Problem 54 - Chemistry: The Central Science

Question

Write balanced molecular and net ionic equations for the reactions of (a) hydrochloric acid with nickel; (b) dilute sulfuric acid with iron; (c) hydrobromic acid with magnesium; (d) acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH}\), with zinc.

Step-by-Step Solution

Answer

In summary, the balanced molecular and net ionic equations for the reactions are: (a) Hydrochloric acid with nickel: Molecular: \( Ni (s) + 2HCl (aq) \rightarrow NiCl_{2} (aq) + H_{2} (g) \) Net ionic: \( Ni (s) + 2H^{+} (aq) \rightarrow Ni^{2+} (aq) + H_{2} (g) \) (b) Dilute sulfuric acid with iron: Molecular: \( Fe (s) + H_{2}SO_{4} (aq) \rightarrow FeSO_{4} (aq) + H_{2} (g) \) Net ionic: \( Fe (s) + 2H^{+} (aq) \rightarrow Fe^{2+} (aq) + H_{2} (g) \) (c) Hydrobromic acid with magnesium: Molecular: \( Mg (s) + 2HBr (aq) \rightarrow MgBr_{2} (aq) + H_{2} (g) \) Net ionic: \( Mg (s) + 2H^{+} (aq) \rightarrow Mg^{2+} (aq) + H_{2} (g) \) (d) Acetic acid with zinc: Molecular: \( Zn (s) + 2CH_{3}COOH (aq) \rightarrow Zn(CH_{3}COO)_{2} (aq) + H_{2} (g) \) Net ionic: \( Zn (s) + 2CH_{3}COOH (aq) \rightarrow Zn^{2+} (aq) + 2CH_{3}COO^{-} (aq) + H_{2} (g) \)

Step 1: Reaction A: Hydrochloric acid with nickel

1. Identify the reactants and possible products. The reactants are hydrochloric acid (HCl) and nickel (Ni). The reaction is between an acid and a metal, producing a salt and hydrogen gas (H₂). 2. Write an unbalanced molecular equation. \[ Ni (s) + HCl (aq) \rightarrow NiCl_{2} (aq) + H_{2} (g) \] 3. Balance the molecular equation. \[ Ni (s) + 2HCl (aq) \rightarrow NiCl_{2} (aq) + H_{2} (g) \] 4. Write the net ionic equation. First, split the strong electrolytes into ions: \[ Ni (s) + 2H^{+} (aq) + 2Cl^{-} (aq) \rightarrow Ni^{2+} (aq) + 2Cl^{-} (aq) + H_{2} (g) \] Then eliminate the spectator ion(s): \[ Ni (s) + 2H^{+} (aq) \rightarrow Ni^{2+} (aq) + H_{2} (g) \]

Step 2: Reaction B: Dilute sulfuric acid with iron

1. Identify the reactants and possible products. The reactants are dilute sulfuric acid (H₂SO₄) and iron (Fe). The reaction is between an acid and a metal, producing a salt and hydrogen gas (H₂). 2. Write an unbalanced molecular equation. \[ Fe (s) + H_{2}SO_{4} (aq) \rightarrow FeSO_{4} (aq) + H_{2} (g) \] 3. Balance the molecular equation. Since the molecular equation is already balanced, there is no need to make any adjustments. 4. Write the net ionic equation. First, split the strong electrolyte(s) into ions: \[ Fe (s) + 2H^{+} (aq) + SO_{4}^{2-} (aq) \rightarrow Fe^{2+} (aq) + SO_{4}^{2-} (aq) + H_{2} (g) \] Then eliminate the spectator ion(s): \[ Fe (s) + 2H^{+} (aq) \rightarrow Fe^{2+} (aq) + H_{2} (g) \]

Step 3: Reaction C: Hydrobromic acid with magnesium

1. Identify the reactants and possible products. The reactants are hydrobromic acid (HBr) and magnesium (Mg). The reaction is between an acid and a metal, producing a salt and hydrogen gas (H₂). 2. Write an unbalanced molecular equation. \[ Mg (s) + HBr (aq) \rightarrow MgBr_{2} (aq) + H_{2} (g) \] 3. Balance the molecular equation. \[ Mg (s) + 2HBr (aq) \rightarrow MgBr_{2} (aq) + H_{2} (g) \] 4. Write the net ionic equation. First, split the strong electrolytes into ions: \[ Mg (s) + 2H^{+} (aq) + 2Br^{-} (aq) \rightarrow Mg^{2+} (aq) + 2Br^{-} (aq) + H_{2} (g) \] Then eliminate the spectator ion(s): \[ Mg (s) + 2H^{+} (aq) \rightarrow Mg^{2+} (aq) + H_{2} (g) \]

Step 4: Reaction D: Acetic acid with zinc

1. Identify the reactants and possible products. The reactants are acetic acid (CH₃COOH) and zinc (Zn). The reaction is between an (weak) acid and a metal, producing a salt and hydrogen gas (H₂). 2. Write an unbalanced molecular equation. \[ Zn (s) + CH_{3}COOH (aq) \rightarrow Zn(CH_{3}COO)_{2} (aq) + H_{2} (g) \] 3. Balance the molecular equation. \[ Zn (s) + 2CH_{3}COOH (aq) \rightarrow Zn(CH_{3}COO)_{2} (aq) + H_{2} (g) \] 4. Write the net ionic equation. Since acetic acid is a weak electrolyte, we will not split it into ions. Only zinc acetate is a strong electrolyte and will be split. \[ Zn (s) + 2CH_{3}COOH (aq) \rightarrow Zn^{2+} (aq) + 2CH_{3}COO^{-} (aq) + H_{2} (g) \]

Key Concepts

Chemical Reactions

Chemical reactions are processes where substances, known as reactants, transform into new substances, referred to as products. These transformations occur due to the breaking and forming of chemical bonds, following the law of conservation of matter.

In the context of the provided exercise, the process of an acid reacting with a metal is a type of chemical reaction that creates a salt and hydrogen gas, demonstrating how reactants recombine in different ways to produce new products.

To help students better visualize and understand these transformations, it's beneficial to use molecular equations, which represent the chemicals in the form they are introduced to the reaction, and net ionic equations, which simplify the reaction to its essence by removing spectator ions that do not partake in the actual chemical change.

Acid-Metal Reactions

Acid-metal reactions are a subset of chemical reactions where an acid reacts with a metal to produce a salt and hydrogen gas. The general form of an acid-metal reaction can be represented as:
\[ \text{Acid (aq) } + \text{ Metal (s) } \rightarrow \text{Salt (aq) } + H_{2}(g) \]

This type of reaction is also an example of a single displacement reaction, where the metal displaces the hydrogen from the acid. The reactivity of the metal and the strength of the acid can influence the reaction rate and products. For instance, hydrochloric acid (a strong acid) will react vigorously with metals, whereas acetic acid (a weak acid) will react less vigorously. It's important for students to distinguish between strong and weak acids, as this affects whether the acid dissociates completely in water, which influences the net ionic equation for the reaction.

Balancing Chemical Equations

Balancing chemical equations is crucial because it ensures that the law of conservation of mass is obeyed in the chemical reaction. The number of atoms for each element in the reactants must equal the number of atoms for those elements in the products.

In practice, balancing an equation involves adjusting coefficients, which are numbers placed in front of chemical formulas, to ensure that the atom count is the same on both sides of an equation. For example, when balancing \( HCl \) with \( Ni \) to form \( NiCl_{2} \) and \( H_{2} \), we find that there are two chlorine atoms on the product side so we need two \( HCl \) molecules on the reactant side to balance the chlorine atoms. Students often struggle with this concept, but remembering the law of conservation of mass can guide them through properly balancing the equations. Visual tools such as atom inventory lists or interactive simulations can be quite beneficial in aiding student understanding of this concept.