Class 10 Acids, Bases and Salts

Reaction of Acids and Bases with Metals

Acids react with metals to form a salt and release hydrogen gas. This happens because metals that are more reactive than hydrogen can displace hydrogen from acids.

Example: When zinc reacts with hydrochloric acid, zinc chloride (a salt) and hydrogen gas are formed, shown by bubbles and a 'pop' sound when tested with a burning splint:

Zn + 2HCl → ZnCl₂ + H₂ ↑

This type of reaction is called a displacement reaction and is common with metals like magnesium, zinc, iron, and aluminum reacting with acids such as hydrochloric acid or sulfuric acid.

Bases generally do not react with most metals to release hydrogen gas, but some metals like zinc and aluminum can react with strong bases (alkalis) like sodium hydroxide to form a salt and hydrogen gas.

Example:

Zn + 2NaOH + 2H₂O → Na₂Zn(OH)₄ + H₂ ↑

Summary

• Acid + Metal → Salt + Hydrogen gas (common and vigorous with reactive metals)
• Base + Metal → Salt + Hydrogen gas (only specific metals like zinc and aluminum react with bases)

Hydrogen gas release can be tested by bringing a burning splint near the gas; it makes a "pop" sound confirming hydrogen presence.

Reaction of Metal Carbonates and Metal Hydrogencarbonates with Acids

When metal carbonates and metal hydrogencarbonates react with acids, they produce a salt, carbon dioxide gas, and water.

General Reaction Equations:

For metal carbonate:

Metal Carbonate + Acid → Salt + Carbon dioxide + Water

Example:

CaCO₃ + 2HCl → CaCl₂ + CO₂ ↑ + H₂O

For metal hydrogencarbonate:

Metal Hydrogencarbonate + Acid → Salt + Carbon dioxide + Water

Example:

NaHCO₃ + HCl → NaCl + CO₂ ↑ + H₂O

Key Points

• The carbon dioxide gas produced causes bubbling or effervescence.
• Carbon dioxide turns lime water milky when passed through it.
• This reaction is exothermic and releases heat.
• Useful for identifying carbonate and hydrogencarbonate ions in substances.

How do Acids and Bases React with Each Other?

When an acid reacts with a base, they form salt and water. This reaction is called a neutralization reaction.

The hydrogen ions (H⁺) from the acid combine with the hydroxide ions (OH⁻) from the base to form water (H₂O).

General Equation:

Acid + Base → Salt + Water

Example:

Hydrochloric acid reacts with sodium hydroxide:

HCl + NaOH → NaCl + H₂O

Explanation:

• The acid provides H⁺ ions.
• The base provides OH⁻ ions.
• H⁺ and OH⁻ combine to form water.
• Remaining ions form a salt (e.g., NaCl in the example).

This reaction usually releases heat and neutralizes the acid and base, often resulting in a solution with a neutral pH of 7 (if strong acid and strong base are used in equal amounts).

Reaction of Metallic Oxides with Acids

Metallic oxides are compounds formed when metals react with oxygen. They usually behave as bases.

When a metallic oxide reacts with an acid, it produces a salt and water. This is a type of neutralization reaction.

General Reaction:

Metal Oxide + Acid → Salt + Water

Examples:

• Sodium oxide reacts with hydrochloric acid:
  Na₂O + 2HCl → 2NaCl + H₂O
• Calcium oxide reacts with sulfuric acid:
  CaO + H₂SO₄ → CaSO₄ + H₂O
• Copper(II) oxide reacts with hydrochloric acid:
  CuO + 2HCl → CuCl₂ + H₂O

These reactions show the basic nature of metallic oxides because they react with acids to form salt and water, neutralizing the acid.

Reaction of Non-metallic Oxides with Bases

Non-metallic oxides are formed when non-metals combine with oxygen. These oxides usually behave as acidic oxides.

When a non-metallic oxide reacts with a base, they produce salt and water. This reaction is a type of neutralization.

General Reaction:

Non-metallic oxide + Base → Salt + Water

Examples:

• Carbon dioxide reacts with sodium hydroxide:
  CO₂ + 2NaOH → Na₂CO₃ + H₂O
• Carbon dioxide reacts with calcium hydroxide (limewater):
  CO₂ + Ca(OH)₂ → CaCO₃ + H₂O
• Sulfur dioxide reacts with potassium hydroxide:
  SO₂ + 2KOH → K₂SO₃ + H₂O

This reaction shows that non-metallic oxides are acidic in nature as they neutralize bases to form salts and water.

What Do All Acids and All Bases Have in Common?

All acids and bases share some key characteristics:

• Both release ions in water, making their solutions conductive (electrolytes).
• Acids release hydrogen ions (H⁺), and bases release hydroxide ions (OH⁻).
• Both change the color of litmus paper:
  • Acids turn blue litmus paper red.
  • Bases turn red litmus paper blue.
• Both can neutralize each other to form salt and water.
• Both can be corrosive in concentrated forms.
• Both have characteristic tastes (acids taste sour, bases taste bitter) and slippery or burning feelings (bases feel slippery).
• Their pH values differ but show they affect solution acidity or alkalinity:
  • Acids have pH less than 7.
  • Bases have pH greater than 7.

What Happens to an Acid or a Base in a Water Solution?

When an acid or base is mixed in water, they dissociate into ions, which are charged particles.

For Acids:

Acids release hydrogen ions (H⁺) into the solution. These ions often combine with water molecules to form hydronium ions (H₃O⁺).

Example reaction:

HCl → H⁺ + Cl⁻
H⁺ + H₂O → H₃O⁺

For Bases:

Bases release hydroxide ions (OH⁻) in water, increasing the solution’s alkalinity.

Example reaction:

NaOH → Na⁺ + OH⁻

Effect on pH:

The pH scale measures how acidic or basic a solution is. Acids have pH less than 7, bases have pH greater than 7, and pure water is neutral with pH about 7.

Important Note: When mixing concentrated acids with water, always add acid to water slowly to avoid heat and splashing, as the reaction is highly exothermic.

How Strong Are Acid or Base Solutions?

The strength of an acid or base solution refers to how completely it ionizes (breaks into ions) in water.

pH Scale:

The strength is measured by the pH scale, which ranges from 0 to 14:

• pH < 7: Acidic solution (lower pH means stronger acid)
• pH = 7: Neutral solution (pure water)
• pH > 7: Basic (alkaline) solution (higher pH means stronger base)

Strong vs Weak Acids and Bases:

• Strong acids/bases completely ionize in water, producing many ions. Example: Hydrochloric acid (HCl), Sodium hydroxide (NaOH).
• Weak acids/bases partially ionize, so fewer ions are present. Example: Acetic acid (CH₃COOH), Ammonia (NH₃).

The more ions an acid or base releases, the stronger it is and the more it affects the pH of the solution.

Concentration vs. Strength

• Strength is an intrinsic property of the chemical itself (how much it dissociates).
• Concentration refers to the amount of solute (acid or base) dissolved in the solvent (water).

A dilute strong acid (low concentration of HCl in water) may have a higher pH than a concentrated weak acid (high concentration of Acetic Acid in water).

Importance of pH in Everyday Life

The pH level measures how acidic or basic a substance is. It plays an important role in various areas of daily life, affecting health, agriculture, environment, and many household activities.

Health

• Our stomach produces hydrochloric acid (pH 1.5 to 3.5) which helps digest food.
• Excess stomach acid causes acidity, which is neutralized by antacids (which are basic).
• The pH of blood is maintained around 7.35-7.45, which is vital for proper bodily functions.
• Tooth decay happens when mouth pH drops below 5.5 due to acid produced by bacteria; toothpaste helps neutralize this acidity.

Agriculture

• Soil pH affects nutrient availability and plant growth. Most plants grow best in soil with pH 6 to 7.
• Farmers adjust soil pH by adding lime to reduce acidity or sulfur to reduce alkalinity.

Environment

• Acid rain lowers the pH of lakes and rivers, harming aquatic life.
• Maintaining water pH between 6.5 and 8.5 helps protect fish and other organisms.

Household Uses

• Soaps and shampoos are made with pH suitable for skin and hair to prevent irritation.
• Cleaning products are designed with specific pH levels for effective stain removal and disinfection.

Understanding and managing pH helps maintain good health, supports agriculture, protects the environment, and ensures safety in household products.

pH of Salts

Salts have different pH values in their aqueous solutions depending on the strength of the acid and base from which they are formed:

• Salts of a strong acid and a strong base are neutral with a pH of about 7. Example: Sodium chloride (NaCl).
• Salts of a strong acid and a weak base are acidic with pH less than 7. Example: Ammonium chloride (NH₄Cl).
• Salts of a strong base and a weak acid are basic with pH greater than 7. Example: Sodium acetate (CH₃COONa).

This happens because ions from weak acids or weak bases hydrolyze (react with water) affecting the pH, while ions from strong acids and strong bases do not.

Chemicals from Common Salt: Preparation and Uses

Sodium Hydroxide (NaOH)

Preparation:

Sodium hydroxide is prepared industrially by the electrolysis of brine (concentrated sodium chloride solution) through the Chlor-Alkali Process:

2NaCl + 2H₂O → Cl₂ + H₂ + 2NaOH

Uses:

• Manufacturing soaps and detergents.
• Used in paper and textile industries.
• Cleaning agent and drain cleaner.
• Used in petroleum refining and metal processing.
• Water treatment for pH control.

Bleaching Powder (CaOCl₂)

Preparation:

Bleaching powder is prepared by passing chlorine gas over dry slaked lime (Ca(OH)₂):

2Cl₂ + 2Ca(OH)₂ → CaOCl₂ + CaCl₂ + 2H₂O

Uses:

• Used as a disinfectant and bleaching agent.
• Purifying water supplies.
• Used in textile industry to bleach fabrics.

Baking Soda (NaHCO₃)

Preparation:

Prepared by passing carbon dioxide into a solution of sodium carbonate:

Na₂CO₃ + CO₂ + H₂O → 2NaHCO₃

Uses:

• Used in baking as a leavening agent.
• Fire extinguishers (to release CO₂ gas).
• Antacid to relieve acidity.

Washing Soda (Na₂CO₃·10H₂O)

Preparation:

Produced by heating baking soda (sodium bicarbonate):

2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O

Uses:

• Used as a water softener in laundry detergents.
• Cleaning agent for removing grease and stains.
• Used in glass manufacturing.

Are the Crystals of Salts Really Dry?

Salt crystals, such as those of common salt (sodium chloride), are not always completely dry. In fact, most salt crystals contain a small amount of moisture, known as hydration water or moisture content.

Even the purest salt crystals usually have a tiny percentage (around 0.2% to 0.5%) of water trapped within their structure or as surface moisture.

This moisture helps to maintain the crystal structure and can affect properties such as solubility and electrical conductivity.

In industrial uses, controlling the moisture content of salt is important because excessive moisture can cause salt to clump or become less free-flowing.

So, while salt crystals may appear dry, they actually contain a small, measurable amount of moisture.