How does a mixture of CH3COOH and CH3COONa act as a buffer when a small amount of NaOH is added?

Do the properties of CH3OOH behave as acidic or basic in a buffer solution?  This article explains in detail the buffer action with NaOH and how it happens.

Introduction:

A fundamental concept related to the topic of this article is Le Chatelier's principle which states that “If a dynamic equilibrium is altered by altering the conditions in a reaction, the position of equilibrium also adjusts itself and moves to counteract the change”. These altered conditions can be many things involved in a reaction including pressure, concentration, temperature, etc. Considering the scope of this article, only change in the concentration is taken into account.

According to this law, if the reaction system between two elements is changed in a way that increases the concentration of one of the reacting species involved in this system, then the reaction will tend to favor the reaction in which that particular species is consumed. This theory is extremely important when understanding buffers that obey the rules defined in this theory.

Is CH₃OOH Acid or Base?

Acetic acid CH₃COOH or C₂H₄O₂ is a weak simple monocarboxylic organic acid containing two carbons in its structure. Also goes by the name ethanoic acid, is around 4 percent by concentration in the vinegar found in almost every kitchen. It is produced from the oxidation of ethanol and is a classic example of a buffer solution when with salt. It can act as a weak monoprotic acid in a buffer solution and a salt act as its conjugate base of acetate anions CH₃COO^-.

Acetic acid experiences ionization during which its hydrogen center in the carboxyl group (−COOH) separates from the molecule and gives away a proton H⁺, thus, giving it acidic properties. So, for the question “Is CH₃OOH acid or base”, this equation can be referred to as an answer:

CH₃COOH + H → CH₃CO⁻²

The Buffer action of NaOH:

The acetic acid and NaOH react to produce sodium acetate CH₃COONa. Introducing more NaOH only results in the creation of more CH₃COOH, and the pH of the solution does not change much. The reaction can be stated as:

CH₃COOH   → CH₃COO^- + Na⁺

The acetate ion, CH3COO-, is produced when sodium acetate fully ionizes. This is where Le Chatelier's theory is applied to explain the buffering action. The addition of OH from NAOH will cause the equilibrium to move to the right, producing more H⁺, more CH₃COO^-, and a decrease in the concentration of CH₃COOH when it reacts with H⁺ to form water.

The concentration of CH₃COO^- ion rises forcing the backward reaction to take place. This further leads to the formation of CH3-COO- ion and H+ ion to form undissociated acetic acid in the solution. The accumulated high concentration of CH3COO- ions from the added NaCH3COO restricts the acetic acid from ionizing. This can be again explained by Le Chatelier's Principle as a weak acid, in this case, which is acetic acid, experiences even lower ionization due to the presence of the acetate ions.

CH₃COOH   <=> H⁺ + CH₃COO^-  (Ka = 1.67x10^-5)

As a result, the pH will remain reasonably steady and the H+ concentration will not significantly change hence acting as a buffer. Note in the above reaction that Na+ is not important to the buffer in this solution and is considered as a spectator ion whereas the main players here to control the pH of the solution are CH₃COOH and the CH₃COO^- ions.

A similar trend of behaving like a buffer can also be observed when it comes to considering concentration levels for a buffer observed before and after the reaction. To understand this, consider another example of a reaction between HNO₃ and CH₃COONa. The strong acid in this reaction is taken in 10 millimoles with 20 millimoles of a salt of weak acid. This results in a solution that will contain 10 millimoles of acetic acid and 10 millimoles of acetate ions, essentially resulting in making a buffer.

Wrapping Up:

In this article, we explained the basics of buffers, the question of the CH₃OOH acid or base nature when reacting with NaOH, and related concepts including Le Chatelier's principle which explains the working of buffer actions in this particular solution.

From the above explanation, we can conclude that a buffer can exist in two configurations: an excessive quantity of unionized acid that can be used to neutralize a base that can be added to the system. The second one can be an excessive quantity of conjugate base that can be used to neutralize any acid that is introduced to the system. From the first example explained above, CH₃COOH is the base used to neutralize and CH₃COO^- was the acid introduced in the system for buffer action.