Question

Using the symbol HA to represent a general acid, write an equation showing how HA forms its conjugate Bronsted-Lowry base when dissolved in water.

Short answer

\( HA (aq) + H2O (l) \rightarrow A^- (aq) + H3O^+ (aq) \) In this equation, the general acid (HA) donates a proton (H+) to water (H2O), forming its conjugate Bronsted-Lowry base (A-) and a hydronium ion (H3O+).

Step-by-step solution

Identify the reactants

The reactants in the reaction are the general acid HA and water (H2O).

Identify the ion donor and the ion acceptor

In this reaction, the general acid HA is the proton (H+) donor, and water (H2O) is the proton (H+) acceptor.

Write the equation for the reaction

An acid donates a proton (H+) to water (H2O), forming its conjugate Bronsted-Lowry base (A-) and a hydronium ion (H3O+). The equation for the reaction is: HA (aq) + H2O (l) → A- (aq) + H3O+ (aq) In this equation: - HA represents the general acid. - H2O represents water. - A- is the conjugate Bronsted-Lowry base formed when HA donates a proton (H+). - H3O+ is the hydronium ion formed when water (H2O) accepts a proton (H+).

Key concepts

Acid-Base Reactions

When we talk about acid-base reactions, we are referring to the process in which an acid donates a proton to a base. According to the Bronsted-Lowry theory, an acid is a substance that can donate a proton, while a base is a substance that can accept a proton. This interaction between an acid and a base is fundamental in chemistry and is what we call a proton transfer reaction.

Let's illustrate this with a simple example. Imagine you have acetic acid, which is a common acid found in vinegar. When it reacts with water, a base in this context, it donates a proton to the water molecules. The acetic acid (CH3COOH) loses a proton and becomes its conjugate base (CH3COO-), while the water gains a proton and becomes a hydronium ion (H3O+). This process is essential in many chemical reactions and is at the heart of understanding acidity and basicity in substances.

Conjugate Base Formation

The formation of a conjugate base is a critical concept in understanding Bronsted-Lowry acid-base theory. During an acid-base reaction, when an acid loses a proton, it becomes what we call a conjugate base. It's called 'conjugate' because it is linked to the acid-base pair by the loss or gain of a single proton.

For instance, if we start with hydrochloric acid (HCl), when it donates a proton, it turns into its conjugate base, chloride (Cl-). The notion of conjugate base is crucial as it helps us predict the outcome of reactions and the relative strengths of acids and bases. Acids with weaker conjugate bases are generally considered to be stronger since they tend to donate protons more readily.

Proton Transfer

Proton transfer is the hallmark of acid-base chemistry. In every acid-base reaction, a proton (H+) is transferred from the acid to the base. This is a very specific type of chemical reaction where the only thing that changes is the ownership of a single hydrogen nucleus without its electron - essentially, a proton.

For example, when ammonia (NH3), a base, comes into contact with hydrogen chloride (HCl), an acid, the hydrogen proton from HCl is transferred to ammonia. This transfer forms ammonium (NH4+), a positively charged ion, and the chloride ion (Cl-), which is the conjugate base of the acid. Remember, whenever we look at an acid-base reaction, we're looking for this key transfer of a proton.

Hydronium Ion

The hydronium ion (H3O+) is often produced during acid-base reactions involving water. It's essentially a water molecule (H2O) that has accepted an extra proton, becoming positively charged. For instance, if we take the reaction of hydrochloric acid (HCl) in water, the HCl donates a proton to water, resulting in the formation of hydronium ions and chloride ions (Cl-).

The presence of hydronium ions is what makes a solution acidic. The concentration of these ions in a solution is measured by the pH scale, which is a measure of acidity and basicity. The more hydronium ions present, the lower the pH, and the more acidic the solution is. Understanding the role of the hydronium ion is essential for grasping the concepts of pH and the acid-base properties of substances in solution.