PH Increase From 5 To 7 Explanation And Answer

Understanding pH and its Relationship to Hydroxide Ion Concentration

When delving into the realm of chemistry, grasping the concept of pH is fundamental. pH, which stands for “potential of Hydrogen,” serves as a yardstick for measuring the acidity or alkalinity of a solution. It operates on a logarithmic scale, ranging from 0 to 14, where values below 7 indicate acidity, 7 signifies neutrality, and values above 7 denote alkalinity or basicity. A cornerstone of understanding pH lies in its inverse relationship with the concentration of hydrogen ions ($H^{+}$) in a solution. A lower pH signifies a higher concentration of $H^{+}$ ions, indicative of an acidic environment, while a higher pH implies a lower concentration of $H^{+}$ ions, characteristic of alkaline conditions. This intricate dance between pH and hydrogen ion concentration lays the groundwork for comprehending the behavior of chemical solutions and reactions.

The hydroxide ion ($OH^{-}$$), on the other hand, plays a pivotal role in determining the alkalinity of a solution. Its concentration is inversely proportional to the hydrogen ion concentration, meaning that as the concentration of $H^{+}$ ions decreases, the concentration of $OH^{-}$$ ions increases, and vice versa. This seesaw relationship is crucial in maintaining the delicate balance of acidity and alkalinity in aqueous solutions. The product of the concentrations of $H^{+}$ and $OH^{-}$$ ions in water remains constant at a given temperature, a principle known as the ion product of water ($K_w$). At 25°C, $K_w$ is $1.0 imes 10^{-14}$, ensuring that the concentrations of these two ions are always intertwined. Understanding the interplay between pH and hydroxide ion concentration is paramount for predicting and manipulating chemical reactions in various fields, ranging from environmental science to biochemistry.

To further illustrate the concept, let's consider a scenario where the pH of a solution shifts from 5 to 7. This seemingly small increment on the pH scale carries significant implications for the concentration of hydroxide ions. Recalling the logarithmic nature of the pH scale, each unit change in pH corresponds to a tenfold change in ion concentration. Therefore, an increase of two pH units, from 5 to 7, translates to a hundredfold change in the concentration of either hydrogen or hydroxide ions. In this specific case, as the pH increases, the concentration of hydroxide ions rises dramatically. This change underscores the sensitivity of pH as an indicator of chemical environment, where even subtle shifts can instigate substantial changes in ion concentrations and reaction dynamics. This knowledge is not just academic; it has practical implications in fields like medicine, where precise pH balance is crucial for biological processes, and in industrial chemistry, where pH adjustments can optimize reaction yields and product quality.

Analyzing the pH Shift from 5 to 7

When the pH of a solution ascends from 5 to 7, a notable transformation occurs in its chemical makeup. This shift signifies a decrease in acidity and a corresponding increase in alkalinity. To dissect this phenomenon, we must revisit the fundamental relationship between pH and ion concentrations. pH is mathematically defined as the negative logarithm (base 10) of the hydrogen ion ($H^+}$$) concentration pH = -log[$H^{+$$]. This equation illuminates the inverse correlation between **pH** and [$$H^{+}$$]; as pH increases, [$H^{+}$$] decreases, and vice versa. The logarithmic nature of this relationship means that each **pH** unit represents a tenfold change in [$$H^{+}$$]. Therefore, a shift from pH 5 to pH 6 signifies a tenfold decrease in [$H^{+}$$], and a further shift to pH 7 denotes another tenfold decrease, resulting in an overall hundredfold reduction in hydrogen ion concentration.

Simultaneously, this pH increase heralds a rise in the concentration of hydroxide ions ($OH^{-}$$). As the concentration of $H^{+}$$ diminishes, the equilibrium of water dissociation shifts to produce more $OH^{-}$$ ions. The ion product of water ($K_w$), which at 25°C is $1.0 imes 10^{-14}$, dictates that the product of [$H^{+}$$] and [$$OH^{-}$$] remains constant. Consequently, as [$H^{+}$$] decreases, [$$OH^{-}$$] must increase to maintain this equilibrium. Mathematically, this can be expressed as [$H^{+}$$][$$OH^{-}$$] = $K_w$. This relationship is the cornerstone for understanding how pH changes affect the balance of ions in aqueous solutions. The shift from pH 5 to pH 7 is not merely a cosmetic change; it represents a fundamental alteration in the ionic landscape of the solution, transitioning it from an acidic to a more neutral or alkaline state. This understanding is critical in fields like environmental science, where pH levels in water bodies can dramatically impact aquatic life, and in chemical engineering, where pH control is crucial for many industrial processes.

Considering the quantitative aspect, at pH 5, the hydrogen ion concentration is $10^{-5}$ M (moles per liter), while at pH 7, it is $10^{-7}$ M. This represents a hundredfold decrease in [$H^{+}$$]. Conversely, at pH 5, the hydroxide ion concentration is $10^{-9}$ M (since [$H^{+}$$][$$OH^{-}$$] = $10^{-14}$), and at pH 7, it is $10^{-7}$ M. This reveals that the hydroxide ion concentration has indeed increased by a factor of 100. This calculation underscores the significant impact of pH changes on the ionic environment of a solution. The increase in hydroxide ion concentration is not just a consequence of the decrease in hydrogen ions; it reflects a fundamental shift in the solution's chemistry. This knowledge is invaluable in fields like biochemistry, where enzymes and proteins function optimally within specific pH ranges, and in analytical chemistry, where pH adjustments are often necessary for accurate measurements and reactions. Understanding these relationships allows chemists and scientists to fine-tune chemical environments for a multitude of applications.

The Correct Answer and Why

Given the scenario of a pH increase from 5 to 7, the correct answer is B. Concentration of $OH^{-}$ is 100 times greater than what it was at pH 5. This conclusion is rooted in the fundamental principles of pH and its relationship to ion concentrations in aqueous solutions. To elucidate this, let’s revisit the mathematical underpinnings and the chemical context.

The pH scale, as previously mentioned, is logarithmic, meaning each unit change corresponds to a tenfold change in hydrogen ion ($H^{+}$$) concentration. An increase in **pH** signifies a decrease in [$$H^{+}$$] and a simultaneous increase in hydroxide ion ($OH^{-}$$) concentration. The **ion product of water** ($$K_w$$), a cornerstone of aqueous chemistry, dictates that the product of [$$H^{+}$$] and [$OH^{-}$$] remains constant at a given temperature. At 25°C, $K_w$ is $1.0 imes 10^{-14}$, a fixed value that governs the equilibrium between these two ions. Therefore, any change in [$H^{+}$$] is mirrored by an inverse change in [$$OH^{-}$$], ensuring that their product remains constant.

At pH 5, the [$H^{+}$$] is $10^{-5}$ M, and the [$OH^{-}$$] is $10^{-9}$ M. At pH 7, the [$H^{+}$$] is $10^{-7}$ M, and the [$OH^{-}$$] is $10^{-7}$ M. Comparing the hydroxide ion concentrations, we see that the concentration at pH 7 ($10^{-7}$ M) is indeed 100 times greater than the concentration at pH 5 ($10^{-9}$ M). This hundredfold increase in [$OH^{-}$$] is a direct consequence of the pH increase and the maintenance of $K_w$. This calculation not only validates the correct answer but also underscores the practical implications of pH adjustments in various chemical processes. Whether in industrial settings, where pH control can optimize reaction yields, or in biological systems, where pH balance is critical for enzymatic activity, understanding these ionic shifts is paramount.

The other options presented are incorrect. Option A suggests that the concentration of $OH^{-}$ is one-hundredth (0.01x) what it was at pH 5, which is the inverse of the actual change. This misunderstands the direct relationship between increasing pH and increasing hydroxide ion concentration. This type of error can arise from not fully grasping the inverse logarithmic scale of pH. In educational settings, students sometimes confuse the relationship, thinking a decrease is associated with an increase, but here, the reverse is true for hydroxide ions. To avoid such misconceptions, it's crucial to reinforce the fundamental principles through practical examples and problem-solving sessions.

Conclusion

In summary, the transition of a solution from pH 5 to pH 7 heralds a significant chemical transformation, primarily characterized by a hundredfold increase in hydroxide ion ($OH^{-}$$) concentration. This phenomenon stems from the inherent logarithmic nature of the **pH** scale and the critical equilibrium maintained by the **ion product of water** ($$K_w$$). The **pH** scale, a cornerstone in chemistry, provides a quantitative measure of a solution's acidity or alkalinity, fundamentally linking to the concentrations of hydrogen ($$H^{+}$$) and hydroxide ions. The inverse relationship between pH and [$H^{+}$$], coupled with the constant product of [$$H^{+}$$] and [$OH^{-}$$] governed by $K_w$, dictates that an increase in pH corresponds to a decrease in acidity and a concurrent rise in alkalinity.

The practical implications of understanding pH shifts are vast and varied, permeating diverse fields from environmental science to biochemistry. In environmental contexts, the pH of water bodies serves as a critical indicator of ecological health, influencing the solubility of nutrients and the viability of aquatic life. Deviations from optimal pH ranges can trigger detrimental effects on ecosystems, impacting biodiversity and water quality. Similarly, in the realm of biochemistry, pH plays a pivotal role in enzymatic reactions, protein stability, and cellular function. Enzymes, the biological catalysts that drive myriad biochemical processes, exhibit optimal activity within narrow pH ranges. Fluctuations outside these ranges can disrupt enzymatic function, impairing metabolic pathways and cellular processes.

Moreover, in industrial chemistry, pH control is often integral to optimizing reaction yields and product quality. Many chemical reactions are pH sensitive, with their rates and equilibrium positions significantly influenced by the acidity or alkalinity of the reaction medium. Careful pH adjustments can enhance reaction efficiency, minimize byproduct formation, and improve the overall process economics. Furthermore, in analytical chemistry, pH plays a crucial role in various analytical techniques, including titrations, spectrophotometry, and chromatography. pH adjustments are often necessary to ensure accurate measurements and reliable results. The understanding of pH and its influence on chemical systems thus extends far beyond theoretical concepts, bridging into practical applications that impact our daily lives and shape technological advancements.