Chapter 10 Acids and Bases (notes)

Chapter 10: Acids and Bases

Acids

• Definition: A substance that forms hydrogen ions (H+) when dissolved in water.

• Strong Acids:

  • Completely ionize in water, leading to a high concentration of H+ ions.

  • Examples and Ionization Equations:

    1. Hydrochloric acid (HCl) ➜ H+ + Cl- (Chloride)

    2. Nitric acid (HNO3) ➜ H+ + NO3- (Nitrate)

    3. Sulfuric acid (H2SO4) ➜ 2H+ + SO4^2- (Sulfate)

• Weak Acids:

  • Partially ionize in water, resulting in a lower concentration of H+ ions.

  • Example: Ethanoic acid (CH3COOH) ➜ CH3COO- + H+ (Ethanoate)

• Hydronium Ions Formation: H+ ions combine with water molecules to form hydroxonium ions (H3O+).

  • Ionization in water can be represented as: HCl + H2O ➜ H3O+ + Cl-

• Acidic Property:

  • Presence of H+ ions gives acids their acidic properties.

  • pH value ranges from 1-6.

• Solution State: All acids exist in aqueous solutions.

Bases

• Definition: Substances that react with acids to form salts and water.

• Types of Bases:

  • Metal oxides (e.g., magnesium oxide, MgO)

  • Metal hydroxides (e.g., sodium hydroxide, NaOH)

  • Ammonia (NH3)

• Solubility:

  • Not all bases dissolve in water.

  • Most metal oxides and hydroxides are insoluble, except those with Group I ions and some Group II elements like Ba2+ and Ca2+.

• Ionization in Water:

  • Metal hydroxide: KOH ➜ K+ + OH-

  • Metal oxide: K2O + H2O ➜ 2KOH ➜ K+ + OH-

• Alkali: A solution formed when a soluble base is dissolved in water that ionizes to form hydroxide ions (OH-).

• Strong Alkali:

  • Ionize completely in water to yield a high concentration of OH- ions.

  • Examples:

    1. Sodium hydroxide (NaOH) ➜ Na+ + OH-

    2. Potassium hydroxide (KOH) ➜ K+ + OH-

• Weak Alkali:

  • Partially ionize in water to yield a lower concentration of OH- ions.

  • Example: Ammonia solution

    • NH3 + H2O ➜ NH4+ + OH-

• Alkaline Properties: OH- confers alkaline properties, with a pH value of 8-14.

  • Exceptions: Bases containing Group I, Ba2+, and Ca2+.

pH Value

• Definition: Indicates the concentration of hydrogen ions (H+) and hydroxide ions (OH-) in a solution.

• Measurement Methods:

  • Universal Indicator:

    • Paper strips soaked in indicator solution; dipped into the solution and color change is matched with a chart.

  • pH Meter and Electrode

• pH Indicators in Titration:

  • Examples:

    • Phenolphthalein: Acidic - colourless, Neutral - colourless, Alkaline - pink

    • Thymolphthalein: Acidic - colourless, Neutral - colourless, Alkaline - blue

    • Methyl Orange: Acidic - red, Neutral - orange, Alkaline - yellow

• Litmus Paper:

  • Must be damp.

  • Acidic: blue ➜ red; Alkaline: red ➜ blue

Titration

• Definition: A method used for neutralization reactions.

• Neutralization Reaction: Acid + Alkali ➜ Salt + Water

• End Point: Point in titration where the reaction is complete, indicated by a color change of a pH indicator.

• Detecting the End Point: Utilization of a pH indicator.

• Steps in Acid-Base Titration:

  1. Add 2-3 drops of pH indicator to 20 cm³ of hydrochloric acid in a conical flask.

  2. Fill a burette with 0.10 mol/dm³ NaOH and note initial volume.

  3. Add NaOH dropwise into acid, swirling to mix.

  4. Stop at color change from red to orange.

  5. Record final volume and repeat thrice for an average reading.

• Result Example: Average volume of NaOH added = 25 cm³.

• Calculate HCl Concentration.

Titration Curve

• Definition: A plot of the pH of the solution versus the volume of titrant added.

• For Strong Acid with Strong Alkali:

  • Example: HCl in conical flask, NaOH added drop by drop.

  • At the end of neutralization, pH becomes 7.

  • Volume of NaOH required for neutralization can be determined from the graph (e.g., 50 cm³).

Chemical Reactions

• Neutralization: Acid + Base ➜ Salt + Water

  • Example Reactions:

    1. 2HCl + MgO ➜ MgCl2 + H2O

    2. HCl + Mg(OH)2 ➜ MgCl2 + 2H2O

    3. HCl + NH3 ➜ NH4Cl

    4. HCl + NH4OH ➜ NH4Cl + H2O

    5. CH3COOH + MgO ➜ (CH3COO)2Mg + H2O

• Acid + Metal: Produces salt + Hydrogen gas.

  • Example:

    1. HCl + Mg ➜ MgCl2 + H2

    2. 2HNO3 + Mg ➜ Mg(NO3)2 + H2

• Acid + Carbonate: Produces salt + Water + Carbon Dioxide.

  • Example:

    1. 2HCl + CaCO3 ➜ CaCl2 + H2O + CO2

• Acid + Hydrogen Carbonate: Produces salt + Water + Carbon Dioxide.

• Ammonium Salt with Alkali: Produces salt + Water + Ammonia Gas.

  • Example: NaOH + NH4Cl ➜ NaCl + H2O + NH3

Controlling Soil Acidity

• Causes of Acidic Soil:

• Acid Rain: Rain with pH < 5

• Neutralizing Agents:

  • Limestone (CaCO3)

  • Lime (CaO): Can be strongly alkaline when dissolved; excess may cause soil to become too alkaline.

  • Ammonia Solution (NH3): Easily evaporates.

Ionic Equation for Neutralization

1. Writing Ionic Equation Steps:

   • Step 1: Write a balanced equation with state symbols.Example: NaOH (aq) + HCl (aq) ➜ NaCl (aq) + H2O (l)

   • Step 2: Split the ions of the aqueous solution. Example: Na+ + OH- + H+ + Cl- ➜ Na+ + Cl- + H2O

   • Step 3: Eliminate similar ions on both sides. Result: H+ + OH- ➜ H2O

2. Proton Donors and Acceptors:

   • H+ ion is a proton and in neutralization:

     • Acid donates protons

     • OH- ions accept protons

   • Acid = Proton donor

   • Base = Proton acceptor

   • Example Reaction:

     • CH3COOH + KOH ➜ CH3COOK + H2O:

       • H+ from ethanoic acid is donated to OH- from KOH.

       • Shows that CH3COOH is an acid and KOH is a base.

• Acid-Base Characteristics:

  • Strong or Weak Acid/Alkali determined by pH value, electrical conductivity, and rate of reaction.

Types of Oxides

• Basic Oxide (Base):

  • Reacts with acids, not alkalis, mainly insoluble in water (except Group I, Ba2+, Ca2+).

• Acidic Oxide:

  • Reacts with alkalis, formed from non-metals, all dissolve in water to form acid solutions.

• Neutral Oxide:

  • Cannot react with acids or alkalis.

    • Examples: N2O, NO, CO, H2O

• Amphoteric Oxide:

  • Can react with both acids and alkalis.

    • Examples: ZnO, Al2O3

  • Symbol Equations:

    • ZnO + 2HNO3 ➜ Zn(NO3)2 + H2O

    • Al2O3 + 6HCl ➜ 2AlCl3 + 3H2O

End of Chapter