Acids & bases

The pH Scale

• Definition:

  • The pH scale measures the acidity or alkalinity of a solution and ranges from below 0 to above 14.

  • Utilized with:

  • pH meters (electronic reading).

  • Indicators (change color depending on the pH value).

• Classification:

  • Acidic Solutions: pH < 7

  • Neutral Solutions: pH = 7

  • Alkaline Solutions: pH > 7

• pH Color Chart:

  • Colors correspond to pH numbers, facilitating identification of solution's pH.

Explaining pH

• Water Properties:

  • Neutral, existing as water molecules (H₂O).

  • Example of a non-metal oxide formed by hydrogen and oxygen.

  • Water dissociates minimally into:

  • Hydrogen ions: H^+

  • Hydroxide ions: OH^-

  H_2O
  ightleftharpoons H^+ + OH^-

• Conductivity of Water:

  • Water is a poor conductor of electricity due to a higher number of molecules compared to ions.

  • Equal concentration of H^+ and OH^- ions: Water is neutral.

• Acid/Alkali Definitions:

  • Acidic Solutions: Higher concentration of hydrogen ions ([H^+] > [OH^-]).

  • Alkaline Solutions: Higher concentration of hydroxide ions ([OH^-] > [H^+]).

Dilution Effects

• Impact of Dilution:

  • Dilution decreases concentration:

  • For acids: As diluted, pH increases (concentration of H^+ decreases).

  • For alkalis: As diluted, pH decreases (concentration of OH^- decreases).

  • Sufficient water can bring any solution to pH = 7.

Common Acids

• Formation:

  • Dissolving non-metal oxides in water (e.g., CO_2 and SO_2).

• Common Laboratory Acids:

  Name of Acid	Chemical Formula
  Hydrochloric acid	HCl
  Nitric acid	HNO₃
  Sulfuric acid	H₂SO₄
  Ethanoic acid	CH₃COOH
  Phosphoric acid	H₃PO₄

• Conductivity:

  • Acidic solutions conduct electricity due to the presence of H^+ ions.

  • Electrolysis of acidic solutions produces hydrogen gas at the negative electrode:

  • Reaction:
    2H^+(aq) + 2e^-
    ightarrow H_2(g)

  • Test: Hydrogen burns with a pop sound.

Common Alkalis

• 
  

• Formation:

  .

  • Dissolved soluble metal oxides in water. Insoluble metal oxides (e.g., copper oxide) do not affect pH.

• 
  

• Common Alkalis:
  

  Name of Alkali	Formula	Ions Present
  Lithium hydroxide	LiOH	Li⁺ and OH^-
  Sodium hydroxide	NaOH	Na⁺ and OH^-
  Potassium hydroxide	KOH	K⁺ and OH^-
  Calcium hydroxide	Ca(OH)₂	Ca²⁺ and OH^-
  Bases and Alkalis

  • Base Definition:

    • Neutralizes an acid by accepting hydrogen ions, forming water. Common bases:

    • Metal oxides

    • Metal hydroxides

    • Metal carbonates

    • Ammonia

  • Alkali Definition:

    • A soluble base (e.g., sodium hydroxide). Copper(II) oxide is a base but not an alkali (insoluble).

  • Neutralization Reaction:

    • Reaction between an acid and a base produces a salt and water:

    • ext{Acid} + ext{Base}
      ightarrow ext{Salt} + ext{Water}

  Neutralization

  • Neutralization Process:

    • Adding an alkali to an acid increases pH towards 7.

    • Adding an acid to an alkali decreases pH towards 7.

  • Chemical Reaction:

    • H^+(aq) + OH^-(aq)
      ightarrow H₂O(l)

  • Salt Formation:

    • Salts are produced when the hydrogen ions of an acid are replaced by a metal ion or ammonium ion.

  • Naming Salts:

    • Dependent on the acid and base used:

    • Example: Sodium hydroxide and nitric acid yield sodium nitrate.

  Example Salts from Reactions

  • 
    

  • Example Table of Reactions:
    

    Alkali	Acid	Salt Produced
    Sodium hydroxide	Hydrochloric acid	Sodium chloride
    Lithium hydroxide	Sulfuric acid	Lithium sulfate
    Calcium hydroxide	Nitric acid	Calcium nitrate
    Potassium hydroxide	Phosphoric acid	Potassium phosphate
    Neutralization with Insoluble Bases

    • Reaction Mechanism:

      • Metal oxides can neutralize acids; excess must be filtered away after the reaction.

      • Salt is formed through evaporation of the resulting solution.

    Ionic Equations

    • Understanding Reactions:

      • When acids react with alkalis, we focus on the reactive species.

      • Example reaction:
        ext{HCl}(aq) + ext{NaOH}(aq)
        ightarrow ext{NaCl}(aq) + H₂O(l)

    • Ionic Equation Example:

      • Both sides contain ions, removing unchanged spectator ions:

      • H^+(aq) + OH^-(aq)
        ightarrow H₂O(l)

    Titrations

    • Purpose:

      • Determines unknown concentrations of acids or bases through neutralization reactions.

      • Involves measuring solution volumes accurately, often with an indicator for color change.

    • Concordant Titre Volumes:

      • Required measurements must be within 0.2 cm³.

    • Standard Solutions:

      • Solutions with accurately known concentrations.

    Titration Equipment

    • Pipette:

      • Measures a specific liquid volume accurately.

    • Burette:

      • Measures varying volumes accurately, with graduations of 0.1 cm³, read at eye level to ensure accuracy.

    Titration Calculations

    • Neutralization Relationships:

      • ext{Concentration}{ ext{acid}} imes ext{Volume}{ ext{acid}} = ext{Concentration}{ ext{alkali}} imes ext{Volume}{ ext{alkali}}

      • Molar relationships between acid and alkali can be derived from the balanced equation.

    • Titration Example 1:

      • 12.5 cm³ of 1 mol I⁻ hydrochloric acid neutralizes 10.0 cm³ of sodium hydroxide.

      • Calculation:

      • ext{Volume}{ ext{acid}} imes ext{Concentration}{ ext{acid}} = ext{Volume}{ ext{alkali}} imes ext{Concentration}{ ext{alkali}}

      • 12.5 imes 1 = 10 imes C
        C = rac{12.5}{10.0}
        ightarrow C = 1.25 ext{ mol} ext{ dm}^{-3}

    • Titration Example 2:

      • Required volume of sulfuric acid with a concentration of 0.1 mol I⁻ to neutralize 10 cm³ lithium hydroxide.

    • Titration Example 3:

      • 10 cm³ of nitric acid neutralized by 5 cm³ of potassium hydroxide solution (0.2 mol I⁻).

      • Calculation required to find the concentration of the acid.