Questions and Answers
What is the formula for the equilibrium constant for a weak acid?
A buffer solution helps maintain stable pH levels.
True
What happens to pH when a strong acid is added to a weak base?
Which of the following definitions of acids and bases is the simplest and most restrictive?
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All reactions that fit the Arrhenius definition also fit the Bronsted-Lowry definition.
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Which species donates a proton in a proton-transfer reaction?
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A strong acid is a weak electrolyte.
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Match the following acids with their strength classification:
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Buffers can resist changes in pH when acids or bases are added.
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Which of the following species acts as a Lewis base?
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In an aqueous solution at 25°C, what is the relationship between [H3O+] and [OH–]?
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What is the pH of a solution with a hydronium-ion concentration of 1.0 × 10−4 M?
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What is the relationship between pH and pOH in an aqueous solution?
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In an aqueous solution of a strong acid, what is the main contributor to the total [H3O+]?
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What is the hydroxide-ion concentration in a solution with a pH of 9?
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What is the normal range of pH?
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What is the relationship between [H3O+] and [OH–] in a basic solution?
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What is the condition for a solution to be neutral?
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In a solution of a strong acid, what is the relationship between the concentration of hydronium ions and hydroxide ions?
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What happens to the concentration of hydroxide ions when the concentration of hydronium ions increases in a solution?
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What is the extent of ionization of a strong acid in water?
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In a neutral solution, what is the relationship between the concentration of hydronium ions and hydroxide ions?
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What is the value of Kw at 25°C?
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If the concentration of hydronium ions in a solution is 0.10 M, what is the concentration of hydroxide ions?
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What is the concentration of hydronium ions in a 0.10 M HCl solution?
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In a solution of a strong acid, what is the relationship between the concentration of the acid and the concentration of hydronium ions?
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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-]?
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A solution with a pH of 11 is considered:
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Which of the following statements correctly describes the autoionization of water?
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What is the pH of a solution that has a hydroxide ion concentration of 1.0 x 10^-3 M?
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Which of the following will increase the pH of a solution?
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A solution has a hydronium ion concentration of 1.0 x 10^-8 M. Is the solution acidic, basic, or neutral?
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Which of the following statements about strong acids is TRUE?
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In the autoionization equilibrium of water, what happens to the concentration of hydronium ions when a strong base is added?
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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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Learn about strong acids and bases, Arrhenius, Bronsted-Lowry, and Lewis definitions. Identify common acids and bases with examples.
