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A topic from the subject of Inorganic Chemistry in Chemistry.

  • 1 year ago
  • 6 min read

Acid-Base Concepts

Acid-base chemistry is a fundamental concept in chemistry dealing with the properties and reactions of acids and bases. Several theories exist to explain these properties and reactions, the most common being:

Arrhenius Theory

The Arrhenius theory defines acids as substances that produce hydrogen ions (H+) when dissolved in water, and bases as substances that produce hydroxide ions (OH-) when dissolved in water. This theory is limited as it only applies to aqueous solutions.

Brønsted-Lowry Theory

The Brønsted-Lowry theory expands upon the Arrhenius theory. It defines acids as proton (H+) donors and bases as proton acceptors. This theory is more general and can apply to non-aqueous solutions.

Lewis Theory

The Lewis theory provides the broadest definition of acids and bases. A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. This theory encompasses many reactions not explained by the Arrhenius or Brønsted-Lowry theories.

Key Concepts

  • pH scale: A logarithmic scale used to express the acidity or basicity of a solution. A pH of 7 is neutral, less than 7 is acidic, and greater than 7 is basic.
  • Neutralization reactions: Reactions between an acid and a base, producing salt and water.
  • Titration: A laboratory technique used to determine the concentration of an unknown acid or base using a solution of known concentration.
  • Strong acids/bases: Completely dissociate in water.
  • Weak acids/bases: Partially dissociate in water.
  • Buffers: Solutions that resist changes in pH upon the addition of small amounts of acid or base.

Understanding acid-base concepts is crucial in various fields, including medicine, environmental science, and industrial processes.

Acid-Base Concepts in Chemistry

Key Points:

  • Acids donate H+ ions (protons), while bases accept H+ ions.
  • pH scale measures acidity or basicity, ranging from 0-14:
    • 0-6: Acidic
    • 7: Neutral
    • 8-14: Basic
  • Strong acids and bases completely dissociate in water, while weak acids and bases only partially dissociate.
  • Acids and bases undergo neutralization reactions, forming salts and water.
  • Conjugate acid-base pairs differ by the exchange of a proton.

Main Concepts:

  • Arrhenius Theory: Defines acids as substances that produce H+ ions and bases as substances that produce OH- ions in water.
  • Brønsted-Lowry Theory: Defines acids as proton donors and bases as proton acceptors.
  • Lewis Theory: Defines acids as electron-pair acceptors and bases as electron-pair donors.
  • pH: Negative logarithm of the H+ ion concentration, representing acidity or basicity. A lower pH indicates a higher concentration of H+ ions and thus a stronger acid.
  • pKa: Negative logarithm of the acid dissociation constant (Ka), indicating the strength of an acid. A lower pKa indicates a stronger acid.
  • Buffer Solutions: Solutions that resist pH changes upon the addition of small amounts of acid or base. They typically consist of a weak acid and its conjugate base, or a weak base and its conjugate acid.
  • Titration: A laboratory technique used to determine the concentration of an unknown acid or base solution by reacting it with a solution of known concentration.
  • Indicators: Substances that change color depending on the pH of a solution, allowing for the visual determination of the endpoint in a titration.
Neutralization Reaction: An Acid-Base Experiment
Materials:
  • Hydrochloric acid (HCl) solution, 0.1 M
  • Sodium hydroxide (NaOH) solution, 0.1 M
  • Phenolphthalein indicator
  • 10 mL graduated cylinder
  • Erlenmeyer flask (250 mL recommended)
  • Buret (50 mL recommended)
  • Pipette (10 mL recommended)
  • Safety goggles and gloves
  • Wash bottle filled with distilled water
  • White background (e.g., white tile or paper) to better observe color change
Procedure:
  1. Wear appropriate safety gear (goggles and gloves).
  2. Clean the buret thoroughly and rinse it with a small amount of the NaOH solution before filling it.
  3. Fill the buret with the NaOH solution. Record the initial buret reading.
  4. Using a pipette, accurately measure 10 mL of the HCl solution into an Erlenmeyer flask.
  5. Add 2-3 drops of phenolphthalein indicator to the HCl solution.
  6. Slowly add the NaOH solution from the buret to the HCl solution, swirling the flask constantly.
  7. Observe the color change of the indicator against a white background.
  8. Continue adding the NaOH solution dropwise near the endpoint until the indicator turns a faint persistent pink color. This is the equivalence point.
  9. Record the final buret reading.
  10. Calculate the volume of NaOH used by subtracting the initial buret reading from the final buret reading.
Key Procedures:
  • Use precise measurements to ensure accurate results.
  • Swirl the flask constantly during the titration to ensure thorough mixing.
  • Observe the color change of the indicator carefully to determine the endpoint. The color change should persist for at least 30 seconds.
  • Rinse the inside walls of the Erlenmeyer flask with distilled water to ensure all the HCl reacts.
Calculations (Optional):

Using the volume of NaOH used and the known concentration of NaOH, the concentration of the HCl solution can be calculated using the following formula (assuming a 1:1 mole ratio between HCl and NaOH):

MHClVHCl = MNaOHVNaOH

Where:

  • MHCl = Molarity of HCl
  • VHCl = Volume of HCl used
  • MNaOH = Molarity of NaOH
  • VNaOH = Volume of NaOH used
Significance:

This experiment demonstrates the neutralization reaction between a strong acid (HCl) and a strong base (NaOH), resulting in the formation of salt (NaCl) and water. The reaction is: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l). It also introduces the concept of titration, a quantitative analytical technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

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