Acids, Bases, Salts, and the pH Scale

Strong vs. Weak Bases

• Strong Bases: These ionise completely to furnish a large amount of $OH^-$ ions, such as Sodium hydroxide ($NaOH$) and Potassium hydroxide ($KOH$).
• Weak Bases: These ionise only partially, such as Magnesium hydroxide ($Mg(OH)_2$) or Ammonium hydroxide ($NH_4OH$).

Exam Tip: Remember that "strength" refers to the degree of ionisation, not how much of the substance you have. A dilute strong acid can still be "stronger" in terms of ion concentration than a concentrated weak acid.

Identifying Substances: Indicators

Since many acids and bases are corrosive and dangerous to taste or touch, scientists use indicators to identify them. Indicators are substances that change color or smell depending on whether they are in an acidic or basic environment.

There are three main types of indicators:

1. Natural Indicators: Derived from nature, such as Litmus (which turns red in acid and blue in base) or Turmeric (which remains unchanged in acid but turns red in base).
2. Synthetic Indicators: Man-made chemicals like Phenolphthalein (which turns pink in a base but remains colorless in acid) and Methyl orange (which turns red in acid and yellow in base).
3. Olfactory Indicators: These change their characteristic smell in different media. For instance, Onion or Vanilla essence retain their smell in acidic solutions but lose it in basic solutions.

Chemical Properties and Reactions

Acids and bases exhibit distinct chemical behaviors when they encounter other substances.

Reactions with Metals

When an acid reacts with a metal, it typically produces a salt and hydrogen gas: $\text{Acid} + \text{Metal} \rightarrow \text{Salt} + H_2(g)$ For example: $2HCl + Zn \rightarrow ZnCl_2 + H_2 \uparrow$. A common test for hydrogen gas is the "pop sound" produced when a burning candle is brought near the gas bubbles.

Bases can also react with certain metals to produce salt and hydrogen gas, such as: $2NaOH + Zn \rightarrow Na_2ZnO_2 + H_2 \uparrow$ where $Na_2ZnO_2$ is sodium zincate.

Reactions with Carbonates

Acids react with metal carbonates or metal hydrogen carbonates to produce salt, water, and carbon dioxide gas: $\text{Acid} + \text{Metal Carbonate} \rightarrow \text{Salt} + CO_2 + H_2O$ An example is: $2HCl + Na_2CO_3 \rightarrow 2NaCl + CO_2 + H_2O$. The $CO_2$ produced can be tested by passing it through lime water, which turns milky due to the formation of calcium carbonate.

Neutralization: The Core Reaction

When an acid and a base react, they neutralize each other's properties. The $H^+$ from the acid combines with the $OH^-$ from the base to form water: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$ $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$ A classic example is the reaction between Hydrochloric acid and Sodium hydroxide to form Sodium chloride (common salt) and water: $NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)$

Salts and Their Types

A salt is an electrically neutral ionic compound formed from a neutralization reaction. Salts are composed of cations (positive ions) and anions (negative ions) in related numbers, ensuring the final product is neutral.

Salts can be classified based on the pH of the solution they form:

• Neutral Salts: Formed from a strong acid and a strong base. They have a pH of 7. Example: $NaCl$.
• Acidic Salts: Formed from a strong acid and a weak base. They have a pH < 7. Example: $NH_4Cl$.
• Basic Salts: Formed from a weak acid and a strong base. They have a pH > 7. Example: $Na_2CO_3$.

The pH Scale and Its Importance

The pH scale is used to measure the concentration of $H^+$ ions in a solution. Because $H^+$ concentrations are often very small, pH is defined as the negative logarithm of the $H^+$ concentration: $\text{pH} = -\log [H^+]$

The scale typically ranges from 0 to 14:

• pH < 7: Acidic solution
• pH = 7: Neutral solution (e.g., pure water)
• pH > 7: Basic solution

Dilution Effects: Diluting an acid increases its pH (moving it closer to 7), while diluting a base decreases its pH (also moving it closer to 7).

pH in Everyday Life

The pH balance is critical for biological and environmental health:

• Human Body: Our body functions within a narrow pH range of 7.0 to 7.8.
• Digestive System: The stomach produces $HCl$ (pH 1.5–3.0) to aid digestion. Excessive acidity can be treated with antacids like milk of magnesia.
• Dental Health: Tooth decay begins when the pH of the mouth falls below 5.5. Basic toothpastes help neutralize this acid.
• Environment: Rainwater with a pH below 5.6 is considered acid rain, which can harm plants and aquatic life.

Note on Safety: When diluting an acid, always add acid to water slowly with constant stirring. Adding water to a concentrated acid can cause a violent, exothermic reaction that may splash acid out of the container.