chapter 14-Acids and bases

Chapter 14 - Acids and Bases

Dr. Geyer - AP Chemistry

Models of Acids and Bases

• Arrhenius Concept:

  • Acids produce H+ ions in a solution.

  • Bases produce OH- ions in a solution.

  • Example: What about NH3?

• Bronsted-Lowry Concept:

  • Acids are H+ donors.

  • Bases are proton acceptors.

  • Example reaction: HCl + H2O → Cl- + H3O+

Conjugate Acid/Base Pairs

• Reaction: HA (aq) + H2O (l) → H3O+ (aq) + A- (aq)

  • Conjugate Base: The remaining part of the acid molecule after it loses a proton.

  • Conjugate Acid: Formed when a proton is transferred to the base.

Equilibrium in Acid/Base Reactions

• Reaction: HA (aq) + H2O (l) ↔ H3O+ (aq) + A- (aq)

  • This illustrates an equilibrium process.

  • Identify bases involved in the reaction.

  • There is competition for H+ ions.

  • The stronger base dictates the direction of the equilibrium.

  • Direction also depends on whether the acid is strong or weak.

Acid Dissociation Constant (Ka)

• Reaction: HA (aq) + H2O (l) → H3O+ (aq) + A- (aq)

  • Ka measures the strength of the acid in a solution.

Strong Acids

• Characteristics:

  • Strong acids dissociate completely.

  • Equilibrium lies far to the right.

  • Ka is large.

  • [H+] = [HA].

  • A- is a weaker base compared to water.

  • Examples:

    • HCl, HNO3, H2SO4

Weak Acids

• Characteristics:

  • Weak acids do not dissociate completely.

  • Equilibrium lies far to the left.

  • Ka is small.

  • [H+] is much less than [HA].

  • A- is a stronger base than water.

  • Examples:

    • Acetic acid, Formic acid

Strong vs Weak Acids

• Figure 14.4:

  • Visual representation of different acid strengths in aqueous solution.

  • Strong acids: complete dissociation, producing significant H+.

  • Weak acids: partial dissociation.

Relative Acid/Base Strength Practice

• Use Table 14.2 to arrange the following bases in order of strength:

  • H2O, ClO2-, HSO4-, NH3, CN-

Types of Acids

• Polyprotic Acids:

  • Possess more than one acidic hydrogen (diprotic, triprotic, etc.).

• Oxyacids:

  • Acidic hydrogen attached to oxygen in an ion (e.g., HNO2).

• Organic Acids:

  • Contain a carbon backbone, usually with a carboxyl group –COOH; generally weak.

Water as an Acid and a Base

• Amphoteric Nature:

  • Water can act as both an acid and a base.

• Autoionization Reaction:

  • H2O (l) + H2O (l) ↔ H3O+ (aq) + OH- (aq)

  • Kw = [OH-][H3O+] = 1 x 10^-14 at 25 °C

Water and Kw Relationships

• Kw:

  • [OH-][H3O+] = 1 x 10^-14 at 25 °C

  • Neutral Solution: [OH-] = [H3O+] = 1 x 10^-7

  • Acidic Solution: [H3O+] > [OH-]

  • Basic Solution: [H3O+] < [OH-]

pH Scale

• Formula: pH = -log [H3O+]

• Ranges from 0 to 14.

• As pH decreases, [H+] increases exponentially.

• Significant figures: Only digits after the decimal of pH are significant.

pH and pOH Scale

• Scale overview:

  • pOH ranging from 0 to 14, inversely related to pH.

Relationships Between pH and Other Values

• pH = -log [H3O+]

• pOH = -log [OH-]

• 14.00 = pH + pOH

  • These equations are crucial for converting between concentrations and pH values.

pH Calculations

• Strong Acids:

  • Fully dissociated, [H+] = [HA].

  • If [HA] < 10^-7, water's contribution to H+ dominates.

• Example Calculations for pH:

  • 0.20 M HCl, 1.0 x 10^-9 M HNO3.

pH Calculations: Weak Acids

• Key Considerations:

  • Ka is generally small.

  • Identify major species to determine if H+ comes from the acid or water based on comparing Ka and Kw.

• Example:

  • Set up ICE table to analyze weak acid equilibria.

Calculating pH of Solutions

• Define change in equilibrium as x.

• Write equilibrium expressions.

• Substitute expressions and solve for pH.

Weak Acid Problems

• Tasks include calculating [H+], [OH-], pH, and pOH of specific concentrations of weak acids (e.g., 2.0 M acetic acid).

Mixture of Acids

• Analyze the dominant species.

• Stronger acid (larger Ka) determines the overall pH.

Polyprotic Acids

• Can donate multiple protons (H+), with approximate dissociation constants (Ka) significant for the first dissociation.

  • Ka1 >> Ka2 >> Ka3, often neglecting the latter ones in calculations.

Polyprotic Acids Titration Curves

• Example of pH changes during the titration of monoprotic versus diprotic acids.

Polyprotic Acid Practice

• Calculate the pH for a specific concentration of Arsenic acid with given Ka values.

Percent Dissociation of Weak Acids

• The percent dissociation increases as the acid becomes more diluted.

Percent Dissociation Practice

• Questions on calculating percent dissociation for given solutions.

Bases

• Strong Bases:

  • Complete dissociation, OH- produced in solution.

  • Example: NaOH → Na+ + OH-

• Weak Bases:

  • Minimal dissociation, often do not contain OH in their formulas.

  • Example: H3CNH2 + H2O ↔ H3CNH3+ + OH-

Base Calculations Practice

• Calculate concentrations and pH values for given weak base solutions.

Salts as Acids and Bases

• Salts are ionic compounds.

• Salts formed from strong bases and strong acids are neutral (e.g., NaCl).

Basic Salts

• Salts with neutral cations and conjugate base of weak acids produce basic solutions.

Ka and Kb Relationships

• Ka x Kb = Kw - relationship at equilibrium, applicable for calculating species concentration in salts like sodium azide.

Acidic Salts

• Salts with neutral anions and conjugate acid of weak bases produce acidic solutions.

Acidic Salts Practice

• Calculate pH of a specific concentration of acidic salts like C5H5NHClO4.

Weak Base Cation/Weak Acid Anion Relationships

• pH classification based on the dominant ions' acidic and basic properties.

Weak Base Cation/Weak Acid Anion Predictions

• Predicting acidity, basicity, or neutrality based on the comparison of Ka and Kb values of the ions involved.

Acidic/Basic Salts Practice

• Review and practice predicting pH of solutions from given salts without memorization.

Structure and Acid-Base Properties

• Two primary factors influencing acidity in binary compounds:

  1. Bond Strength: Stronger acid correlates with weaker bond strength.

  2. Bond Polarity: Greater polarity leads to stronger acidity due to electronegative atoms.

Oxyacids

• Acids like HClO and its derivatives: the strength of these acids increases with more oxygen atoms, pulling electrons away from hydrogen.

Acid-Base Properties of Oxides

• Non-metal oxides can form acids in water, and metallic oxides can form bases.