Question

When hydrochloric acid, HCl , is added dropwise to a buffered solution, the component of the buffer consumes the added hydrogen ion.

Short answer

In summary, when hydrochloric acid (HCl) is added dropwise to a buffered solution, the conjugate base (A-) present in the buffer consumes the added hydrogen ions (H+) by reacting with them and forming the weak acid HA. This reaction prevents a significant change in pH and maintains the buffered solution's pH within the desired range. In this process, the Henderson-Hasselbalch equation helps to understand the relationship between the pH and the concentrations of the buffer components.

Step-by-step solution

Understand the concept of buffer solutions

Buffer solutions are aqueous solutions that can resist changes in pH when small amounts of acid or base are added to them. They contain a weak acid and its conjugate base or a weak base and its conjugate acid. The main function of a buffer solution is to maintain a relatively constant pH in the presence of added acids or bases. In our case, we have a buffer solution and HCl, a strong acid, is being added to it. We need to find out how the component of the buffer consumes the added hydrogen ion from HCl.

Introduce the concept of the Henderson-Hasselbalch equation

The Henderson-Hasselbalch equation describes the relationship between the pH of a buffer solution and the concentration of its components. For a buffer composed of a weak acid (HA) and its conjugate base (A-), the equation is: \(pH = pK_a + log \frac{[A^-]}{[HA]}\) where - \(pH\) is the pH of the buffer solution - \(pK_a\) is the acid dissociation constant of the weak acid (HA) - \([A^-]\) is the concentration of the conjugate base (A-) - \([HA]\) is the concentration of the weak acid (HA) The expression helps us to understand how the buffer components interact with the added acid or base and maintain the pH of the solution.

The reaction between the buffer solution and added HCl

When HCl is added to the buffered solution, it will dissociate into H+ and Cl- ions: \(HCl \rightarrow H^+ + Cl^-\) In the buffered solution, there is a weak acid HA and its conjugate base A-. The added H+ ions from the HCl will react with the conjugate base A-, which acts as a proton acceptor. The reaction can be written as: \(A^- + H^+ \rightarrow HA\) This reaction removes the added H+ ions from the solution, preventing a significant change in pH.

Explain how the buffer solution maintains its pH

As HCl is added dropwise, continuous small amounts of H+ ions are added to the solution. However, because the proton acceptor A- is present in the buffered solution, it reacts with the added H+ ions, forming the weak acid HA. This consumption of H+ ions prevents a significant change in pH. The buffer system effectively maintains the pH as the added H+ ions are consumed by the A- ions without greatly affecting the overall concentration of HA and A- in the solution. In conclusion, when hydrochloric acid (HCl) is added dropwise to a buffered solution, the component of the buffer, specifically the conjugate base (A-), consumes the added hydrogen ions, preventing a significant change in pH and maintaining the buffered solution's pH in the desired range.

Key concepts

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation is a foundational formula in chemistry that helps us understand buffer solutions. It describes how the pH of a buffer solution depends on the concentration of its components: the weak acid and its conjugate base. The equation is expressed as:\[ pH = pK_a + \log \frac{[A^-]}{[HA]} \]In this equation, \(pH\) represents the acidity level of the buffer solution. The term \(pK_a\) is the acid dissociation constant specific to the weak acid involved, reflecting its strength. Meanwhile, \([A^-]\) stands for the concentration of the conjugate base, and \([HA]\) represents the weak acid concentration.This equation is instrumental in showing us how a buffer solution can manage to keep its pH relatively stable. By helping us see how changes in the concentration of the acid or base affect the solution, this formula informs us about how the buffer components interact to resist pH shifts upon the addition of small amounts of acid or base. It's a crucial tool for anyone looking to understand the balancing act involved in buffer solutions.

Weak Acids and Conjugate Bases

In the world of buffer solutions, weak acids and their conjugate bases play central roles. A weak acid, unlike a strong acid, does not completely dissociate in solution. Instead, it partially ionizes, creating a delicate balance in the solution. The conjugate base of a weak acid is the species that remains after the acid has donated its hydrogen ion (proton). This base is crucial in buffer systems because it acts as a proton acceptor. When strong acids like hydrochloric acid (HCl) are added to a buffer solution, it is the conjugate base that steps in to neutralize the added hydrogen ions.

• Weak Acid (HA) - Only partially disassociates in water, establishing an equilibrium between its undissociated form and ions.
• Conjugate Base (A^-) - Accepts protons, crucial for resisting pH change when acids are introduced into the buffer.

This teamwork between the weak acid and its conjugate base is what makes buffers so effective, reinforcing the buffer solution’s ability to maintain a steady pH even when small amounts of acids or bases are added.

pH Resistance Mechanism

The pH resistance mechanism is the heart of how buffer solutions operate effectively. When faced with either an acidic or basic challenge, buffers have an innate ability to "neutralize" these additions, keeping the pH from shifting dramatically. This is due to the reaction of its components - the weak acid and conjugate base.When a strong acid like HCl is added to a buffer solution, it releases hydrogen ions (H⁺). The conjugate base in the buffer quickly reacts with these ions:\[ A^- + H^+ \rightarrow HA \]This reaction effectively consumes these hydrogen ions, forming more of the weak acid (HA) and minimizing the pH change. On the flip side, if a base is added, the weak acid component can donate hydrogen ions (H⁺) to counteract the effect of the base, thus again stabilizing the pH.

• Conjugate base neutralizes added strong acids.
• Weak acid provides protons to neutralize added bases.
• Continual reactions adjust {HA} and {A^-} levels without significant pH alteration.

This dynamic equilibrium enables buffer solutions to be resilient and reliable in maintaining a consistent pH, which is crucial for many scientific and industrial applications.