acids bases

Unit 10: Acids & Bases

Lesson 1: Introduction to Acids & Bases

Learning Objectives:

• Compare and contrast the properties of acids and bases.

• Classify a substance as an acid or base depending upon characteristic properties or the pH of the sample.

Properties of Acids and Bases

• Acids:

  • Turn litmus paper red.

  • pH < 7.

  • React with metals and some carbonates.

  • Have a sour taste (e.g., citric acid in lemons).

  • Electrolytes (conduct electricity).

  • Concentrated acids can burn tissue.

• Bases (Alkalines):

  • Turn litmus paper blue.

  • pH > 7.

  • Have a slippery feel (e.g., soaps).

  • Taste bitter.

  • React with acids to form salts and water.

• Properties Shared by Acids and Bases:

  • Both can act as electrolytes.

  • Neutralization occurs between acids and bases, producing salt and water.

Lesson 2: Acids-Base Theories

Learning Objective:

• Describe and identify acids and bases according to the Arrhenius and alternative acid-base theory.

Acid-Base Theories

1. Arrhenius Acid:

   • A substance that yields H+ or hydrogen ions in solution.

   • Examples:

     • Hydrochloric Acid: $HCl(aq) <br>ightarrow H^+(aq) + Cl^-(aq)$.

     • Sulfuric Acid: $H<em>2SO</em>4(aq) <br>ightarrow 2H^+(aq) + SO_4^{2-}(aq)$ (diprotic).

2. Arrhenius Base:

   • A substance that yields OH- (hydroxide ions) in solution.

   • Examples:

     • Sodium Hydroxide: $NaOH(aq) <br>ightarrow Na^+(aq) + OH^-(aq)$.

     • Calcium Hydroxide: $Ca(OH)_2(aq) <br>ightarrow Ca^{2+}(aq) + 2OH^-(aq)$.

3. Formation of Hydronium Ion:

   • The hydrogen ion $H^+$ attaches to water to form the hydronium ion $H<em>3O^+$: $H^+(aq) + H</em>2O(l) <br>ightarrow H_3O^+(aq)$.

Lesson 3: The Power of Hydrogen (pH) Scale

Learning Objective:

• Relate pH to the concentration of hydronium ions.

Understanding pH

1. Definition of pH:

   • pH = power of Hydrogen.

   • LOW pH indicates higher H+, less OH- (acidic).

   • HIGH pH indicates higher OH-, less H+ (basic).

2. Mathematical Equation:

   • $pH = - ext{log}[H_3O^+]$

   • Example: If $[H_3O^+] = 0.01 M$, then (pH = 2).

   • Inversely related, each decrease of one pH unit corresponds to a tenfold increase in $[H_3O^+]$ concentration.

3. Niceties of Acidic vs. Basic Solutions:

   • Acidic solutions have more $H^+$ than $OH^-$.

   • Basic solutions have more $OH^-$ than $H^+$.

   • Neutral solutions have equal amounts of $H^+$ and $OH^-$ (pH 7 at 25°C).

Lesson 4: Indicators (Table M)

Learning Objectives:

• Identify ways in which the pH of a solution can be measured.

• Use Table M to determine the pH of a given solution.

Measuring pH

1. Methods of Measurement:

   • pH meters.

   • pH paper.

   • Acid-base indicators (dyes) that change color with pH.

2. Color Changes of Indicators:

   • For example, Methyl Orange changes from:

     • Red below pH 3.1

     • Yellow above pH 4.4

     • Orange between pH 3.3 and 4.4.

3. Color Expectations:

   • Litmus turns red below pH 4.5 and blue above pH 8.3.

Lesson 5: Neutralization Reactions

Aims:

• Write balanced chemical equations for simple neutralization reactions.

Neutralization Reactions

1. Definition:

   • Neutralization: A special type of double displacement reaction where equal quantities of an acid and a base react to produce salt and water:
     $ext{Acid} + ext{Base} <br>ightarrow ext{Salt} + ext{Water}$

   • Resulting solution has a neutral pH (pH ≈ 7).

2. Reaction Examples:

• $HCl(aq) + KOH(aq) <br>ightarrow KCl(aq) + H_2O(l)$

• $HNO<em>3(aq) + NaOH(aq) ightarrow NaNO</em>3(aq) + H_2O(l)$

Lesson 6: Titrations

Aim:

• Calculate the concentration or volume of a solution using titration data.

Titration Basics

1. Definition:

   • Titration: A laboratory technique where a solution of known concentration (titrant) is added to a known volume of a solution of unknown concentration (analyte) until neutralization occurs.

   • Equivalence Point: The moment when neutralization happens, detectable via an indicator.

2. Titration Formula:

   • $M<em>A imes V</em>A = M<em>B imes V</em>B$

   • Where:

     • $M_A$ = molarity of acid

     • $V_A$ = volume of acid

     • $M_B$ = molarity of base

     • $V_B$ = volume of base

   • For polyprotic acids (e.g., $H<em>2SO</em>4$), multiply molarity by number of H+ ions released.

3. Example Titration Problems:

   • How many mL of 2.0 M NaOH are needed to neutralize 24 mL of 1.0 M HCl?

   • Calculate molarity when 10.0 mL of an acid is needed to neutralize 30.0 mL of 0.50 M KOH.