Net Ionic Equation For Hydrolysis Of Nac2h3o2

The hydrolysis of sodium acetate (NaC₂H₃O₂) is a classic example of how salts derived from weak acids and strong bases can affect the pH of a solution. Understanding the net ionic equation for this process is crucial for grasping acid-base chemistry principles and predicting the behavior of various salt solutions.

Understanding Sodium Acetate and Hydrolysis

Sodium acetate (NaC₂H₃O₂) is a salt formed by the reaction of a strong base, sodium hydroxide (NaOH), and a weak acid, acetic acid (HC₂H₃O₂). When sodium acetate is dissolved in water, it dissociates completely into its ions:

NaC₂H₃O₂(s) → Na⁺(aq) + C₂H₃O₂⁻(aq)

The sodium ions (Na⁺) do not undergo hydrolysis because they are the conjugate acid of a strong base (NaOH). However, the acetate ions (C₂H₃O₂⁻) are the conjugate base of a weak acid (acetic acid) and therefore react with water in a process called hydrolysis. This reaction changes the pH of the solution, making it basic.

What is Hydrolysis?

Hydrolysis, in the context of salts, is the reaction of an ion with water, leading to the formation of either hydroxide ions (OH⁻) or hydronium ions (H₃O⁺), thereby affecting the pH of the solution. In the case of sodium acetate, the acetate ion reacts with water to produce acetic acid and hydroxide ions.

The Role of Weak Acids and Strong Bases

Salts formed from weak acids and strong bases undergo hydrolysis because the conjugate base of the weak acid has a significant affinity for protons (H⁺). This affinity allows the conjugate base to react with water, pulling protons from H₂O molecules and generating hydroxide ions.

The Net Ionic Equation for Hydrolysis of Sodium Acetate

To write the net ionic equation for the hydrolysis of sodium acetate, we focus only on the species that are directly involved in the reaction. The sodium ion (Na⁺) is a spectator ion, meaning it does not participate in the reaction and is present on both sides of the equation. Therefore, we can exclude it from the net ionic equation.

The acetate ion (C₂H₃O₂⁻) reacts with water (H₂O) to form acetic acid (HC₂H₃O₂) and hydroxide ions (OH⁻). The balanced net ionic equation for this reaction is:

C₂H₃O₂⁻(aq) + H₂O(l) ⇌ HC₂H₃O₂(aq) + OH⁻(aq)

This equation shows that the acetate ion acts as a base, accepting a proton from water and producing hydroxide ions, which increase the pH of the solution. The double arrow (⇌) indicates that this is an equilibrium reaction, meaning it proceeds in both directions.

Step-by-Step Breakdown

1. Dissociation of Sodium Acetate: NaC₂H₃O₂(s) → Na⁺(aq) + C₂H₃O₂⁻(aq)

2. Identification of Reactive Ion: The acetate ion (C₂H₃O₂⁻) is the reactive ion because it is the conjugate base of a weak acid.

3. Reaction with Water: C₂H₃O₂⁻(aq) + H₂O(l) ⇌ HC₂H₃O₂(aq) + OH⁻(aq)

4. Exclusion of Spectator Ions: The sodium ion (Na⁺) is a spectator ion and is not included in the net ionic equation.

Explanation of the Equation

• C₂H₃O₂⁻(aq): This represents the acetate ion in aqueous solution.
• H₂O(l): This represents liquid water.
• HC₂H₃O₂(aq): This represents acetic acid in aqueous solution, formed when the acetate ion accepts a proton from water.
• OH⁻(aq): This represents the hydroxide ion in aqueous solution, which is responsible for the basic pH of the solution.

Factors Affecting Hydrolysis

Several factors can influence the extent of hydrolysis and the resulting pH of the solution. These include:

1. The Strength of the Acid or Base

The weaker the acid from which the salt is derived, the stronger its conjugate base, and the greater the extent of hydrolysis. In the case of sodium acetate, acetic acid is a weak acid, so its conjugate base, the acetate ion, is relatively strong and undergoes significant hydrolysis.

2. Concentration of the Salt

The concentration of the salt affects the number of ions available to react with water. Higher concentrations of sodium acetate will result in more acetate ions in solution, leading to a greater production of hydroxide ions and a higher pH.

3. Temperature

Temperature can influence the equilibrium constant for the hydrolysis reaction. Generally, higher temperatures favor the hydrolysis reaction, increasing the concentration of hydroxide ions and raising the pH.

4. Presence of Other Ions

The presence of other ions in the solution can also affect the hydrolysis of sodium acetate. For example, the addition of a strong acid will neutralize the hydroxide ions, shifting the equilibrium of the hydrolysis reaction to the left and decreasing the pH.

Calculating the pH of a Sodium Acetate Solution

To calculate the pH of a sodium acetate solution, you need to consider the equilibrium constant for the hydrolysis reaction, known as the base dissociation constant (Kb). The Kb for the acetate ion can be calculated from the acid dissociation constant (Ka) of acetic acid using the following relationship:

Kw = Ka * Kb

Where Kw is the ion product of water (1.0 x 10⁻¹⁴ at 25°C). The Ka of acetic acid is approximately 1.8 x 10⁻⁵. Therefore, the Kb for the acetate ion can be calculated as:

Kb = Kw / Ka = (1.0 x 10⁻¹⁴) / (1.8 x 10⁻⁵) ≈ 5.6 x 10⁻¹⁰

Using the Kb value, you can set up an ICE (Initial, Change, Equilibrium) table to determine the equilibrium concentrations of the species involved in the hydrolysis reaction. Let's consider a solution of 0.1 M sodium acetate:

The equilibrium expression for the hydrolysis reaction is:

Kb = [HC₂H₃O₂][OH⁻] / [C₂H₃O₂⁻]

Substituting the equilibrium concentrations from the ICE table:

5. 6 x 10⁻¹⁰ = (x)(x) / (0.1 - x)

Since Kb is very small, we can assume that x is much smaller than 0.1, so we can simplify the equation to:

5. 6 x 10⁻¹⁰ = x² / 0.1

Solving for x:

x² = (5.6 x 10⁻¹⁰) * 0.1 = 5.6 x 10⁻¹¹ x = √(5.6 x 10⁻¹¹) ≈ 7.48 x 10⁻⁶ M

Therefore, [OH⁻] = 7.48 x 10⁻⁶ M. Now we can calculate the pOH:

pOH = -log[OH⁻] = -log(7.48 x 10⁻⁶) ≈ 5.13

Finally, we can calculate the pH:

pH = 14 - pOH = 14 - 5.13 ≈ 8.87

So, a 0.1 M solution of sodium acetate has a pH of approximately 8.87, indicating that it is a basic solution.

Applications of Sodium Acetate Hydrolysis

The hydrolysis of sodium acetate has several practical applications in various fields:

1. Buffers

Sodium acetate is often used in combination with acetic acid to create buffer solutions. A buffer solution resists changes in pH upon the addition of small amounts of acid or base. The acetate ion can neutralize added acid, while acetic acid can neutralize added base, maintaining a relatively stable pH.

2. Chemical Synthesis

Sodium acetate is used as a base in various organic reactions. Its ability to generate hydroxide ions through hydrolysis makes it a mild base suitable for reactions where strong bases would cause unwanted side reactions.

3. Food Industry

Sodium acetate is used as a food additive to regulate acidity and as a preservative. Its buffering capacity helps maintain the pH of food products, preventing spoilage and preserving their quality.

4. Textile Industry

In the textile industry, sodium acetate is used as a mordant, helping dyes adhere to fabrics. It also acts as a buffering agent in dyeing processes, ensuring consistent and even coloration.

5. Laboratory Applications

Sodium acetate is commonly used in laboratories for various purposes, including DNA extraction, protein purification, and as a component of electrophoresis buffers. Its buffering properties are essential for maintaining the integrity of biological molecules during these procedures.

Common Mistakes to Avoid

When working with hydrolysis reactions, it's important to avoid common mistakes that can lead to incorrect predictions or calculations. Some of these mistakes include:

1. Forgetting to Exclude Spectator Ions

Spectator ions do not participate in the reaction and should not be included in the net ionic equation. Including them can lead to confusion and incorrect interpretations of the reaction.

2. Incorrectly Identifying Weak Acids and Bases

It's crucial to correctly identify whether a salt is derived from a weak acid or a weak base. Only salts derived from weak acids or bases undergo significant hydrolysis.

3. Neglecting the Equilibrium Nature of the Reaction

Hydrolysis reactions are equilibrium reactions, and the extent of hydrolysis depends on the equilibrium constant (Kb or Ka). Neglecting this can lead to inaccurate calculations of pH.

4. Assuming Complete Hydrolysis

Hydrolysis reactions do not proceed to completion. The extent of hydrolysis is determined by the strength of the conjugate acid or base and the reaction's equilibrium constant.

5. Incorrectly Applying the Kw = Ka * Kb Relationship

Ensure that you correctly use the relationship Kw = Ka * Kb to calculate the Kb for the conjugate base or the Ka for the conjugate acid. Using the wrong value can lead to significant errors in pH calculations.

Examples of Other Hydrolysis Reactions

Besides sodium acetate, several other salts undergo hydrolysis in water, affecting the pH of the solution. Here are a few examples:

1. Ammonium Chloride (NH₄Cl)

Ammonium chloride is a salt formed from a weak base (ammonia, NH₃) and a strong acid (hydrochloric acid, HCl). When dissolved in water, the ammonium ion (NH₄⁺) acts as an acid, donating a proton to water and forming hydronium ions (H₃O⁺):

NH₄⁺(aq) + H₂O(l) ⇌ NH₃(aq) + H₃O⁺(aq)

This reaction lowers the pH of the solution, making it acidic.

2. Sodium Carbonate (Na₂CO₃)

Sodium carbonate is a salt formed from a strong base (sodium hydroxide, NaOH) and a weak acid (carbonic acid, H₂CO₃). The carbonate ion (CO₃²⁻) acts as a base, accepting protons from water and forming hydroxide ions (OH⁻):

CO₃²⁻(aq) + H₂O(l) ⇌ HCO₃⁻(aq) + OH⁻(aq)

This reaction raises the pH of the solution, making it basic.

3. Aluminum Chloride (AlCl₃)

Aluminum chloride is a salt formed from a weak base (aluminum hydroxide, Al(OH)₃) and a strong acid (hydrochloric acid, HCl). The aluminum ion (Al³⁺) undergoes hydrolysis, forming hydronium ions (H₃O⁺):

Al³⁺(aq) + 3H₂O(l) ⇌ Al(OH)₃(s) + 3H⁺(aq)

This reaction lowers the pH of the solution, making it acidic.

Conclusion

The hydrolysis of sodium acetate is a fundamental concept in acid-base chemistry, illustrating how salts derived from weak acids and strong bases can affect the pH of a solution. The net ionic equation for this process, C₂H₃O₂⁻(aq) + H₂O(l) ⇌ HC₂H₃O₂(aq) + OH⁻(aq), highlights the reaction between the acetate ion and water, leading to the formation of acetic acid and hydroxide ions. Understanding the factors that influence hydrolysis and being able to calculate the pH of sodium acetate solutions are essential skills for anyone studying chemistry or related fields. Furthermore, recognizing the applications of sodium acetate hydrolysis in various industries underscores the practical importance of this chemical principle. By avoiding common mistakes and expanding your knowledge to other hydrolysis reactions, you can gain a deeper appreciation for the complexities and applications of acid-base chemistry.