Carbon dioxide in the atmosphere dissolves in raindrops to produce carbonic acid (H2CO3), causing the pH of clean, unpolluted rain to range from about 5.2 to 5.6. What are the ranges of [H+] and [OH-] in the raindrops?

Verified step by step guidance

Step 1: Understand the relationship between pH and [H+]. The pH of a solution is defined as the negative logarithm (base 10) of the hydrogen ion concentration: \( \text{pH} = -\log[H^+] \).

Step 2: Calculate the range of [H+] using the given pH range. For pH 5.2, use the formula \( [H^+] = 10^{-\text{pH}} \) to find the concentration of hydrogen ions. Repeat this calculation for pH 5.6.

Step 3: Understand the relationship between [H+] and [OH-] using the water dissociation constant. At 25°C, the product of the concentrations of hydrogen ions and hydroxide ions in water is constant: \( [H^+][OH^-] = 1.0 \times 10^{-14} \).

Step 4: Calculate the range of [OH-] using the [H+] values obtained in Step 2. Rearrange the equation from Step 3 to solve for [OH-]: \( [OH^-] = \frac{1.0 \times 10^{-14}}{[H^+]} \).

Step 5: Verify the consistency of your results. Ensure that the calculated [H+] and [OH-] values are consistent with the pH range and the water dissociation constant.

Key Concepts

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

pH Scale

The pH scale measures the acidity or basicity of a solution, ranging from 0 to 14. A pH of 7 is neutral, while values below 7 indicate acidity and above 7 indicate basicity. The pH is inversely related to the concentration of hydrogen ions [H+]; as [H+] increases, pH decreases, indicating a more acidic solution.

Ion Product of Water

The ion product of water (Kw) is the equilibrium constant for the self-ionization of water, defined as Kw = [H+][OH-] = 1.0 x 10^-14 at 25°C. This relationship allows us to calculate the concentration of hydroxide ions [OH-] when the concentration of hydrogen ions [H+] is known, and vice versa, maintaining the balance in aqueous solutions.

Acid-Base Equilibrium

In an aqueous solution, the presence of acids and bases establishes an equilibrium between [H+] and [OH-]. For a solution with a known pH, the concentration of [H+] can be calculated using the formula [H+] = 10^(-pH). Consequently, the concentration of [OH-] can be derived from the ion product of water, ensuring that the product of [H+] and [OH-] remains constant.

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The average pH of normal arterial blood is 7.40. At normal body temperature 137 °C2, Kw = 2.4 * 10-14. Calculate 3H+4, 3OH-4, and pOH for blood at this temperature.

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Addition of the indicator methyl orange to an unknown solution leads to a yellow color. The addition of bromthymol blue to the same solution also leads to a yellow color. (b) What is the range (in whole numbers) of possible pH values for the solution?

Textbook Question

Addition of phenolphthalein to an unknown colorless solution does not cause a color change. The addition of bromthymol blue to the same solution leads to a yellow color. (b) Which of the following can you establish about the solution: (i) A minimum pH, (ii) A maximum pH, or (iii) A specific range of pH values?

Textbook Question

Addition of phenolphthalein to an unknown colorless solution does not cause a color change. The addition of bromthymol blue to the same solution leads to a yellow color. (c) What other indicator or indicators would you want to use to determine the pH of the solution more precisely?