An aqueous solution contains 1.2 mM of total ions. (a) If the solution is NaCl(aq), what is the concentration of chloride ion?

Verified step by step guidance

Step 1: Understand the dissociation of NaCl in water. NaCl dissociates into Na^+ and Cl^- ions in a 1:1 ratio.

Step 2: Recognize that the total ion concentration is the sum of the concentrations of Na^+ and Cl^- ions.

Step 3: Set up the equation for total ion concentration: [Na^+] + [Cl^-] = 1.2 \(\text{ mM}\).

Step 4: Since NaCl dissociates in a 1:1 ratio, [Na^+] = [Cl^-].

Step 5: Substitute [Cl^-] for [Na^+] in the total ion concentration equation and solve for [Cl^-].

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Molarity and Concentration

Molarity is a measure of concentration defined as the number of moles of solute per liter of solution. In this context, a concentration of 1.2 mM (millimolar) indicates that there are 1.2 millimoles of total ions in one liter of the solution. Understanding molarity is essential for calculating the concentration of individual ions in a solution.

Dissociation of Ionic Compounds

Ionic compounds like NaCl dissociate into their constituent ions when dissolved in water. For NaCl, it separates into sodium ions (Na⁺) and chloride ions (Cl⁻). Therefore, the total concentration of ions in the solution is the sum of the concentrations of these individual ions, which is crucial for determining the concentration of chloride ions in this scenario.

Stoichiometry of Dissociation

The stoichiometry of dissociation refers to the ratio in which ions are produced from a compound when it dissolves. For NaCl, one mole of NaCl produces one mole of Na⁺ and one mole of Cl⁻. Thus, if the total ion concentration is 1.2 mM, the concentration of each ion (Na⁺ and Cl⁻) will be equal, allowing for straightforward calculations of individual ion concentrations.

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You are titrating an acidic solution with a basic one, and just realized you forgot to add the indicator that tells you when the equivalence point is reached. In this titration, the indicator turns blue at the equivalence point from an initially colorless solution. You quickly grab a bottle of indicator and add some to your titration beaker, and the whole solution turns dark blue. What do you do now?

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Textbook Question

Which data set, of the two graphed here, would you expect to observe from a titration like that shown in Figure 4.18?

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An aqueous solution contains 1.2 mM of total ions. (b) If the solution is FeCl3(aq), what is the concentration of chloride ion?

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