Acids and Bases: Properties and Theories

• Acids and bases are fundamental chemical concepts

Properties of Acids:

• Acids taste sour
• Acids turn blue litmus paper red
• Acids react with metals to produce hydrogen gas
• Acids conduct electricity in solution
• Acids donate H⁺ ions, according to the Arrhenius & Brønsted-Lowry definitions

Properties of Bases:

• Bases taste bitter
• Bases feel slippery
• Bases turn red litmus paper blue
• Bases conduct electricity in solution
• Bases accept H⁺ ions, according to the Brønsted-Lowry definition

Acid-Base Theories:

• Arrhenius Definition: Acids produce H⁺ ions in solution, while bases produce OH⁻ ions in solution
• Brønsted-Lowry Definition: Acids are proton (H⁺) donors, whereas bases are proton (H⁺) acceptors
• Lewis Definition: Acids accept an electron pair, while bases donate an electron pair

Common Acids & Bases:

• Hydrochloric Acid (HCl) is a common acid
• Sulfuric Acid (H₂SO₄) is a common acid
• Nitric Acid (HNO₃) is a common acid
• Acetic Acid (CH₃COOH) is a common acid
• Sodium Hydroxide (NaOH) is a common base
• Potassium Hydroxide (KOH) is a common base
• Ammonia (NH₃) is a common base
• Calcium Hydroxide (Ca(OH)₂) is a common base

Strengths of Acids and Bases:

• Strong acids completely ionize in water (e.g., HCl, HNO₃, H₂SO₄)
• Weak acids only partially ionize (e.g., CH₃COOH)
• Strong bases completely dissociate in water (e.g., NaOH, KOH)
• Weak bases only partially dissociate (e.g., NH₃)

Dissociation Equations:

• Strong Acid: ( HCl \rightarrow H^+ + Cl^- )
• Weak Acid: ( CH_3COOH \rightleftharpoons H^+ + CH_3COO^- )
• Strong Base: ( NaOH \rightarrow Na^+ + OH^- )
• Weak Base: ( NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^- )

Equilibrium Constant (Ka and Kb):

• A large ( K_a ) indicates a stronger acid
• A large ( K_b ) indicates a stronger base

Problem-Solving for Weak Acids:

• To find the dissociation constant: ( K_a = \frac{[H^+][A^-]}{[HA]} )
• Steps to calculate the pH of a weak acid:
  • Write the dissociation equation
  • Use an ICE table (Initial, Change, Equilibrium)
  • Solve for H⁺ concentration using Ka
  • Calculate pH using ( pH = -\log[H^+] )

pH & pOH Scale:

• pH measures H⁺ concentration: ( pH = -\log [H^+] )
• pOH measures OH⁻ concentration: ( pOH = -\log [OH^-] )
• Relationship between pH and pOH: ( pH + pOH = 14 )

Neutral, Acidic, and Basic Solutions:

• Acidic solutions have a pH < 7
• Neutral solutions have a pH = 7
• Basic solutions have a pH > 7

Ion Product Constant for Water:

• ( K_w = [H^+][OH^-] = 1.0 \times 10^{-14} )

Problem-Solving with pH and pOH:

• To find pH given H⁺ concentration: ( pH = -\log(1.0 \times 10^{-3}) = 3 ) if ([H^+] = 1.0 \times 10^{-3})
• To find pOH given OH⁻ concentration: ( pOH = -\log(2.0 \times 10^{-5}) \approx 4.7 ) if ([OH^-] = 2.0 \times 10^{-5})
• To find H⁺ or OH⁻ given pH or pOH:
  • ([H^+] = 10^{-pH})
  • ([OH^-] = 10^{-pOH})

Neutralization Reaction:

• Acid + Base → Salt + Water
• Example: ( HCl + NaOH \rightarrow NaCl + H_2O )

Titration:

• Titration is used to determine an unknown concentration
• Titration uses a burette to add a base to an acid (or vice versa)
• Equivalence Point: Moles of ( H^+ ) = Moles of ( OH^- )

Titration Formula:

• ( M_a V_a = M_b V_b )
  • ( M_a ) = molarity of acid
  • ( V_a ) = volume of acid
  • ( M_b ) = molarity of base
  • ( V_b ) = volume of base