Question

Write the formulas for three combinations of weak acid and salt that would act as buffered solutions. For each of your combinations, write chemical equations showing how the components of the buffered solution would consume added acid and base.

Short answer

Three combinations of weak acid and salt that would act as buffered solutions are: 1. Acetic acid (CH3COOH) and sodium acetate (CH3COONa) with the following reactions: a) Added Acid: \(CH3COO^- + H^+ \to CH3COOH\) b) Added Base: \(CH3COOH + OH^- \to CH3COO^- + H2O\) 2. Ammonium chloride (NH4Cl) and ammonia (NH3) with the following reactions: a) Added Acid: \(NH3 + H^+ \to NH4^+\) b) Added Base: \(NH4^+ + OH^- \to NH3 + H2O\) 3. Sodium dihydrogen phosphate (NaH2PO4) and sodium hydrogen phosphate (Na2HPO4) with the following reactions: a) Added Acid: \(HPO4^{2-} + H^+ \to H2PO4^-\) b) Added Base: \(H2PO4^- + OH^- \to HPO4^{2-} + H2O\)

Step-by-step solution

1. Combination of Acetic Acid and Sodium Acetate

One example of a buffered solution is a combination of a weak acid, acetic acid (CH3COOH), and its salt, sodium acetate (CH3COONa).

1a. Added Acid Reaction

When an acid (H+) is added to the buffered solution, the acetate ion from sodium acetate (CH3COO-) will consume the acid by forming acetic acid (CH3COOH). The reaction is: \(CH3COO^- + H^+ \to CH3COOH\)

1b. Added Base Reaction

If a base (OH-) is added to the buffered solution, the acetic acid (CH3COOH) will consume the base by forming water and an acetate ion. The reaction is: \(CH3COOH + OH^- \to CH3COO^- + H2O\)

2. Combination of Ammonium Chloride and Ammonia

Another example of a buffered solution is a combination of a weak acid, ammonium chloride (NH4Cl), and a weak base, ammonia (NH3).

2a. Added Acid Reaction

When an acid (H+) is added to the buffered solution, the ammonia (NH3) will consume the acid by forming ammonium ions. The reaction is: \(NH3 + H^+ \to NH4^+\)

2b. Added Base Reaction

If a base (OH-) is added to the buffered solution, the ammonium ion from ammonium chloride (NH4+) will consume the base by forming ammonia (NH3) and water. The reaction is: \(NH4^+ + OH^- \to NH3 + H2O\)

3. Combination of Sodium Dihydrogen Phosphate and Sodium Hydrogen Phosphate

A third example of a buffered solution is a combination of a weak acid, sodium dihydrogen phosphate (NaH2PO4), and its salt, sodium hydrogen phosphate (Na2HPO4).

3a. Added Acid Reaction

When an acid (H+) is added to the buffered solution, the hydrogen phosphate ion from sodium hydrogen phosphate (HPO4^2-) will consume the acid by forming dihydrogen phosphate ion (H2PO4^-). The reaction is: \(HPO4^{2-} + H^+ \to H2PO4^-\)

3b. Added Base Reaction

If a base (OH-) is added to the buffered solution, the dihydrogen phosphate ion from sodium dihydrogen phosphate (H2PO4^-) will consume the base by forming hydrogen phosphate ion (HPO4^2-) and water. The reaction is: \(H2PO4^- + OH^- \to HPO4^{2-} + H2O\)

Key concepts

Weak Acid and Salt Combinations

Buffered solutions are pivotal in chemistry for maintaining a stable pH environment. A common way to create a buffer is to combine a weak acid with a salt of its conjugate base. This pairing is crucial because the weak acid can donate protons (H⁺ ions) when the solution becomes too basic, while the conjugate base can absorb protons when the solution becomes too acidic.

For instance, acetic acid (CH₃COOH) combined with sodium acetate (CH₃COONa) forms an effective buffer. Acetic acid, being a weak acid, only partially dissociates in water, leaving a mixture of acetic acid and acetate ions. The sodium acetate dissociates completely, increasing the concentration of acetate ions. Together, these prevent significant pH changes when small amounts of acids or bases are added to the solution.

Similarly, the combination of a weak base like ammonia (NH₃) with its salt, ammonium chloride (NH₄Cl), forms another buffer. Ammonia reacts with H⁺ to form ammonium ions (NH₄⁺), adjusting the pH when required. These simple combinations indicate how buffers stabilize pH by harnessing the properties of weak acids or bases and their corresponding salts.

Chemical Equations for Buffers

Writing chemical equations for buffered solutions helps demonstrate how they mitigate changes in pH. Chemical equations represent the specific reactions that occur when acids or bases are added to a buffer.

A case to examine is the aforementioned acetic acid and sodium acetate buffer. When an acid is introduced, the equation is:

Acetate Buffer Reacts with Acid:

\(CH₃COO^- + H⁺ \to CH₃COOH\)

This reaction illustrates how the excess H⁺ ions are consumed by the acetate ions to form more acetic acid. Conversely, when a base is introduced, such as OH^-, acetic acid donates a proton, neutralizing the base and forming water, as shown below:

Acetate Buffer Reacts with Base:

\(CH₃COOH + OH^- \to CH₃COO^- + H₂O\)

The chemical equations for other buffer systems like the ammonium chloride-ammonia buffer or the sodium dihydrogen phosphate-sodium hydrogen phosphate buffer showcase the same principle: the absorption or release of H⁺ to resist changes in pH.

Acid and Base Reaction in Buffers

The beauty of buffer systems lies in their ability to react with added acids or bases, thus preventing drastic pH changes. When a buffer encounters an acid, H⁺ ions are added to the solution, but the buffer's base component reacts with these ions, assimilating them and minimizing their impact on pH.

For example, when H⁺ is added to a solution containing sodium hydrogen phosphate (Na₂HPO₄), it reacts with the hydrogen phosphate ions (HPO₄^2-), forming dihydrogen phosphate (H₂PO₄^-). This reaction can be represented as:

Hydrogen Phosphate Reacts with Acid:

\(HPO₄^2- + H⁺ \to H₂PO₄^-\)

If the situation is reversed and a base is introduced, then the dihydrogen phosphate ion can release a H⁺, neutralizing the added OH^- and maintaining the solution's pH:

Dihydrogen Phosphate Reacts with Base:

\(H₂PO₄^- + OH^- \to HPO₄^2- + H₂O\)

Through these reactions, buffers are capable of 'absorbing' the effects of added acids and bases, effectively 'cushioning' the solution against pH shifts and ensuring that the environment remains conducive for chemical reactions or biological processes that are sensitive to pH changes.