Question

The \(\mathrm{pH}\) of a solution is 6.7 . From this statement alone, can you conclude that the solution is acidic? If not, what additional information would you need? Can the \(\mathrm{pH}\) of a solution be zero or negative? If so, give examples to illustrate these values.

Short answer

Yes, the solution is acidic. pH can be zero or negative in strong acids, like 1 M HCl.

Step-by-step solution

Understand pH and Acidic Solutions

The pH scale ranges from 0 to 14, where a pH less than 7 indicates an acidic solution, a pH of 7 represents a neutral solution, and a pH greater than 7 indicates a basic (alkaline) solution. Thus, if a solution has a pH of 6.7, it is considered acidic because it is less than 7.

The pH Scale Limits

The pH scale is commonly understood to extend from 0 to 14; however, it can exceed these boundaries in highly concentrated acidic or basic solutions. A pH value of 0 or negative is possible in very strong acids. For example, a solution of approximately 1 M hydrochloric acid (HCl) can have a pH close to 0, while even more concentrated acids can result in negative pH values.

Key concepts

Acidic Solutions

The concept of acidic solutions is central to understanding pH. When we say a solution is acidic, we are talking about its concentration of hydrogen ions (H extsuperscript{+}).
An acidic solution has a higher concentration of hydrogen ions compared to a neutral substance, like pure water. The pH scale, which ranges from 0 to 14, helps us measure how acidic or basic a solution is.

• A pH less than 7 means the solution is acidic.
• A pH of exactly 7 is considered neutral, like pure water.
• And a pH greater than 7 points to a basic (alkaline) solution.

Since the solution in our exercise has a pH of 6.7, it clearly falls into the acidic category, as it is below 7. This simple comparison to the neutral point indicates the solution's acidic nature. Understanding this classification helps in grasping how substances interact chemically.
Acidic solutions are widespread, from household vinegar to lemon juice, each with varying pH levels but all classified as being acidic by their pH.

pH Scale Limits

The pH scale is a fundamental aspect when discussing acidity and alkalinity. While typically taught as going from 0 to 14, these are not absolute limits due to the varying nature of chemical solutions.
Normal pH scale value representations are:

• 0 to 6.99 for acidic solutions,
• 7 for neutral solutions,
• and greater than 7 for basic solutions.

However, in certain situations, the pH can go beyond these norms. Highly concentrated strong acids or bases can skew this understanding. For example, strong acids like sulfuric acid or hydrochloric acid in high concentrations can result in a pH that is equal to or below 0. This occurs because the logarithmic nature of the pH scale allows negative values when the concentration of hydrogen ions becomes quite high.
It's crucial to remember that the typical pH scale limits are a guideline, with exceptions arising in cases of extreme acidity or basicity.

Strong Acids

When discussing pH and acidity, strong acids play a particularly significant role. A strong acid is one that completely dissociates in water, releasing a large number of hydrogen ions. This complete dissociation is what leads to the possibility of very low or even negative pH values.
Common examples of strong acids include:

• Hydrochloric acid (HCl)
• Nitric acid (HNO₃)
• Sulfuric acid (H₂SO₄)

These acids are highly reactive and their solutions can have extremely low pH values, sometimes approaching or dropping below 0.
The presence of a high concentration of hydrogen ions in these solutions is what contributes to their strength and their ability to lower pH beyond the standard scale limits. Concentrated forms of these acids can thus alter the typical boundaries of the pH scale, demonstrating the scale's flexibility in real-world applications.
Understanding strong acids not only helps in predicting pH but also in anticipating their behavior in chemical reactions, which is key in fields like chemistry and biochemistry.