Problem 95 - Chemistry: The Central Science in SI Units, Global Edition

Question

Antacids are often used to relieve pain and promote healing in the treatment of mild ulcers. Write balanced net ionic equations for the reactions between the aqueous \(\mathrm{HCl}\) in the stomach and each of the following substances used in various antacids: (a) \(\mathrm{Al}(\mathrm{OH})_{3}(s),(\mathbf{b}) \mathrm{Mg}(\mathrm{OH})_{2}(s),(\mathbf{c}) \mathrm{MgCO}_{3}(s),\) (d) \(\mathrm{NaAl}\left(\mathrm{CO}_{3}\right)(\mathrm{OH})_{2}(s)\) (e) \(\mathrm{CaCO}_{3}(s)\)

Step-by-Step Solution

Answer

(a) Al(OH)3(s) + 3H+(aq) → 3H2O(l) (b) Mg(OH)2(s) + 2H+(aq) → 2H2O(l) (c) CO3^2-(aq) + 2H+(aq) → H2O(l) + CO2(g) (d) CO3^2-(aq) + 2OH-(aq) + 6H+(aq) → 2H2O(l) + 3CO2(g) + Na+(aq) (e) CO3^2-(aq) + 2H+(aq) → H2O(l) + CO2(g)

Step 1: (a) Reaction between HCl and Al(OH)3

First, we will write the balanced molecular equation for the reaction between Al(OH)3 and HCl: Al(OH)3(s) + 3HCl(aq) → AlCl3(aq) + 3H2O(l) To write the net ionic equation, let's break down the molecular equation into its ionic components: Al(OH)3(s) → Al^3+ (aq) + 3OH- (aq) 3HCl(aq) → 3H+ (aq) + 3Cl- (aq) AlCl3(aq) → Al^3+ (aq) + 3Cl- (aq) Since Al^3+(aq) and 3Cl-(aq) are the same on both sides of the equation, they are called spectator ions and can be eliminated from the net ionic equation. The net ionic equation for this reaction will be: Al(OH)3(s) + 3H+ (aq) → 3H2O(l)

Step 2: (b) Reaction between HCl and Mg(OH)2

First, we will write the balanced molecular equation for the reaction between Mg(OH)2 and HCl: Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + 2H2O(l) Now, let's break down the molecular equation into its ionic components: Mg(OH)2(s) → Mg^2+ (aq) + 2OH- (aq) 2HCl(aq) → 2H+ (aq) + 2Cl- (aq) MgCl2(aq) → Mg^2+ (aq) + 2Cl- (aq) The spectator ions in this case are Mg^2+(aq) and 2Cl-(aq). The net ionic equation for this reaction will be: Mg(OH)2(s) + 2H+ (aq) → 2H2O(l)

Step 3: (c) Reaction between HCl and MgCO3

First, we will write the balanced molecular equation for the reaction between MgCO3 and HCl: MgCO3(s) + 2HCl(aq) → MgCl2(aq) + H2O(l) + CO2(g) Now, let's break down the molecular equation into its ionic components: MgCO3(s) → Mg^2+ (aq) + CO3^2- (aq) 2HCl(aq) → 2H+ (aq) + 2Cl- (aq) MgCl2(aq) → Mg^2+ (aq) + 2Cl- (aq) The spectator ions in this case are Mg^2+(aq) and 2Cl-(aq). The net ionic equation for this reaction will be: CO3^2- (aq) + 2H+ (aq) → H2O(l) + CO2(g)

Step 4: (d) Reaction between HCl and NaAl(CO3)(OH)2

First, we will write the balanced molecular equation for the reaction between NaAl(CO3)(OH)2 and HCl: NaAl(CO3)(OH)2(s) + 6HCl(aq) → AlCl3(aq) + NaCl(aq) + 2H2O(l) + 3CO2(g) Now, let's break down the molecular equation into its ionic components: NaAl(CO3)(OH)2(s) → Na+ (aq) + Al^3+ (aq) + CO3^2- (aq) + 2OH- (aq) 6HCl(aq) → 6H+ (aq) + 6Cl- (aq) AlCl3(aq) → Al^3+ (aq) + 3Cl- (aq) NaCl(aq) → Na+ (aq) + Cl- (aq) The spectator ions in this case are Al^3+(aq) and Cl-(aq). The net ionic equation for this reaction will be: CO3^2- (aq) + 2OH- (aq) + 6H+ (aq) → 2H2O(l) + 3CO2(g) + Na+ (aq)

Step 5: (e) Reaction between HCl and CaCO3

First, we will write the balanced molecular equation for the reaction between CaCO3 and HCl: CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g) Now, let's break down the molecular equation into its ionic components: CaCO3(s) → Ca^2+ (aq) + CO3^2- (aq) 2HCl(aq) → 2H+ (aq) + 2Cl- (aq) CaCl2(aq) → Ca^2+ (aq) + 2Cl- (aq) The spectator ions in this case are Ca^2+(aq) and 2Cl-(aq). The net ionic equation for this reaction will be: CO3^2- (aq) + 2H+ (aq) → H2O(l) + CO2(g)

Key Concepts

Acid-Base Reactions

In acid-base reactions, acids react with bases to produce water and a salt. A good example of this is when hydrochloric acid (HCl), an acid present in the stomach, reacts with bases such as Mg(OH)_2 or Al(OH)_3. These reactions help neutralize excess stomach acid.

• Acid: Substances that increase the concentration of H⁺ ions in a solution.
• Base: Substances that increase the concentration of OH⁻ ions.

To identify the net ionic equation, it's crucial to focus on the ions that directly participate in forming water, leaving out spectator ions that do not take part in the reaction. When Al(OH)_3 reacts with HCl, the net ionic equation is:\[\mathrm{Al(OH)_3 (s) + 3H^+ (aq) \to 3H_2O (l)}\]Here, OH⁻ ions combine with H⁺ ions to form H_2O, demonstrating a fundamental principle of acid-base chemistry.

Antacids

Antacids are medications that help neutralize stomach acid, providing relief from symptoms like indigestion and heartburn. These agents often contain compounds like Al(OH)_3, Mg(OH)_2, or CaCO_3. They work by buffering the acidic environment of the stomach.
The role of antacids is to bring balance by reducing the acidity level. They undergo reactions that increase the pH of the stomach contents, making it less acidic. For instance:

• \( \mathrm{Mg(OH)_2 (s) + 2H^+ (aq) \to 2H_2O (l)} \)
• \( \mathrm{CaCO_3 (s) + 2H^+ (aq) \to H_2O (l) + CO_2 (g)} \)

In these reactions, the OH⁻ from the bases neutralizes the H⁺ ions in the stomach acid, forming water. The reaction with calcium carbonate (CaCO_3) also releases carbon dioxide gas, which can sometimes cause belching.

Chemical Equilibrium

Chemical equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction, resulting in a stable state where the concentrations of reactants and products remain constant.
In the context of acid-base reactions with antacids, reaching equilibrium means the antacid has effectively neutralized the stomach acid to where any additional acid doesn’t further react or change pH levels significantly. This concept helps to understand why only a specific amount of an antacid is needed to reach therapeutic effects. Dynamic Equilibrium in Reactions:

• The rates of the forward and backward reactions become equal.
• Concentrations of reactants and products stay constant but not necessarily equal.

For antacid reactions, once enough base has reacted to neutralize the acid, the system reaches an equilibrium where H⁺ ions are continuously formed and neutralized at the same rate, maintaining a balanced pH. Understanding equilibrium is key to correctly dosing antacids and ensuring effective relief without overuse.