Domanda del libro di testo
Which of the following solutions has the greater buffer capacity: 100 mL of 0.30 M HNO2-0.30 M NaNO2 or 100 mL of 0.10 M HNO2-0.10 M NaNO2? Explain.
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Ch.17 - Applications of Aqueous Equilibria
McMurry8th EditionChemistryISBN: 9781292336145
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TOC.tableOfContents
- Ch.1 - Chemical Tools: Experimentation & Measurement170
- Ch.2 - Atoms, Molecules & Ions160
- Ch.3 - Mass Relationships in Chemical Reactions169
- Ch.4 - Reactions in Aqueous Solution199
- Ch.5 - Periodicity & Electronic Structure of Atoms102
- Ch.6 - Ionic Compounds: Periodic Trends and Bonding Theory92
- Ch.7 - Covalent Bonding and Electron-Dot Structures61
- Ch.8 - Covalent Compounds: Bonding Theories and Molecular Structure59
- Ch.9 - Thermochemistry: Chemical Energy122
- Ch.10 - Gases: Their Properties & Behavior155
- Ch.11 - Liquids & Phase Changes54
- Ch.12 - Solids and Solid-State Materials73
- Ch.13 - Solutions & Their Properties120
- Ch.14 - Chemical Kinetics98
- Ch.15 - Chemical Equilibrium125
- Ch.16 - Aqueous Equilibria: Acids & Bases110
- Ch.17 - Applications of Aqueous Equilibria147
- Ch.18 - Thermodynamics: Entropy, Free Energy & Equilibrium86
- Ch.19 - Electrochemistry154
- Ch.20 - Nuclear Chemistry96
- Ch.21 - Transition Elements and Coordination Chemistry156
- Ch.22 - The Main Group Elements187
- Ch.23 - Organic and Biological Chemistry51
Capitolo 17, Problema 60
The pH of a solution of NH3 and NH4Br is 8.90. What is the molarity of NH4Br if the molarity of NH3 is 0.016 M?
Guida verificata passo dopo passo1
Identify the relevant chemical equilibrium: \(\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-\). This is the equilibrium for ammonia acting as a base.
Use the given pH to find the pOH: \(\text{pOH} = 14 - \text{pH}\). Calculate the hydroxide ion concentration, \([\text{OH}^-]\), using \([\text{OH}^-] = 10^{-\text{pOH}}\).
Write the expression for the base dissociation constant, \(K_b\), for ammonia: \(K_b = \frac{[\text{NH}_4^+][\text{OH}^-]}{[\text{NH}_3]}\). Use the known \(K_b\) value for ammonia.
Assume that the change in concentration of \(\text{NH}_3\) due to dissociation is negligible, so \([\text{NH}_3] \approx 0.016 \text{ M}\). Substitute \([\text{OH}^-]\) and \([\text{NH}_3]\) into the \(K_b\) expression to solve for \([\text{NH}_4^+]\).
Since \([\text{NH}_4^+]\) comes from the dissociation of \(\text{NH}_4\text{Br}\), the molarity of \(\text{NH}_4\text{Br}\) is equal to \([\text{NH}_4^+]\).
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Buffer Solutions
A buffer solution is a system that resists changes in pH upon the addition of small amounts of acid or base. It typically consists of a weak acid and its conjugate base or a weak base and its conjugate acid. In this case, the NH3/NH4Br system acts as a buffer, maintaining a pH of 8.90, which is crucial for understanding the equilibrium between the ammonia and ammonium ions.
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Percorso guidato
03:02
Buffer Solutions
Henderson-Hasselbalch Equation
The Henderson-Hasselbalch equation relates the pH of a buffer solution to the concentration of its acidic and basic components. It is expressed as pH = pKa + log([A-]/[HA]), where [A-] is the concentration of the base (NH3) and [HA] is the concentration of the acid (NH4+). This equation is essential for calculating the molarity of NH4Br based on the given pH and the concentration of NH3.
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Percorso guidato
02:40Henderson-Hasselbalch Equation
Equilibrium and Ka/Kb Values
The equilibrium constant (Ka for acids and Kb for bases) quantifies the strength of an acid or base in solution. For NH3, a weak base, its Kb value can be used to find the pKa of its conjugate acid (NH4+). Understanding these values is important for determining the relationship between the concentrations of NH3 and NH4Br in the buffer solution, which ultimately helps in calculating the molarity of NH4Br.
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03:50Characteristics of Ka and Kb