Chapter 7: Problem 110
The \(\mathrm{p} K_{\mathrm{a}}\) of certain weak acid is \(4.0 .\) What should be the salt to acid ratio if we have to prepare a buffer with \(\mathrm{pH}=5\) using the acid and of its salts? (1) \(4: 5\) (2) \(5: 4\) (3) \(10: 1\) (4) \(1: 10\)
Short Answer
Expert verified
The salt to acid ratio is 10:1, which is option (3).
Step by step solution
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Henderson-Hasselbalch equation
The Henderson-Hasselbalch equation is a powerful tool for understanding the pH of buffer solutions. It relates the pH of a solution to the pKa (the acid dissociation constant) and the concentrations of the acid and its conjugate base. The equation is written as: \ \( \text{pH} = \text{p}K_{\text{a}} + \log \frac{[\text{A}^-]}{[\text{HA}]} \) \
Here, \( [\text{A}^-] \) represents the concentration of the conjugate base (salt), and \( [\text{HA}] \) represents the concentration of the weak acid.
This equation is invaluable because it helps us calculate the pH of buffer solutions and thus predict how the solution will react when acids or bases are added. Essentially, it provides a way to understand the balance between acidic and basic components in a solution.
For example, if you know the pKa of an acid (which is a constant for a given acid) and the pH you want to achieve, you can use the equation to find the necessary ratio of salt to acid. This makes it much simpler to prepare buffer solutions in laboratory settings or even in industrial processes.
Understanding this equation is crucial for students learning about buffer solutions and acid-base equilibrium, as it forms the foundation of many calculations and concepts in acid-base chemistry.
Here, \( [\text{A}^-] \) represents the concentration of the conjugate base (salt), and \( [\text{HA}] \) represents the concentration of the weak acid.
This equation is invaluable because it helps us calculate the pH of buffer solutions and thus predict how the solution will react when acids or bases are added. Essentially, it provides a way to understand the balance between acidic and basic components in a solution.
For example, if you know the pKa of an acid (which is a constant for a given acid) and the pH you want to achieve, you can use the equation to find the necessary ratio of salt to acid. This makes it much simpler to prepare buffer solutions in laboratory settings or even in industrial processes.
Understanding this equation is crucial for students learning about buffer solutions and acid-base equilibrium, as it forms the foundation of many calculations and concepts in acid-base chemistry.
Buffer solutions
A buffer solution is a special type of solution that resists changes in its pH when small amounts of an acid or a base are added. Buffer solutions are typically made from a weak acid and its conjugate base (or a weak base and its conjugate acid).
The key role of a buffer is to maintain a stable pH in a system, which is crucial in many biological and chemical processes. For instance, your blood contains a bicarbonate buffer system that balances the pH around 7.4, a necessary condition for most biochemical reactions in the human body.
Buffer solutions work because the weak acid component can neutralize added bases, while the conjugate base component can neutralize added acids. This dual capacity allows the solution to absorb minor fluctuations in pH effectively.
When preparing a buffer solution, the desired pH is usually close to the pKa of the weak acid or base being used. This is because the buffer system has its maximum buffering capacity at a pH equal to the pKa. If the ratio of the concentrations of the conjugate base to the acid is 1:1, the pH of the buffer will be exactly equal to the pKa of the acid. This is one of the reasons why understanding and calculating these ratios is essential for anyone working with buffer systems.
The key role of a buffer is to maintain a stable pH in a system, which is crucial in many biological and chemical processes. For instance, your blood contains a bicarbonate buffer system that balances the pH around 7.4, a necessary condition for most biochemical reactions in the human body.
Buffer solutions work because the weak acid component can neutralize added bases, while the conjugate base component can neutralize added acids. This dual capacity allows the solution to absorb minor fluctuations in pH effectively.
When preparing a buffer solution, the desired pH is usually close to the pKa of the weak acid or base being used. This is because the buffer system has its maximum buffering capacity at a pH equal to the pKa. If the ratio of the concentrations of the conjugate base to the acid is 1:1, the pH of the buffer will be exactly equal to the pKa of the acid. This is one of the reasons why understanding and calculating these ratios is essential for anyone working with buffer systems.
Acid-base equilibrium
Acid-base equilibrium refers to the balance between the concentrations of acids and bases in a solution. This equilibrium is essential because it determines the pH of the solution, which in turn affects many chemical reactions and biological processes.
In solution, acids donate protons (H⁺ ions), whereas bases accept protons. The strength of an acid or a base is typically expressed using the acid dissociation constant (Ka) or the base dissociation constant (Kb). For any given acid, the pKa is the negative logarithm of the Ka value and represents the pH at which the acid is half dissociated. The same concept applies to bases, where pKb is the negative logarithm of Kb.
The Henderson-Hasselbalch equation is a practical application of acid-base equilibrium principles. It allows us to link the pH of a solution to the concentrations of acid and base components. By understanding and manipulating these concentrations, chemists can control the pH of a solution accurately.
Acid-base equilibrium is not just a fundamental concept in chemistry; it's also a critical concept in biology, environmental science, medicine, and many industrial applications. For instance, maintaining the proper pH balance in the human body is vital for enzyme function and metabolic processes. Similarly, environmental scientists monitor the pH of natural water sources to assess ecosystem health. In industrial settings, acid-base equilibrium principles are often applied in processes like fermentation, drug formulation, and waste treatment.
In solution, acids donate protons (H⁺ ions), whereas bases accept protons. The strength of an acid or a base is typically expressed using the acid dissociation constant (Ka) or the base dissociation constant (Kb). For any given acid, the pKa is the negative logarithm of the Ka value and represents the pH at which the acid is half dissociated. The same concept applies to bases, where pKb is the negative logarithm of Kb.
The Henderson-Hasselbalch equation is a practical application of acid-base equilibrium principles. It allows us to link the pH of a solution to the concentrations of acid and base components. By understanding and manipulating these concentrations, chemists can control the pH of a solution accurately.
Acid-base equilibrium is not just a fundamental concept in chemistry; it's also a critical concept in biology, environmental science, medicine, and many industrial applications. For instance, maintaining the proper pH balance in the human body is vital for enzyme function and metabolic processes. Similarly, environmental scientists monitor the pH of natural water sources to assess ecosystem health. In industrial settings, acid-base equilibrium principles are often applied in processes like fermentation, drug formulation, and waste treatment.
