}(s)+\mathrm{CuSO}_{4}(a q) \rightarrow \mathrm{FeSO}_{4}(a q)+\mathrm{Cu}(s)\) b. \(\mathrm{Cl}_{2}(g)+2 \mathrm{NaBr}(a q) \rightarrow 2 \mathrm{NaCl}(a q)+\mathrm{Br}_{2}(l)\) c. \(3 \mathrm{CuS}(s)+8 \mathrm{HNO}_{3}(a q) \rightarrow 3 \mathrm{CuSO}_{4}(a q)+8 \mathrm{NO}(g)+4 \mathrm{H}_{2} \mathrm{O}(l)\) d. \(2 \mathrm{Zn}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{ZnO}(s)\)

Short answer

a. Fe is oxidized (oxidation state: 0 -> +2), and Cu is reduced (oxidation state: +2 -> 0). b. Cl is reduced (oxidation state: 0 -> -1), and Br is oxidized (oxidation state: -1 -> 0). c. Cu is oxidized (oxidation state: +1 -> +2), and N is reduced (oxidation state: +5 -> +2). d. Zn is oxidized (oxidation state: 0 -> +2), and O is reduced (oxidation state: 0 -> -2).

Step-by-step solution

a. Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

First, let's determine the oxidation states for each element: Fe(s): Initial oxidation state = 0 (since it's in its elemental form) CuSO₄(aq): Initial oxidation states: Cu = +2, S = +6, and O = -2 FeSO4(aq): Final oxidation states: Fe = +2, S = +6, and O = -2 Cu(s): Final oxidation state = 0 (since it's in elemental form) Next, we can see that Fe's oxidation state has increased from 0 to +2, which means it's been oxidized. Cu's oxidation state has decreased from +2 to 0, which means it's been reduced.

b. Cl2(g) + 2 NaBr(aq) → 2 NaCl(aq) + Br2(l)

First, let's determine the initial and final oxidation states for each element: Cl2(g): Initial oxidation state = 0 (since it's in its elemental form) NaBr(aq): Initial oxidation states: Na = +1, Br = -1 NaCl(aq): Final oxidation states: Na = +1, Cl = -1 Br2(l): Final oxidation state = 0 (since it's in its elemental form) Next, we can see that Cl's oxidation state has decreased from 0 to -1, which means it's been reduced. Br's oxidation state has increased from -1 to 0, which means it's been oxidized.

c. 3 CuS(s) + 8 HNO3(aq) → 3 CuSO4(aq) + 8 NO(g) + 4 H2O(l)

First, let's determine the initial and final oxidation states for each element: CuS(s): Initial oxidation states: Cu = +1, S = -2 HNO3(aq): Initial oxidation states: H = +1, N = +5, O = -2 CuSO4(aq): Final oxidation states: Cu = +2, S = +6, and O = -2 (as previously determined in part a) NO(g): Final oxidation states: N = +2, O = -2 H2O(l): Final oxidation states: H = +1, O = -2 Next, we can see that Cu's oxidation state has increased from +1 to +2, which means it's been oxidized. N's oxidation state has decreased from +5 to +2, which means it's been reduced.

d. 2 Zn(s) + O2(g) → 2 ZnO(s)

First, let's determine the initial and final oxidation states for each element: Zn(s): Initial oxidation state = 0 (since it's in its elemental form) O2(g): Initial oxidation state = 0 (since it's in its elemental form) ZnO(s): Final oxidation states: Zn = +2, O = -2 Next, we can see that Zn's oxidation state has increased from 0 to +2, which means it's been oxidized. O's oxidation state has decreased from 0 to -2, which means it's been reduced.

Key concepts

Assigning Oxidation Numbers

Understanding how to assign oxidation numbers is crucial for identifying changes that occur to elements during chemical reactions. Oxidation numbers, often referred to as oxidation states, are a hypothetical charge that an atom would have if all bonds were ionic and electrons were transferred completely. The rules for assigning oxidation numbers include:

• For an atom in its elemental form, the oxidation number is always 0.
• For a monoatomic ion, the oxidation number is equal to the ion's charge.
• Oxygen usually has an oxidation number of -2, except in peroxides.
• Hydrogen generally has an oxidation number of +1 when bonded to nonmetals, and -1 when bonded to metals.
• The sum of the oxidation numbers for all atoms in a molecule or polyatomic ion equals the charge on the molecule or ion.

In a reaction, tracking the changes in these numbers helps determine which elements are oxidized or reduced.

Chemical Oxidation

Chemical oxidation is a process where an element's oxidation number increases, indicating a loss of electrons. In reactions, this can be conceptualized as the element donating electrons to another species. An element that is oxidized is often referred to as the reducing agent, because it causes reduction to occur to another species by providing electrons. For instance, in the reaction between iron (Fe) and copper sulfate (CuSO4), the oxidation number of Fe increases from 0 to +2, indicating that Fe is oxidized as it donates electrons to copper (Cu).

Chemical Reduction

Conversely, chemical reduction involves a decrease in an element's oxidation number, which signifies a gain of electrons. A species that undergoes reduction is gaining electrons and is referred to as the oxidizing agent, as it is causing oxidation to happen to another species by accepting electrons from it. Take the reaction between chlorine gas (Cl2) and sodium bromide (NaBr): chlorine is reduced because its oxidation number decreases from 0 to -1 as it accepts electrons from bromide (Br-), which is being oxidized.

Redox Reactions

Redox reactions, or oxidation-reduction reactions, are processes where both oxidation and reduction occur simultaneously. These reactions play a fundamental role in energy transfer and chemical transformation in biological and industrial systems. In a redox process, the element that gets oxidized transfers electrons to the element that gets reduced. This can be seen in the reaction between zinc (Zn) and oxygen (O2) to form zinc oxide (ZnO). Zn goes from an oxidation state of 0 to +2, losing electrons and hence, getting oxidized; while O2 is reduced, going from 0 to -2 as it gains electrons.