Acids and Bases Chapter 14
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Questions and Answers

What is the formula for the equilibrium constant for a weak acid?

  • Ka = [A-] / [H+]
  • Ka = [HA] / [H+][A-]
  • Ka = [H+][A-] / [HA] (correct)
  • Ka = [OH-][B+] / [BH]

    • A buffer solution helps maintain stable pH levels.

    True

    What happens to pH when a strong acid is added to a weak base?

  • pH increases
  • pH decreases (correct)
  • pH remains constant
  • pH fluctuates wildly

    • Which of the following definitions of acids and bases is the simplest and most restrictive?

    <p>Arrhenius</p> Signup and view all the answers

    All reactions that fit the Arrhenius definition also fit the Bronsted-Lowry definition.

    <p>True</p> Signup and view all the answers

    Which species donates a proton in a proton-transfer reaction?

    <p>Acid</p> Signup and view all the answers

    A strong acid is a weak electrolyte.

    <p>False</p> Signup and view all the answers

    Match the following acids with their strength classification:

    <p>HCl = Strong Acid H2SO4 = Strong Acid HF = Weak Acid Acetic Acid (CH3COOH) = Weak Acid</p> Signup and view all the answers

    Buffers can resist changes in pH when acids or bases are added.

    <p>True</p> Signup and view all the answers

    Which of the following species acts as a Lewis base?

    <p>Electron pair donor</p> Signup and view all the answers

    In an aqueous solution at 25°C, what is the relationship between [H3O+] and [OH–]?

    <p>[H3O+] = [OH–] = 1.00 × 10−7 M</p> Signup and view all the answers

    What is the pH of a solution with a hydronium-ion concentration of 1.0 × 10−4 M?

    <p>4</p> Signup and view all the answers

    What is the relationship between pH and pOH in an aqueous solution?

    <p>pH + pOH = 14</p> Signup and view all the answers

    In an aqueous solution of a strong acid, what is the main contributor to the total [H3O+]?

    <p>The acid</p> Signup and view all the answers

    What is the hydroxide-ion concentration in a solution with a pH of 9?

    <p>1.00 × 10−9 M</p> Signup and view all the answers

    What is the normal range of pH?

    <p>0 to 14</p> Signup and view all the answers

    What is the relationship between [H3O+] and [OH–] in a basic solution?

    <p>[H3O+] &lt; [OH–]</p> Signup and view all the answers

    What is the condition for a solution to be neutral?

    <p>[H3O+] = [OH–]</p> Signup and view all the answers

    In a solution of a strong acid, what is the relationship between the concentration of hydronium ions and hydroxide ions?

    <p>The concentration of hydronium ions is inversely proportional to the concentration of hydroxide ions.</p> Signup and view all the answers

    What happens to the concentration of hydroxide ions when the concentration of hydronium ions increases in a solution?

    <p>It decreases</p> Signup and view all the answers

    What is the extent of ionization of a strong acid in water?

    <p>Near 100% ionization</p> Signup and view all the answers

    In a neutral solution, what is the relationship between the concentration of hydronium ions and hydroxide ions?

    <p>The concentration of hydronium ions is equal to the concentration of hydroxide ions.</p> Signup and view all the answers

    What is the value of Kw at 25°C?

    <p>1.0 x 10^(-14)</p> Signup and view all the answers

    If the concentration of hydronium ions in a solution is 0.10 M, what is the concentration of hydroxide ions?

    <p>6.7 x 10^(-13) M</p> Signup and view all the answers

    What is the concentration of hydronium ions in a 0.10 M HCl solution?

    <p>0.10 M</p> Signup and view all the answers

    In a solution of a strong acid, what is the relationship between the concentration of the acid and the concentration of hydronium ions?

    <p>The concentration of the acid is directly proportional to the concentration of hydronium ions.</p> Signup and view all the answers

    If the concentration of hydronium ions, [H3O+], in a solution is 1.0 x 10^-5 M, what is the concentration of hydroxide ions, [OH-]?

    <p>1.0 x 10^-9 M</p> Signup and view all the answers

    A solution with a pH of 11 is considered:

    <p>Strongly basic</p> Signup and view all the answers

    Which of the following statements correctly describes the autoionization of water?

    <p>The reaction produces equal concentrations of hydronium and hydroxide ions.</p> Signup and view all the answers

    What is the pH of a solution that has a hydroxide ion concentration of 1.0 x 10^-3 M?

    <p>11</p> Signup and view all the answers

    Which of the following will increase the pH of a solution?

    <p>Adding a strong base</p> Signup and view all the answers

    A solution has a hydronium ion concentration of 1.0 x 10^-8 M. Is the solution acidic, basic, or neutral?

    <p>Basic</p> Signup and view all the answers

    Which of the following statements about strong acids is TRUE?

    <p>Strong acids are completely ionized in solution.</p> Signup and view all the answers

    In the autoionization equilibrium of water, what happens to the concentration of hydronium ions when a strong base is added?

    <p>It decreases.</p> Signup and view all the answers

    Study Notes

    Concepts of Acids and Bases

    • Three primary definitions: Arrhenius, Brønsted-Lowry, and Lewis.
    • Arrhenius Definition: Acids produce H3O+ in water; bases produce OH− in water.
    • Brønsted-Lowry Definition: Acids donate H+, while bases accept H+ during reactions.
    • Lewis Definition: Acids accept electron pairs; bases donate electron pairs.

    Arrhenius Concept

    • Strong acids like HCl fully dissociate in water: HCl → H+ + Cl−.
    • Strong bases like NaOH fully dissociate in water: NaOH → Na+ + OH−.
    • Acid-base reactions produce water and a salt: acid + base → salt + water (exothermic).

    Brønsted-Lowry Theory

    • Key focus on H+ transfer.
    • Acid: proton donor; Base: proton acceptor.
    • Reactions involving proton transfer define acid-base interactions.

    Conjugate Acid-Base Pairs

    • Pair consists of acid and base differing by one proton.
    • Example: HCl (acid) and Cl− (conjugate base).

    Brønsted-Lowry Acids and Bases

    • Any substance with H can act as a Brønsted-Lowry acid.
    • Brønsted-Lowry bases are substances with lone pairs to accept protons.

    Lewis Acids and Bases

    • Lewis acids are electron pair acceptors (electron deficient).
    • Lewis bases are electron pair donors (must possess lone pairs).

    Strong vs Weak Acids and Bases

    • Strong acids/bases fully ionize in solution, acting as strong electrolytes.
    • Weak acids/bases partially ionize, resulting in weak electrolytes.

    Relative Strengths of Acids and Bases

    • Strongest acids have weakest conjugate bases and vice versa.
    • Equilibrium favors formation of weaker acid/base.

    Molecular Structure Influencing Acid Strength

    • Binary acids (H—Y): Acidity increases across a period due to electronegativity (e.g., HCl > HF).
    • Acidity increases down a group as H—Y bond strength decreases.

    Strengths of Oxoacids

    • Acid strength increases with electronegativity of Y in H-O-Y- compounds.
    • For oxoacids with varying O atoms, strength increases with the number of O atoms.

    Autoionization of Water

    • Pure water undergoes autoionization: 2 H2O ⇌ H3O+ + OH−.
    • Concentration of H3O+ and OH− in pure water is 1 × 10−7 M at 25 °C.

    Ion-Product Constant for Water (Kw)

    • Kw = [H3O+][OH−] = 1.0 × 10−14 at 25 °C.
    • Adding acid/base alters [H3O+] and [OH−], maintaining Kw constant but in opposite proportions.

    pH Measurement

    • pH is calculated as: pH = -log[H3O+].
    • pH < 7 indicates acidic conditions; pH > 7 indicates basic conditions; pH = 7 indicates neutrality.

    Solutions of a Strong Acid

    • Strong acids contribute significantly to H3O+ concentration in solution.
    • Calculating [OH−] involves considering autoionization: Kw = [H+][OH−].

    Solutions of a Weak Acid

    • Weak acids ionize partially, producing less H3O+.
    • Their conjugate bases can contribute to the basicity of a solution.

    Cations and Weak Acids

    • Some cations derived from weak bases can act as acids.
    • Cations from strong bases typically do not alter pH.

    Assessing Salt Solution pH

    • Neutral: salt from strong acid and strong base.
    • Basic: salt from strong base and weak acid.
    • Acidic: salt from weak base and strong acid.

    Common Ion Effect

    • Addition of a common ion shifts equilibrium to favor reactants, reducing ionization of weak acids.
    • Example: Adding HCl to acetic acid reduces H3O+ production.

    Buffer Solutions

    • Buffers resist changes in pH when acid or base is added.
    • Composed of weak acid and conjugate base or weak base and conjugate acid.
    • Example: Blood buffer system involves H2CO3 and HCO3−.### Acidic Buffer Solution
    • An acidic buffer requires significant amounts of a weak acid and its conjugate base for effective buffering.
    • Key components: acetic acid (HAc) and sodium acetate (NaAc).
    • Addition of strong acid (HCl) or strong base (NaOH) leads to neutralization reactions without significant pH changes if the quantities are appropriate.

    Buffer Dynamics

    • Strong Base Neutralization:
      • NaOH neutralizes weak acid (HC2H3O2) to form water and sodium acetate.
    • Strong Acid Neutralization:
      • HCl neutralizes conjugate base (NaC2H3O2) to form acetic acid and sodium chloride.

    Henderson-Hasselbalch Equation

    • pH = pKa + log([A-]/[HA])
    • Utilizes concentrations of the conjugate base ([A-]) and weak acid ([HA]) to calculate the pH of a buffer solution.
    • Valid under the "x is small" approximation.

    Buffer Effectiveness and Capacity

    • Effective buffers can neutralize moderate amounts of acid or base without significant pH changes.
    • Buffering range: Effective when 0.1 < [base]:[acid] < 10.
    • Buffering capacity increases with the concentration of acid and base components.

    Choosing a Buffer

    • Select an acid with a pKa close to the desired pH for optimal buffering.

    Titration Process

    • A titration involves adding a titrant of unknown concentration from a burette to a known solution until the reaction reaches the endpoint.
    • Endpoints may be indicated by pH-sensitive color changes (indicators).
    • Equivalence point occurs when moles of H3O+ equal moles of OH−.

    Titration Curves

    • Plotting pH against titrant volume shows shifts, with the inflection point marking the equivalence point.
    • pH at equivalence depends on the nature of the salt produced:
      • Neutral salt: pH = 7
      • Acidic salt: pH < 7
      • Basic salt: pH > 7

    Buffering with Different Acids and Bases

    • Buffer regions are critical for weak acids versus strong bases and vice versa, demonstrating varying pH responses to added titrant.
    • Equal concentrations of acid and base enhance buffer effectiveness.

    Summary of Buffer Behavior

    • Buffering capacity contributes significantly to pH stability.
    • Effective when both components are present in similar concentrations, optimizing the ability to neutralize added acids or bases.

    Acid-Base Chemistry Overview

    • The concentration of hydronium ions [H3O+] and hydroxide ions [OH–] in pure water is equal to 1.00 × 10−7 M at 25 °C.
    • Acidic solutions: [H3O+] > 1.00 × 10−7 M, [OH–] < 1.00 × 10−7 M.
    • Basic solutions: [OH–] > 1.00 × 10−7 M, [H3O+] < 1.00 × 10−7 M.

    pH Concept

    • pH indicates the acidity or basicity of a solution and is calculated using the formula: pH = -log[H3O+].
    • A solution with [H3O+] of 1.0 x 10-3 M has a pH of 3.
    • pH scale:
      • pH < 7 indicates acidity
      • pH > 7 indicates basicity
      • pH = 7 indicates neutrality
    • Normal pH range is 0 to 14.

    Hydroxide Ion Concentration (pOH)

    • pOH can be calculated using the formula: pOH = -log[OH–].
    • The relationship between pH and pOH is given by: pH + pOH = 14.
    • At 25 °C, the ion product of water (Kw) is: Kw = [H3O+][OH–] = 1.0 × 10−14.

    Strong Acids and Bases

    • Strong acids dissociate almost completely in water, contributing significantly to H3O+ concentration.
    • Strong bases contribute to OH– concentration similarly.
    • In a solution of a strong acid, the contribution from water is negligible as [H3O+] significantly outweighs [OH–].
    • The ion product remains constant, affecting inverse relationships between [H3O+] and [OH–].

    Examples of Strong Acids and Bases

    • For 0.10 M HCl, the hydroxide concentration can be determined using the ion product constant:
      • Kw = [H+] [OH–].
    • Calculating H3O+ and OH– concentrations for specific examples:
      • 0.15 M HNO3 leads to specific H3O+ and OH– values.
      • 0.010 M Ca(OH)2 can also be used for calculations.

    Characteristics of Acid Strength

    • Acid strength increases with the increase in atom size and decreases in H—X bond strength down a group.
    • Binary acid strength: H—C < H—N < H—O < H—F.
    • Acid strength trends for oxoacids:
      • Strength increases with the electronegativity of atom Y in H—O—Y.
      • Example: HIO < HBrO < HClO.
      • Strength increases with the number of oxygen atoms bonded to Y: HClO < HClO2 < HClO3 < HClO4.

    Autoionization of Water

    • Pure water behaves as a nonelectrolyte and minimally conducts electricity.
    • Autoionization reaction:
      • H2O(l) + H2O(l) ⇌ H3O+(aq) + OH−(aq).
    • In aqueous solutions, concentrations of H3O+ and OH– are equal: [H3O+] = [OH–] = 10−7 M at 25 °C.

    Equilibrium Concepts

    • The equilibrium constant expression for the autoionization of water can be represented as:
      • Kw = [H3O+][OH–].
    • The equilibrium constant Kw is consistently 1.0 × 10−14 at 25 °C, indicating the relationship between H3O+ and OH– concentrations in water.

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