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Chapter 14: Problem 27

What is a buffer?

Short Answer

Expert verified
A buffer is a solution that resists pH changes when acids or bases are added, typically made of a weak acid and its conjugate base or vice versa.

Step by step solution

01

Definition of a Buffer

A buffer is a solution that can resist pH change upon the addition of an acidic or basic components. It typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid.
02

Function of a Buffer

Buffers work by reacting with any added acid or base to prevent drastic changes in pH. This is important for processes that require a constant pH, such as biological systems and chemical reactions.
03

Buffer Capacity

Buffer capacity is a measure of the efficiency of a buffer in resisting changes in pH. It depends on the amount of acid and conjugate base (or base and conjugate acid) the buffer contains and the pH range over which the buffer acts effectively.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

pH Change Resistance
Understanding the concept of pH change resistance is essential in chemistry, especially when studying buffers. A buffer essentially acts like a chemical sponge, soaking up excess hydrogen ions (H+) or hydroxide ions (OH-) to maintain a stable pH environment. This is crucial in many biological systems where enzymes and other macromolecules require a specific pH to function correctly. The pH resistance comes from the buffer's ability to neutralize added acids or bases through a set of reversible reactions, which prevents the pH from swinging wildly and maintains equilibrium within a narrow range.
Acid-Base Conjugate Pairs
The beauty of a buffer lies in its components: the acid-base conjugate pairs. These pairs consist of a weak acid and its conjugate base, or conversely, a weak base and its conjugate acid. This duo works in tandem to counteract pH changes. When an external acidic substance is introduced, the conjugate base will neutralize it, forming the weak acid. Conversely, when a base is added, the weak acid component can donate a proton to neutralize the base, again forming the conjugate base. This pivot between acid and base within the pair allows the solution to resist drastic pH changes, underpinning the buffer's role in chemical and biological systems.
Buffer Capacity
Buffer capacity is a term that denotes how well a buffer can keep the pH stable when acids or bases are added. It's a bit like the buffer's strength or endurance. This capacity is influenced by the concentrations of the conjugate acid-base pairs; the higher their concentration, the greater the buffer's capacity to maintain pH. However, buffer capacity isn't infinite and has its limits based on the amounts and proportions of the components present. Within a certain pH range, the buffer can efficiently neutralize added acids or bases, but once this capacity is exceeded, the pH will start to change more noticeably.
Chemical Reactions
Buffers play an essential role in controlling the environment for chemical reactions. Many reactions in chemistry are pH-dependent, meaning that they require a constant pH to proceed correctly or to optimize yields. Buffers ensure that the pH remains steady, even if some of the reactants or products are acidic or basic. This stable pH environment can significantly affect reaction rates and the equilibrium of reversible reactions. In essence, a buffer acts as a safeguard, preserving the optimal conditions for a chemical reaction to occur effectively and predictably.

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Most popular questions from this chapter

Determine the \(\mathrm{pH}\) of each solution and classify it as acidic, basic, or neutral. (a) \(\mathrm{pOH}=12.5\) (b) \(\mathrm{pOH}=5.5\) (c) \(\mathrm{pOH}=0.55\) (d) \(\mathrm{pOH}=7.98\)

Calculate the \(\mathrm{pH}\) of each solution: (a) \(1.34 \times 10^{-3} \mathrm{M} \mathrm{HClO}_{4}\) (b) \(0.0211 \mathrm{M} \mathrm{NaOH}\) (c) \(0.0109 \mathrm{MHBr}\) (d) \(7.02 \times 10^{-5} \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\)

Calculate \(\left[\mathrm{OH}^{-}\right]\)given \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\)in each aqueous solution and classify the solution as acidic or basic. (a) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=1.3 \times 10^{-3} \mathrm{M}\) (b) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=9.1 \times 10^{-12} \mathrm{M}\) (c) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=5.2 \times 10^{-4} \mathrm{M}\) (d) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=6.1 \times 10^{-9} \mathrm{M}\)

For each reaction, identify the Bronsted-Lowry acid, the Bronsted-Lowry base, the conjugate acid, and the conjugate base. (a) \(\mathrm{HBr}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{3} \mathrm{O}^{+}(a q)+\mathrm{Br}^{-}(a q)\) (b) \(\mathrm{NH}_{3}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{NH}_{4}^{+}(a q)+\mathrm{OH}^{-}(a q)\) (c) \(\mathrm{HNO}_{3}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{3} \mathrm{O}^{+}(a q)+\mathrm{NO}_{3}^{-}(a q)\) (d) \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{~N}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}^{+}(a q)+\mathrm{OH}^{-}(a q)\)

For each \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\), determine the \(\mathrm{pH}\) and state whether the solution is acidic or basic. (a) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=0.0025 \mathrm{M}\) (b) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=1.8 \times 10^{-12} \mathrm{M}\) (c) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=9.6 \times 10^{-9} \mathrm{M}\) (d) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=0.0195 \mathrm{M}\)
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