Problem 109 - Chemistry: An Atoms-Focused Approach

Question

For each of the molecular equations, write net ionic equations and identify the Bronsted-Lowry acids and bases: a. \(2 \mathrm{HNO}_{3}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(a q) \rightarrow\) \(2 \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(a q)\) b. \(\mathrm{Na}_{2} \mathrm{CO}_{3}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow\) \(\mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(\ell)\) c. \(\mathrm{CH}_{3} \mathrm{NH}_{2}(a q)+\mathrm{HBr}(a q) \rightarrow\left(\mathrm{CH}_{3} \mathrm{NH}_{3}\right) \mathrm{Br}(a q)\) d. \(2 \mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{Mg}(\mathrm{OH})_{2}(s) \rightarrow\) \(\left(\mathrm{CH}_{3} \mathrm{COO}\right)_{2} \mathrm{Mg}(a q)+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) e. \(\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Ca}(\mathrm{OH})_{2}(s)\) f. \(\operatorname{LiH}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \operatorname{LiOH}(a q)+\mathrm{H}_{2}(g)\) g. \(\mathrm{Ba}(\mathrm{OH})_{2}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow \mathrm{BaSO}_{4}(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) h. \(\mathrm{NaSH}(a q)+\mathrm{HNO}_{3}(a q) \rightarrow \mathrm{NaNO}_{3}(a q)+\mathrm{H}_{2} \mathrm{S}(g)\)

Step-by-Step Solution

Answer

Question: Write net ionic equations and identify the Bronsted-Lowry acids and bases for the following molecular equations. a. \(2 \mathrm{HNO}_{3}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(a q) \rightarrow 2 \mathrm{H}_{2}\mathrm{O}(\ell)+\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(a q)\) Net Ionic Equation: \(2 \mathrm{H}^{+}(a q) + 2\mathrm{OH}^{-}(a q) \rightarrow 2 \mathrm{H}_{2}\mathrm{O}(\ell)\) Acid: \(2 \mathrm{HNO}_{3}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{Ca}(\mathrm{OH})_{2}\) (accepts \(\mathrm{H}^{+}\)) b. \(\mathrm{Na}_{2} \mathrm{CO}_{3}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2}\mathrm{O}(\ell)\) Net Ionic Equation: \(\mathrm{CO}_{3}^{2-}(a q) + 2\mathrm{H}^{+}(a q) \rightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2}\mathrm{O}(\ell)\) Acid: \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) (accepts \(\mathrm{H}^{+}\)) c. \(\mathrm{CH}_{3} \mathrm{NH}_{2}(a q)+\mathrm{HBr}(a q) \rightarrow\left(\mathrm{CH}_{3} \mathrm{NH}_{3}\right) \mathrm{Br}(a q)\) Net Ionic Equation: \(\mathrm{CH}_{3}\mathrm{NH}_{2}(a q) + \mathrm{H}^{+}(a q) \rightarrow\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}(a q)\) Acid: \(\mathrm{HBr}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) (accepts \(\mathrm{H}^{+}\))

Step 1: Balanced Molecular Equation

The equation is already balanced.

Step 2: Total Ionic Equation

Split the aqueous compounds into their respective ions: $2 \mathrm{H}^{+}(a q) + 2 \mathrm{NO}_{3}^{-}(a q) + \mathrm{Ca}^{2+}(a q) + 2\mathrm{OH}^{-}(a q) \rightarrow 2 \mathrm{H}_{2}\mathrm{O}(\ell) + \mathrm{Ca}^{2+}(a q) + 2\mathrm{NO}_{3}^{-}(a q)$

Step 3: Net Ionic Equation

Eliminate the spectator ions, which are \(\mathrm{Ca}^{2+}\) and \(\mathrm{NO}_{3}^{-}\): \(2 \mathrm{H}^{+}(a q) + 2\mathrm{OH}^{-}(a q) \rightarrow 2 \mathrm{H}_{2}\mathrm{O}(\ell)\)

Step 4: Identify Bronsted-Lowry acids and bases

Acid: \(2 \mathrm{HNO}_{3}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{Ca}(\mathrm{OH})_{2}\) (accepts \(\mathrm{H}^{+}\)) We proceed with the other equations in the same manner: b. \(\mathrm{Na}_{2} \mathrm{CO}_{3}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2}\mathrm{O}(\ell)\) Total Ionic Equation: \(2 \mathrm{Na}^{+}(a q) + \mathrm{CO}_{3}^{2-}(a q) + 2\mathrm{H}^{+}(a q) + \mathrm{SO}_{4}^{2-}(a q) \rightarrow 2 \mathrm{Na}^{+}(a q) + \mathrm{SO}_{4}^{2-}(a q) + \mathrm{CO}_{2}(g)+\mathrm{H}_{2}\mathrm{O}(\ell)\) Net Ionic Equation: \(\mathrm{CO}_{3}^{2-}(a q) + 2\mathrm{H}^{+}(a q) \rightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2}\mathrm{O}(\ell)\) Acid: \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) (accepts \(\mathrm{H}^{+}\)) c. \(\mathrm{CH}_{3} \mathrm{NH}_{2}(a q)+\mathrm{HBr}(a q) \rightarrow\left(\mathrm{CH}_{3} \mathrm{NH}_{3}\right) \mathrm{Br}(a q)\) Total Ionic Equation: \(\mathrm{CH}_{3}\mathrm{NH}_{2}(a q) + \mathrm{H}^{+}(a q) + \mathrm{Br}^{-}(a q) \rightarrow\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}(a q) + \mathrm{Br}^{-}(a q)\) Net Ionic Equation: \(\mathrm{CH}_{3}\mathrm{NH}_{2}(a q) + \mathrm{H}^{+}(a q) \rightarrow\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}(a q)\) Acid: \(\mathrm{HBr}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) (accepts \(\mathrm{H}^{+}\)) d. \(2 \mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{Mg}(\mathrm{OH})_{2}(s) \rightarrow \left(\mathrm{CH}_{3} \mathrm{COO}\right)_{2} \mathrm{Mg}(a q)+2\mathrm{H}_{2}\mathrm{O}(\ell)\) Total Ionic Equation: \(2 \mathrm{CH}_{3}\mathrm{COO}^{-}(a q)+ 2\mathrm{H}^{+}(a q) + \mathrm{Mg}^{2+}(s) + 2\mathrm{OH}^{-}(s) \rightarrow \left(\mathrm{CH}_{3} \mathrm{COO}\right)_{2} \mathrm{Mg}^{2+}(a q) +2\mathrm{H}_{2}\mathrm{O}(\ell)\) Net Ionic Equation: \(2 \mathrm{CH}_{3}\mathrm{COO}^{-}(a q)+ 2\mathrm{H}^{+}(a q) + \mathrm{Mg}^{2+}(s) + 2\mathrm{OH}^{-}(s) \rightarrow \left(\mathrm{CH}_{3} \mathrm{COO}\right)_{2} \mathrm{Mg}^{2+}(a q) +2\mathrm{H}_{2}\mathrm{O}(\ell)\) Acid: \(2 \mathrm{CH}_{3} \mathrm{COOH}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{Mg}(\mathrm{OH})_{2}\) (accepts \(\mathrm{H}^{+}\)) e. \(\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Ca}(\mathrm{OH})_{2}(s)\) Total Ionic Equation: \(\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Ca}^{2+}(s) + 2\mathrm{OH}^{-}(s)\) Net Ionic Equation: \(\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Ca}^{2+}(s) + 2\mathrm{OH}^{-}(s)\) Acid: \(\mathrm{H}_{2} \mathrm{O}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{CaO}\) (accepts \(\mathrm{H}^{+}\)) f. \(\operatorname{LiH}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \operatorname{LiOH}(a q)+\mathrm{H}_{2}(g)\) Total Ionic Equation: \(\operatorname{LiH}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Li}^{+}(a q) + \mathrm{OH}^{-}(a q)+\mathrm{H}_{2}(g)\) Net Ionic Equation: \(\operatorname{LiH}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Li}^{+}(a q) + \mathrm{OH}^{-}(a q)+\mathrm{H}_{2}(g)\) Acid: \(\mathrm{H}_{2} \mathrm{O}\) (donates \(\mathrm{H}^{+}\)) Base: \(\operatorname{LiH}\) (accepts \(\mathrm{H}^{+}\)) g. \(\mathrm{Ba}(\mathrm{OH})_{2}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow \mathrm{BaSO}_{4}(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) Total Ionic Equation: \(\mathrm{Ba}^{2+}(a q) + 2\mathrm{OH}^{-}(a q) + 2\mathrm{H}^{+}(a q) + \mathrm{SO}_{4}^{2-}(a q) \rightarrow \mathrm{Ba}^{2+}(s) + \mathrm{SO}_{4}^{2-}(s) +2 \mathrm{H}_{2}\mathrm{O}(\ell)\) Net Ionic Equation: \(2\mathrm{OH}^{-}(a q) + 2\mathrm{H}^{+}(a q) \rightarrow 2 \mathrm{H}_{2}\mathrm{O}(\ell)\) Acid: \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{Ba}(\mathrm{OH})_{2}\) (accepts \(\mathrm{H}^{+}\)) h. \(\mathrm{NaSH}(a q)+\mathrm{HNO}_{3}(a q) \rightarrow \mathrm{NaNO}_{3}(a q)+\mathrm{H}_{2} \mathrm{S}(g)\) Total Ionic Equation: \(\mathrm{Na}^{+}(a q) + \mathrm{HS}^{-}(a q) + \mathrm{H}^{+}(a q) + \mathrm{NO}_{3}^{-}(a q) \rightarrow \mathrm{Na}^{+}(a q) + \mathrm{NO}_{3}^{-}(a q)+\mathrm{H}_{2} \mathrm{S}(g)\) Net Ionic Equation: \(\mathrm{HS}^{-}(a q) + \mathrm{H}^{+}(a q) \rightarrow \mathrm{H}_{2} \mathrm{S}(g)\) Acid: \(\mathrm{HNO}_{3}\) (donates \(\mathrm{H}^{+}\)) Base: \(\mathrm{NaSH}\) (accepts \(\mathrm{H}^{+}\))

Key Concepts

Bronsted-Lowry acids and bases

In chemistry, acids and bases play critical roles in reactions, especially in solution chemistry. The Bronsted-Lowry theory defines acids as substances that donate protons ( \(H^{+}\)), while bases are those that accept protons. This concept helps us identify the nature of a substance in a reaction. For example, in reaction \(a.\) \(\text{2 HNO}_{3}\) acts as a Bronsted-Lowry acid because it donates \(\text{2 H}^{+}\). Correspondingly, \(\text{Ca} \text{(OH)}_{2}\) is a Bronsted-Lowry base as it accepts these protons. Understanding acids and bases in this context allows us to predict and balance equations effectively by recognizing which molecules will donate or accept protons. This approach can be straightforward for common strong acids and bases—like sulfuric acid and sodium carbonate in reaction \(b.\)—where the donation and acceptance of protons usually involve complete ionization or dissociation.

Chemical Reactions

Chemical reactions involve the transformation of reactants into products. In many cases, including those involving Bronsted-Lowry acids and bases, the main goal is to reach a state of equilibrium. Key characteristics of these reactions include:

• Reactants and Products: Reactants undergo a chemical change to form products, such as in the equation for reaction \(g.\) where \(\text{Ba} \text{(OH)}_2\) and \(\text{H}_2\text{SO}_4\) form \(\text{BaSO}_4\) and water.
• Energy Changes: Energy is either absorbed or released during reactions. Typically, acid-base reactions like \(a.\) and \(b.\) are exothermic, releasing energy.
• Reaction Conditions: Temperature, pressure, and concentration can affect the direction and speed of a reaction.

Understanding these elements is crucial for accurately balancing equations and predicting reaction behavior.

Ionic Compounds

Ionic compounds are formed when metals transfer electrons to non-metals, resulting in positive and negative ions. These ions often end up in aqueous solutions, dissociating into their respective charges. For instance:

• In reaction \(f.\), solid \(\text{LiH}\) reacts with water to form \(\text{LiOH}\), an ionic compound in solution.
• In reaction \(b.\), the combination of \(\text{Na}_2\text{CO}_3\) and \(\text{H}_2\text{SO}_4\) results in the production of sodium sulfate, another ionic compound.

Understanding how ionic compounds form and behave in solutions helps in visualizing reactions, particularly when writing net ionic equations. It allows chemists to focus only on the actual chemical changes, without the interference of unreactive spectator ions.

Solution Chemistry

Solution chemistry is an essential part of understanding reactions. It involves studying how substances dissolve, dissociate, and interact in solvent environments. In aqueous solutions:

• Dissolution: Ionic compounds dissolve and dissociate into ions, providing a medium for rapid reactions, seen in our reactions like \(a.\) and \(b.\).
• Concentration: The amount of solute affects reaction rates and equilibrium. More concentrated solutions tend to react more quickly.
• Dynamic Equilibria: Acids and bases reacting in solutions are typically reversible, seeking a balance between forward and backward reactions as shown in some weak acid-base interactions.

For students, mastering the principles of solution chemistry helps in predicting how substances behave in different conditions and understanding the broader implications of reactions in aqueous environments.