11.08 Chapter Summary

1. Properties of Acids and Bases

Acids:

• pH: Below 7
• Taste: Sour (when safe to taste)
• Corrosive: Can cause damage to materials
• Chemical Behavior:
  • Neutralization: Acids neutralize bases to form salts and water.
  • Ionization in Water: Release hydrogen ions (H⁺), making the solution acidic.

• Example: Hydrochloric Acid (HCl)
• HCl (aq)→H⁺ (aq)+Cl⁻ (aq)

Bases (Alkalis):

• pH: Above 7
• Properties:
  • Taste: Bitter (when safe to taste)
  • Feel: Slippery
  • Corrosive: Can cause damage to materials
• Chemical Behavior:
  • Neutralization: Bases neutralize acids to form salts and water.
  • Ionization in Water: Release hydroxide ions (OH⁻), making the solution alkaline.

• Example: Sodium Hydroxide (NaOH) NaOH (s)→Na⁺ (aq)+OH⁻ (aq)

2. Reactions Involving Acids

a. Acids and Metals:

• Condition: Only metals above hydrogen in the reactivity series react with dilute acids.
• Products: Salt and hydrogen gas (H₂).

• General Equation: Acid+Metal→Salt+H₂

Examples:

• Hydrochloric Acid and Magnesium: Mg + 2HCl → MgCl₂ + H₂ ​
• Sulfuric Acid and Magnesium: Mg + H₂SO₄ → MgSO₄ + H₂
• Nitric Acid and Magnesium: Mg + 2HNO₃ → Mg(NO₃)₂ + H₂

b. Acids and Bases (Neutralization):

• Reaction: Produces salt and water.

• General Equation: Acid+Base→Salt+H2O

Examples:

• Hydrochloric Acid and Magnesium Hydroxide:
  Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O
• Sulfuric Acid and Magnesium Oxide:
  MgO + H₂SO₄ → MgSO₄ + H₂O
• Nitric Acid and Magnesium Hydroxide: Mg(OH)₂+2HNO₃→Mg(NO₃)₂+2H₂O

c. Acids and Metal Carbonates:

• Reaction: Produces salt, carbon dioxide (CO₂), and water.

• General Equation: Acid+Metal Carbonate→Salt+CO₂+H₂O

Examples:

• Hydrochloric Acid and Magnesium Carbonate:
  MgCO₃ + 2HCl → MgCl₂ + CO₂ + H₂O
• Sulfuric Acid and Magnesium Carbonate:
  MgCO₃ + H₂SO4 → MgSO₄ + CO₂ + H₂O
• Nitric Acid and Magnesium Carbonate:
  MgCO₃ + 2HNO₃ → Mg(NO₃)₂ + CO₂ + H₂O

3. Indicators

Types:

• Natural Indicators: Extracted from plants (e.g., litmus from lichens).
• Synthetic Indicators: Organic compounds with distinct color changes at specific pH levels (e.g., thymolphthalein, methyl orange).

Common Indicators and Color Changes:

Usage:

• Universal Indicator: A mixture of indicators that display a range of colors corresponding to different pH levels. A drop is added to the solution, and the color is matched to a pH chart.

• Note: Synthetic indicators are preferred in titrations due to their sharp color changes, which help identify the endpoint accurately.

4. pH Scale and Hydrogen Ion Concentration

pH Scale:

• Range: 0 to 14
  • pH < 7: Acidic
  • pH = 7: Neutral
  • pH > 7: Alkaline

Strength of Acids and Bases:

• Acids:
  • Strong Acids: Completely dissociate in water (e.g., HCl, H₂SO₄), resulting in low pH.
  • Weak Acids: Partially dissociate in water (e.g., CH₃CH₂COOH), resulting in higher pH compared to strong acids.
• Bases:
  • Strong Bases (Alkalis): Completely dissociate in water (e.g., NaOH).
  • Weak Bases: Partially dissociate in water.

Logarithmic Nature:

• Each pH unit represents a tenfold change in hydrogen ion concentration.
  • Example: pH 3 has ten times more H⁺ ions than pH 4; pH 2 has 100 times more H⁺ ions than pH 4.

5. Oxides Classification

• Oxides: Compounds consisting of oxygen combined with another element.

Types:

• Acidic Oxides:
  • Formed with non-metals (e.g., CO₂, SO₂).
  • React with bases to form salts and water.
  • Produce acidic solutions in water.
• Basic Oxides:
  • Formed with metals (e.g., MgO, CaO).
  • React with acids to form salts and water.
  • Produce alkaline solutions in water.
• Amphoteric Oxides: Can behave as both acids and bases depending on the reactant (e.g., ZnO, Al₂O₃).
• Neutral Oxides: Do not react with acids or bases (e.g., N₂O, CO).

6. Preparing Salts

a. Soluble Salts:

• Definition: Compounds formed by the neutralization of an acid and a base, soluble in water.

Naming:

• Two-Part Name:
  1. Cation: Derived from the base (e.g., sodium from NaOH).
  2. Anion: Derived from the acid (e.g., chloride from HCl).

Examples:

• HCl and NaOH: HCl + NaOH → NaCl + H₂O
• ZnO and H₂SO₄: ZnO + H₂SO₄ → ZnSO₄ + H₂O

Methods:

• Method A: Reacting an acid with a metal, insoluble base, or insoluble carbonate.
  1. Add dilute acid to a beaker and heat.
  2. Gradually add the solid reactant until excess forms.
  3. Filter and evaporate the solution to crystallize the salt.
• Method B: Titration between a dilute acid and a soluble base.
  1. Use a pipette to add base to a flask with indicator.
  2. Slowly add acid from a burette until the endpoint is reached (color change).
  3. Calculate the volume of acid used and crystallize the salt from the mixture.

b. Insoluble Salts:

• Definition: Formed through precipitation reactions where the product salt is insoluble in water.

Preparation:

1. Dissolve two soluble salts in water.
2. Mix the solutions using a stirring rod.
3. Filter out the precipitate (insoluble salt).
4. Wash, dry, and collect the pure insoluble salt.

Example: Lead(II) Sulfate Formation
Pb(NO₃)₂(aq) + K₂SO₄(aq) → PbSO₄(s) + 2KNO₃(aq)

7. Solubility Rules

General Guidelines:

8. Hydrated and Anhydrous Salts

Hydrated Salts:

• Definition: Contain water molecules within their crystal structure (water of crystallization).
• Example: Hydrated Copper(II) Sulfate (CuSO₄⋅5H₂O) is blue.

Anhydrous Salts:

• Definition: Contain no water in their structure.
• Example: Anhydrous Copper(II) Sulfate (CuSO₄) is white.

Conversion:

• Dehydration: Heating hydrated salts removes water, forming anhydrous salts. CuSO₄⋅5H₂O → CuSO₄ + 5H₂O
• Hydration: Adding water to anhydrous salts recrystallizes hydrated salts. CuSO₄ + 5H₂O → CuSO₄⋅5H₂O

9. Additional Concepts

Proton Transfer (Advanced):

• Acids: Proton donors (release H⁺ ions).
• Bases: Proton acceptors (accept H⁺ ions).

Neutralization Reaction (Net Ionic Equation):

H⁺(aq) + OH⁻(aq) → H₂O(l)

Ammonium Salts and Alkalis:

• Reaction: Ammonium salts react with alkalis to produce ammonia gas (NH₃), water, and a salt.

• Example: NH₄Cl+NaOH→NaCl+H₂O+NH₃

Testing for Ammonia:

• Method: Pass ammonia gas over damp red litmus paper; it turns blue if NH₃ is present.

10. Examiner Advice

• Neutralization Definition: Only reactions producing salt and water are neutralizations.
• pH Determination: Understand the logarithmic relationship between pH and H⁺ concentration.
• Indicator Selection: Use appropriate indicators based on the sharpness of color change required.
• Solubility Rules: Memorize key solubility rules to predict reaction products.
• Hydration States: Recognize hydrated vs. anhydrous forms and their formulas.