Problem 56 - Chemistry The Central Science

Question

Write a balanced net ionic equation for each of the following reactions: (a) Dilute nitric acid reacts with zinc metal with formation of nitrous oxide. (b) Concentrated nitric acid reacts with sulfur with formation of nitrogen dioxide. (c) Concentrated nitric acid oxidizes sulfur dioxide with formation of nitric oxide. (d) Hydrazine is burned in excess fluorine gas, forming \(\mathrm{NF}_{3}\). (e) Hydrazine reduces \(\mathrm{CrO}_{4}^{2-}\) to \(\mathrm{Cr}(\mathrm{OH})_{4}^{-}\) in base (hydrazine is oxidized to \(\mathrm{N}_{2}\) ).

Step-by-Step Solution

Answer

(a) \(Zn + 4H^+ + 4NO^-_2 \rightarrow Zn^{2+} + 2NO + 2NO_2 + 2H_2O\) (b) \(6H^+ + 6NO^-_3 + S \rightarrow 2H^+ + SO^{2-}_4 + 6NO_2 + 2H_2O\) (c) \(2NO^-_3 + SO_2 \rightarrow NO + NO_2 + SO^{2-}_4\) (d) \(2F_2 + N_2H_4 \rightarrow 4NF_3\) (e) \(N_2H_4 + 2Cr^{6+} + 8O^{2-} + OH^- \rightarrow N_2 + 2Cr^{3+} + 8OH^-\)

Step 1: Reaction (a) - Dilute nitric acid reacts with zinc metal with formation of nitrous oxide

First, write the balanced molecular equation: \(Zn + 4HNO_2 \rightarrow Zn^{2+} + 2NO + 2NO_2 + 2H_2O\) Now, break apart the strong electrolytes into their ions to get the total ionic equation: \(Zn + 4H^+ + 4NO^-_2 \rightarrow Zn^{2+} + 2NO + 2NO_2 + 2H_2O\) Group the same ions on both sides and cross out the spectator ions (none in this case). The remaining equation is the net ionic equation: \(Zn + 4H^+ + 4NO^-_2 \rightarrow Zn^{2+} + 2NO + 2NO_2 + 2H_2O\)

Step 2: Reaction (b) - Concentrated nitric acid reacts with sulfur with formation of nitrogen dioxide

First, write the balanced molecular equation: \(6HNO_3 + S \rightarrow H_2SO_4 + 6NO_2 + 2H_2O\) Next, break apart the strong electrolytes into their ions to get the total ionic equation: \(6H^+ + 6NO^-_3 + S \rightarrow 2H^+ + SO^{2-}_4 + 6NO_2 + 2H_2O\) Group the same ions on both sides and cross out the spectator ions (none in this case). The remaining equation is the net ionic equation: \(6H^+ + 6NO^-_3 + S \rightarrow 2H^+ + SO^{2-}_4 + 6NO_2 + 2H_2O\)

Step 3: Reaction (c) - Concentrated nitric acid oxidizes sulfur dioxide with formation of nitric oxide

First, write the balanced molecular equation: \(2HNO_3 + SO_2 \rightarrow NO + NO_2 + H_2SO_4\) Next, break apart the strong electrolytes into their ions to get the total ionic equation: \(2H^+ + 2NO^-_3 + SO_2 \rightarrow NO + NO_2 + 2H^+ + SO^{2-}_4\) Group the same ions on both sides and cross out the spectator ions (H^+ in this case). The remaining equation is the net ionic equation: \(2NO^-_3 + SO_2 \rightarrow NO + NO_2 + SO^{2-}_4\)

Step 4: Reaction (d) - Hydrazine is burned in excess fluorine gas, forming \(\mathrm{NF}_{3}\)

First, write the balanced molecular equation: \(2F_2 + N_2H_4 \rightarrow 4NF_3\) This reaction does not involve ions, so no further decomposition is needed. The net ionic equation is the same as the molecular equation: \(2F_2 + N_2H_4 \rightarrow 4NF_3\)

Step 5: Reaction (e) - Hydrazine reduces \(\mathrm{CrO}_{4}^{2-}\) to \(\mathrm{Cr}(\mathrm{OH})_{4}^{-}\) in base (hydrazine is oxidized to \(\mathrm{N}_{2}\) )

First, write the balanced molecular equation: \(N_2H_4 + 2CrO^{2-}_4 + OH^- \rightarrow N_2 + 2Cr(OH^-)_4\) Next, break apart the strong electrolytes into their ions to get the total ionic equation: \(N_2H_4 + 2Cr^{6+} + 8O^{2-} + OH^- \rightarrow N_2 + 2Cr^{3+} + 8OH^-\) Group the same ions on both sides and cross out the spectator ions (none in this case). The remaining equation is the net ionic equation: \(N_2H_4 + 2Cr^{6+} + 8O^{2-} + OH^- \rightarrow N_2 + 2Cr^{3+} + 8OH^-\)

Key Concepts

Acid-Base Reactions

Acid-base reactions are foundational concepts in chemistry where an acid reacts with a base to produce water and a salt. An acid is a substance that can donate a proton ( H^+ ) and a base is a substance that can accept a proton. When an acid is mixed with a base, the H^+ ions from the acid and the OH^- ions from the base combine to form water. This is often observed as a neutralization reaction.

Key points to remember when dealing with acid-base reactions are:

• Neutralization: Occurs when equal amounts of an acid and base react to form water and a salt.
• Indicators: Substances that change color in the presence of an acid or base, helping to identify the endpoint of a titration.
• pH scale: A scale used to measure how acidic or basic a solution is.

In the exercise, reactions involving nitric acid and metals are examples of how acids can react with other substances to produce different results, highlighting different aspects of their properties.

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between chemical species. In these reactions, one element or compound gets oxidized (loses electrons) while another gets reduced (gains electrons). This simultaneous process often involves changes in the oxidation state of the elements involved.

A helpful way to remember the process is the acronym OIL RIG: Oxidation Is Loss, Reduction Is Gain.

• Oxidation: A process where a substance loses electrons.
• Reduction: A process where a substance gains electrons.
• Oxidizing agent: The species that is reduced and thus helps another to be oxidized.
• Reducing agent: The species that is oxidized and thus helps another to be reduced.

The reactions in the exercise show various redox processes, such as the oxidation of sulfur and the reducing action of hydrazine, illustrating how redox reactions play a key role in chemical transformations.

Chemical Equations

Chemical equations are symbolic representations of chemical reactions where reactants are transformed into products. They consist of chemical formulas for the substances involved, separated by an arrow indicating the direction of the transformation.

An essential aspect of chemical equations is balancing. Here, the number of atoms for each element must be the same on both sides of the equation. This satisfies the law of conservation of mass which states that mass in an isolated system is neither created nor destroyed by chemical reactions or physical transformations.

• Reactants: Substances that start the reaction. They are found on the left side of the equation.
• Products: Substances formed as a result of the reaction. They are found on the right side.
• Coefficients: Numbers placed before formulas to balance the equation.
• States of matter: Often designated in parentheses, such as (s) for solids, (l) for liquids, (g) for gases, and (aq) for aqueous solutions.

In the exercise, each reaction starts with writing the balanced molecular equation before breaking it down into ionic forms, emphasizing the importance of correctly balancing chemical equations.

Oxidation States

Oxidation states, also known as oxidation numbers, are a way of keeping track of electrons in atoms participating in chemical reactions. Each element in a compound is assigned an oxidation state which helps in determining how electrons are distributed in a chemical reaction.

The rules to determine oxidation states include:

• Free elements have an oxidation state of zero.
• For ions, the oxidation state is equal to the charge of the ion.
• Oxygen usually has an oxidation state of -2, while hydrogen is typically +1.
• The sum of oxidation states in a neutral compound is zero; in a polyatomic ion, it equals the ion's charge.

In the step-by-step solutions, understanding oxidation states is crucial to identifying which elements are oxidized and which are reduced. This knowledge is essential for balancing redox reactions, as it reveals the electron transfer process, a vital part of these reactions.