pH and pH scale 

In chemistry, pH is used to gauge the acidity or basicity of an aqueous solution. It stands for "Potential of Hydrogen" and is defined as the negative logarithm of the hydrogen ion (H⁺) concentration. The pH scale values extends from 0 to 14, with a value of seven considered neutral. Solutions with a pH lower than 7 are acidic, whereas those with a pH higher than 7 are basic.

1.0What is pH?

pH measures the acidity or basicity of aqueous or other liquid solutions. It converts the concentration of hydrogen ions—typically between 1 and 10⁻¹⁴ grams per litre—into a scale ranging from 0 to 14. 

The concept of pH was introduced by Danish biochemist S.P.L. Sorensen, who defined it as the negative logarithm of the hydrogen ion concentration in a solution:

pH = −log (H⁺).

pH is important because it helps control reactions, processes, and balances in a variety of settings. It is used in water purification, chemistry, medicine, biology, agriculture, and other scientific applications.

2.0pH Scale for Acids and Bases

Hydrogen ion (H+) concentration determines whether a solution is basic or acidic. H⁺ ions are more prevalent in acidic solutions than in basic solutions, which have a lower concentration than neutral water.

The pH scale is a numerical indicator of how basic or acidic a solution is. It typically ranges from 0 to 14.

This concentration is measured on the pH scale and is computed as follows:

pH=−log⁡[H⁺]

The pH scale, the negative logarithm of the hydrogen ion activity, makes it easier to depict the concentration of hydronium ions. When the solution is diluted (less than 0.01 M), the activity is almost equivalent to the molarity [H⁺]. A pH scale change of one unit indicates a tenfold change in [H⁺] concentration. For example, the pH changes by two units for every 100-fold change in [H⁺]. Because of its logarithmic nature, slight pH fluctuations caused by temperature are insignificant.

pH measurement is crucial in industries like cosmetics and biology. pH paper, which changes colour depending on the pH of the solution, provides a rapid estimate. With an accuracy of roughly 0.5, modern pH paper with four strips can measure pH from 1 to 14.

Note: Activity, a dimensionless quantity, is:

$a_H^{+}=\frac{[H^{+}]}{mol{L}^{-1}}$

From the definition of pH, the following equation can be written:

$pH=-log{a}_H^{+}=-log(\frac{[H^{+}]}{mol{L}^{-1}})$

For example, an acidic solution of HCl with a concentration of 10⁻² M will have a pH of 2. Similarly, a basic solution of NaOH with [OH⁻] = 10⁻⁴ M and [H₃O⁺] = 10⁻¹⁰ M will have a pH of 10. At 25°C, pure water has a hydrogen ion concentration of [H⁺] = 10⁻⁷ M, so its pH is:

$pH=-log(10^{-7})=7$

Acidic solutions have [H⁺] > 10⁻⁷ M and a pH < 7, while essential solutions have [H⁺] < 10⁻⁷ M and a pH > 7. A neutral solution has a pH of 7.

Although K_w may vary slightly with temperature, these changes are typically small and often ignored in calculations. In aqueous solutions, there must always be a link between pH and pOH. Since the pH scale is logarithmic, even small pH variations can signal significant changes in hydrogen ion concentration.

3.0The Effective Range of the pH Scale

The pH scale is commonly thought to range from 0 to 14, but this is not strictly correct. The pH scale has no fixed upper or lower limit and represents the concentration of hydrogen ions (H⁺).

For instance, at pH 0, the hydronium ion concentration is 1 molar, and at pH 14, the hydroxide ion concentration is 1 molar. Most aqueous solutions typically have H⁺ concentrations between 1 M (pH 0) and 10⁻¹⁴ M (pH 14), making 0 to 14 a practical but not absolute range.

However, pH values can exceed these limits, as hydronium or hydroxide ion concentrations can be greater than 1 molar in extreme conditions.

4.0pH Scale Properties

The pH scale is a logarithmic measure of the acidity or basicity of a solution. Here are some fundamental properties of the pH scale:

• The pH scale ranges from 0 to 14.
• pH value 7 is considered neutral.
• pH values below 7 indicate acidity.
• pH values above 7 indicate basicity.
• Each unit change on the pH scale represents a tenfold difference in hydrogen ion concentration.
• Temperature dependence: pH values can vary slightly with temperature changes.
• Different substances change color at specific pH levels, enabling visual determination of pH.

5.0Limitations of pH Scale

While the pH scale is widely used, it's important to understand its limitations:

• The pH scale is not meaningful for non-aqueous solutions as it is only applicable to aqueous solutions.
• Very strong acids or bases can fall outside the 0-14 range.
• Concentration effects: H measurements may need more accurate in highly concentrated solutions.
• Interference from other ions: Some ions can interfere with pH measurements, leading to inaccurate readings.
• Limited precision: pH measurements typically have about ± 0.1 pH units.

6.0Common Examples of pH in Daily Life

• The pH of pool water is carefully observed and adjusted with acidic or basic chemicals to maintain safety and comfort.
• Antacid tablets, which are basic, neutralize excess stomach acid and alleviate discomfort.
• Blood has a slightly basic pH. Maintaining this balance is essential, as any significant change can damage vital organs.
• pH levels of blood and urine are used to diagnose various health conditions. Abnormal pH can signal underlying diseases.
• Crops thrive at specific pH levels. Adjusting soil pH optimizes plant health and increases yield.
• Enzymes in the body work best at certain pH levels, which is crucial for efficient metabolic processes.

7.0Sample Question on pH and pH scale

Ques. What is pH?

Ans. The pH of a solution is defined as the negative logarithm to base 10 of the activity of hydrogen ions

$a_H^{+}=\frac{[H^{+}]}{mol{L}^{-1}}$