Question

A typical solution used in general chemistry laboratories is 3.0 \(M\) HCl. Describe, in detail, the composition of \(2.0 \mathrm{L}\) of a 3.0-M HCI solution. How would 2.0 L of a \(3.0-M \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) solution differ from the same quantity of the HCl solution?

Short answer

In 2.0 L of a 3.0 M HCl solution, there are 6.0 moles of HCl solute and approximately 2.0 L of water as the solvent. In 2.0 L of a 3.0 M HC₂H₃O₂ solution, there are 6.0 moles of HC₂H₃O₂ solute and approximately 2.0 L of water as the solvent. The main difference between these two solutions is the solute, with HCl being a strong acid and HC₂H₃O₂ (acetic acid) being a weak acid. The HCl solution will have a lower pH, while the HC₂H₃O₂ solution will have a higher pH and fewer dissolved ions due to its partial ionization.

Step-by-step solution

Understanding Molarity

Molarity is a measure of the concentration of a solute in a solution, defined as the number of moles of solute per liter of solution. Molarity is commonly denoted as "M". In this problem, we are given two solutions with 3.0 M concentration: one with HCl as the solute and one with HC₂H₃O₂.

Analyzing the HCl solution

To find the composition of 2.0 L of a 3.0 M HCl solution, first, we need to calculate the moles of HCl using the molarity formula: Moles of solute = Molarity × Volume of solution In our case: Moles of HCl = 3.0 M × 2.0 L = 6.0 moles of HCl The composition of the 2.0 L HCl solution: - 6.0 moles of HCl as solute - Approximately 2.0 L of water as the solvent (assuming HCl is dissolved in water)

Analyzing the HC₂H₃O₂ solution

Now we'll analyze the 2.0 L of a 3.0 M HC₂H₃O₂ solution. Moles of HC₂H₃O₂ = 3.0 M × 2.0 L = 6.0 moles of HC₂H₃O₂ The composition of the 2.0 L HC₂H₃O₂ solution: - 6.0 moles of HC₂H₃O₂ as solute - Approximately 2.0 L of water as the solvent (assuming HC₂H₃O₂ is dissolved in water)

Comparing both solutions

When comparing the two solutions, we can see that both solutions have the same volume (2.0 L) and molarity (3.0 M). The main difference between these two solutions is the solute: - In the 2.0 L of 3.0 M HCl solution, the solute is Hydrochloric acid (HCl) which is a strong acid. - In the 2.0 L of 3.0 M HC₂H₃O₂ solution, the solute is Acetic acid (HC₂H₃O₂) which is a weak acid. The difference in the strength of the acids determines their degree of ionization in the solution, which in turn determines the pH values and other chemical properties of the solutions. In general, the HCl solution will have a lower pH (more acidic) and will have more dissolved ions due to its complete ionization, while the HC₂H₃O₂ solution will have a higher pH (less acidic) and will have fewer dissolved ions due to its partial ionization.

Key concepts

Solution Composition

In chemistry, a solution is a homogeneous mixture composed of two components: the solute and the solvent. The solute is the substance that is dissolved, while the solvent is the medium in which the solute is dissolved. In our example, we are dealing with two different solutes, Hydrochloric acid (HCl) and Acetic acid (HC₂H₃O₂), both forming solutions when mixed with water, which acts as the solvent.

To describe the composition of these solutions, we look at their molarity. Molarity (\(M\) ) is a crucial factor to consider. It expresses the concentration of the solute in a given volume of the solution. Specifically, it is the number of moles of the solute divided by the volume of the solution in liters. For example, a 3.0 M HCl solution contains 3 moles of HCl per liter. Therefore, in a 2.0 L solution of 3.0 M HCl, there are \(3.0 \times 2.0 = 6.0\) moles of HCl present.

Similarly, a 2.0 L solution of 3.0 M HC₂H₃O₂ also contains 6.0 moles of the acetic acid solute. While the water acts as a solvent, providing the volume for the solution's diluteness, it is crucial to note that both solutions, theoretically, have approximately 2.0 L of water as the solvent.

Strong vs Weak Acids

Acids are substances that can donate a proton (\(H^+\) ) when dissolved in water. The strength of an acid refers to its ability to donate these protons. This ability significantly affects the chemical properties and behavior of acid solutions.

Strong acids, like Hydrochloric acid (HCl), dissociate completely in water. This dissociation means almost all HCl molecules break apart into \(H^+\) and \(Cl^-\) ions in solution. Therefore, strong acid solutions usually have a higher concentration of \(H^+\) ions, resulting in a lower pH and a more acidic environment.

On the other hand, weak acids, such as Acetic acid (HC₂H₃O₂), do not completely dissociate in water. Only a small fraction of HC₂H₃O₂ molecules release \(H^+\) ions into the solution, with many of them remaining intact. This behavior suggests that weak acids result in fewer \(H^+\) ions in the solution, leading to a higher pH when compared with a strong acid of the same concentration.

Understanding the difference between strong and weak acids is essential for predicting how different solutions will react chemically and physically.

Acid Ionization

The concept of ionization describes the process by which an acid splits into its ions in a solution. For acids, especially, this involves the release of \(H^+\) ions. Acid ionization is a vital property used to classify acids into strong or weak categories.

In terms of ionization:

• Strong acids, such as HCl, exhibit complete ionization. This means that in a solution, almost all acid molecules dissociate into their respective ions. As a result, there are many free \(H^+\) ions, which contribute to the solution's acidic characteristics.
• Weak acids, like Acetic acid (HC₂H₃O₂), only partially ionize in solutions. A significant portion of the acid remains in its original, un-ionized form, which results in fewer free \(H^+\) ions in the solution, compared to a strong acid of the same molarity.

The degree of acid ionization affects several factors, including the acidity of the solution (pH) and its electrical conductivity. A greater ionization degree leads to a more robust acidic environment and increased electrical conductivity.