Question

Calculate the normality of each of the following solutions. a. \(0.50 M\) acetic acid, \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) b. \(0.00250 M\) sulfuric acid, \(\mathrm{H}_{2} \mathrm{SO}_{4}\) c. 0.10 \(M\) potassium hydroxide, KOH

Short answer

The normality of the given solutions are: a. \(0.50 M\) acetic acid, \(HC_{2}H_{3}O_{2}\): \(0.50 N\) b. \(0.00250 M\) sulfuric acid, \(H_{2}SO_{4}\): \(0.00500 N\) c. \(0.10 M\) potassium hydroxide, KOH: \(0.10 N\)

Step-by-step solution

a. 0.50 M acetic acid, HC2H3O2

Step 1: Identify the number of acidic hydrogen ions (protons) Acetic acid (HC2H3O2) has one acidic hydrogen ion in the reaction (H+). So, n = 1. Step 2: Calculate normality using the formula \(N = M \times n\) \(N_{acetic} = 0.50 \times 1\) \(N_{acetic} = 0.50 N\) The normality of the 0.50 M acetic acid solution is \(0.50 N\).

b. 0.00250 M sulfuric acid, H2SO4

Step 1: Identify the number of acidic hydrogen ions (protons) Sulfuric acid (H2SO4) has two acidic hydrogen ions in the reaction (2H+). So, n = 2. Step 2: Calculate normality using the formula \(N = M \times n\) \(N_{sulfuric} = 0.00250 \times 2\) \(N_{sulfuric} = 0.00500 N\) The normality of the 0.00250 M sulfuric acid solution is \(0.00500 N\).

c. 0.10 M potassium hydroxide, KOH

Step 1: Identify the number of basic hydrogen ions (protons) Potassium hydroxide (KOH) has one basic hydrogen ion in the reaction (OH-). So, n = 1. Step 2: Calculate normality using the formula \(N = M \times n\) \(N_{KOH} = 0.10 \times 1\) \(N_{KOH} = 0.10 N\) The normality of the 0.10 M potassium hydroxide solution is \(0.10 N\).

Key concepts

molarity

Molarity is a way to express concentration that tells us how many moles of a solute are present in one liter of solution. It's abbreviated as "M". For example, a solution with a molarity of 1 M has one mole of solute per liter of solution. Molarity helps chemists understand how much of a substance is in a given volume, which is crucial for creating solutions and reactions in the lab.

To calculate molarity, use the formula \[M = \frac{\text{moles of solute}}{\text{liters of solution}}.\]This formula emphasizes the importance of measuring the volume of the entire solution, not just the solute. When adjusting molarity, like diluting a solution, the number of moles of solute remains constant, as only the volume changes. Keeping track of these variables is vital for maintaining accurate chemical reactions.

acidic hydrogen ions

Acidic hydrogen ions, often represented as \(H^+\), play a key role in acid-base chemistry. They are protons released by acids when dissolved in water, which defines a substance as an acid. The strength of an acid is determined by how readily it dissociates to release these protons.

For example:

• Acetic acid \( (\mathrm{HC}_2\mathrm{H}_3\mathrm{O}_2)\) releases one \(H^+\) ion, making it a monoprotic acid.
• Sulfuric acid \((\mathrm{H}_2\mathrm{SO}_4)\) releases two \(H^+\) ions, categorizing it as a diprotic acid.

Understanding the number of acidic hydrogen ions helps in calculating normality and preparing solutions with specific reactive capacities.

basic hydrogen ions

Basic hydrogen ions actually refer to the hydroxide ions \( (OH^-) \) released by bases when dissolved in water. These ions make a solution basic or alkaline. For example, potassium hydroxide \((\mathrm{KOH})\) releases one \(OH^-\) ion, defining its basicity.

Here's how this works in acid-base reactions:

• Bases accept \(H^+\) ions by neutralizing them and forming water or other compounds.
• The more \(OH^-\) ions released, the stronger the base and its neutralizing power.

The count of basic hydrogen ions in a solution influences its normality, determining how much it can counteract an acid.

acid-base reactions

Acid-base reactions are chemical processes where acids donate protons \((H^+)\), and bases accept them. This process results in the formation of water and possibly salts. When these reactions reach completion, they neutralize the reactive properties of both acids and bases.

The reaction is expressed generally as:\[\text{Acid} + \text{Base} \rightarrow \text{Water} + \text{Salt}\]

• Acids lose \(H^+\) ions, which bases accept.
• The \(OH^-\) ions from bases combine with \(H^+\) ions to form water.
• The neutralization often results in a pH closer to 7.

These reactions are fundamental in many real-world applications, from biological systems to industrial processes, illustrating the balance within chemical equilibria.