pH and Solutions

1.0Introduction

The term pH, short for "potential of hydrogen," was introduced by Danish biochemist Søren Sørensen in 1909. He conceived this concept as a more straightforward method to represent the exceedingly minute concentrations of hydrogen (H+) and hydroxide (OH−) ions found in aqueous solutions.

pH is defined as:

$\boxed{\text{pH}=-{\log }_{10}[{\text{H}}^{+}]}$

This means that the pH of a solution is the negative logarithm (base 10) of the hydrogen ion concentration.

For example, if the hydrogen ion concentration is

 1×10⁻³  the pH would be:

$\text{pH}=-{\log }_{10}(1\times 10^{-3})=3$

2.0pOH and Its Relation with pH

Just like pH measures the concentration of hydrogen ions, pOH measures the concentration of hydroxide ions:

$\boxed{\text{pOH}=-{\log }_{10}[{\text{OH}}^{-}]}$

For aqueous solutions at 25°C, pH and pOH are related by:

$\boxed{\text{pH}+\text{pOH}=14}$

3.0The pH Scale   

The pH scale (0 to 14) helps to determine whether a solution is acidic, neutral, or basic:

4.0Different Solutions  and pH 

Acidic Solutions

• High concentration of hydronium ions (H₃O⁺)
• Low concentration of hydroxide ions (OH⁻)
• pH is less than 7

Basic (Alkaline) Solutions

• High concentration of hydroxide ions (OH⁻)
• Low concentration of hydronium ions (H₃O⁺)
• pH is greater than 7

5.0The pH of Pure Water

Characterized by a pH of 7, pure water's neutrality is attributed to its autoionization equilibrium. In this reversible reaction, water molecules (H2​O) dissociate, continuously forming precisely equivalent molar concentrations of hydrogen (H+) and hydroxide (OH−) ions.

H₂​O ⇌ H⁺+OH^-

This equilibrium ensures that the concentration of H+ ions is exactly equal to that of OH− ions, preventing any dominant acidic or basic properties and establishing a neutral pH.

Neutral Solution and pH 

The pH of a neutral solution is 7.

• In such solutions, the number of hydrogen ions [H⁺] is equal to the number of hydroxide ions [OH⁻].
• Examples include pure water and human blood, which is slightly basic but close to neutral (around pH 7.4).

When the [H⁺] = [OH⁻], the solution is neutral, and the pH = 7. This is the benchmark for determining whether a solution is acidic (<7) or basic (>7).

6.0pH Role in Everyday Life 

One of the most relatable examples of pH in daily life is Oral Health. The normal pH of the mouth is around 6, which is slightly acidic; however, when we consume sugary foods or drinks, bacteria in the mouth break down these sugars and produce acids as byproducts.

If the pH level drops below 5.5, these acids begin to corrode  the tooth enamel (which is made of calcium phosphate), leading to tooth decay and cavities.

To combat this:

• Toothpastes are mildly alkaline to help neutralize  mouth acids.
• This helps restore a safe pH, protecting the enamel and maintaining oral health.

7.0Salt Solution and pH

Salts are generally formed through the neutralisation reaction between an acid and a base. While we might expect all salts to be neutral with a pH of 7, in reality, the pH of salts can vary—they may be acidic, basic, or neutral, depending on the strength of the parent acid and base.

Not All Salts Are Neutral: While some salts achieve the ideal pH of 7, the majority will have a pH either below 7 (acidic) or above 7 (basic). This means salts can be neutral, acidic, or basic in nature.

When Salts Are Neutral (pH = 7):

• Neutral salts are formed when a strong acid perfectly reacts with a strong base.
• They are the result of a complete neutralization reaction.
• Example: Common table salt, sodium chloride (NaCl), is a perfect example, with a pH of 7.

When Salts Are Acidic (pH < 7):

• Acidic salts are created when a strong acid encounters a weak or mild base.
• The reason they're acidic is that the weak base isn't strong enough to fully neutralise the strong acid; it only achieves partial neutralisation.
• Examples: You'll find acidic salts in compounds like ammonium sulfate and ammonium chloride.

When Salts Are Basic (pH > 7):

• Basic salts are formed when a strong base reacts with a weaker acid.
• In this scenario, the weaker acid can't completely overcome the strong base, leading to only partial neutralization of the base.
• Examples: Familiar basic salts include sodium carbonate, calcium carbonate, and magnesium carbonate.

8.0Importance of pH in Daily Life

The pH of a substance plays a crucial role in maintaining balance in living systems. Most living organisms can survive only within a narrow pH range. Any significant deviation can disrupt biological processes and even threaten life.

• Environmental Impact – Acid Rain: Acid rain has a pH lower than 7. When this acidic water enters rivers and lakes, it lowers the natural pH of the water bodies. This creates a harsh environment for aquatic organisms, making it difficult for them to survive.
• Digestion and Stomach Health: Our stomach contains hydrochloric acid (HCl), which aids in the digestion of food. However, during indigestion, excess acid is produced, causing pain and discomfort. To neutralize this acidity, antacids (mild bases like magnesium hydroxide) are used. These increase the pH and provide relief from the irritation.
• Dental Care: Bacteria in the mouth break down food particles and release acids, which lower the pH inside the mouth. This acidic environment can lead to tooth decay. To prevent this, we use toothpaste, which is basic in nature, to neutralize the acids and maintain a healthy pH balance in the mouth.
• Insect Stings – Bee Sting Relief: A bee sting injects methanoic acid (formic acid) into the skin, causing a burning sensation and pain. To relieve this, we apply baking soda (a mild base), which helps neutralize the acid and restore the normal pH of the skin, providing quick relief.