CHM2046L Final Exam UF Spring 2026 Study Guide
Course Overview
CHM2046L is the General Chemistry II Laboratory course at the University of Florida. This guide
is designed to help you prepare for the final exam by reviewing key concepts, providing practice
questions, and summarizing important experiments.
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Core Lab Experiments
1. Beer’s Law & Spectrophotometry – Calibration curves, absorption spectroscopy, and
concentration determination.
2. Chemical Kinetics – Reaction rates, rate laws, pseudo‑first‑order conditions, temperature
dependence.
3. Chemical Equilibrium – The iron thiocyanate system, equilibrium constant (Kc) determination.
4. Acids, Bases, & Buffers – pH measurements, buffer preparation, weak acid titrations.
5. Solubility Equilibria – Ksp determination, common ion effect.
6. Electrochemistry – Voltaic cells, the Nernst equation, concentration cells.
7. Thermodynamics – ∆G°, ∆H°, ∆S° relationships, van’t Hoff plots.
8. Transition Metal Complexes – Ligand field theory, color, and absorption.
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Key Formulas & Concepts
,Beer‑Lambert Law
\[
A = \varepsilon c d
\]
- A = absorbance (unitless)
- ε = molar absorptivity (L mol⁻¹ cm⁻¹)
- c = concentration (mol/L)
- d = path length (cm)
Calibration curve: \( A = m \cdot c + b \), with \( m = \varepsilon d \). The y‑intercept should be
near zero; otherwise, there is systematic error.
Kinetics
- Rate law: rate = \( k [A]^m [B]^n \)
- Pseudo‑first‑order: when one reactant is in large excess, rate = \( k_{obs} [B]^n \)
- First‑order integrated rate law: \(\ln[A]_t = -kt + \ln[A]_0\)
- Second‑order integrated rate law: \(\frac{1}{[A]_t} = kt + \frac{1}{[A]_0}\)
Equilibrium
- Equilibrium constant: \( K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b} \) for \( aA + bB \rightleftharpoons
cC + dD \)
- Reaction quotient (Q): same form as Kc but with initial concentrations
- ICE tables: used to solve for equilibrium concentrations
Acids, Bases & Buffers
- Henderson‑Hasselbalch equation: \( pH = pK_a + \log\left(\frac{[base]}{[acid]}\right) \)
- At half‑equivalence point: \( pH = pK_a \)
, - Buffer capacity: greatest when \([base]/[acid] = 1\)
Solubility Equilibria
- Ksp expression: \( K_{sp} = [M^{n+}][X^{m-}] \) (for \( M_aX_b \))
- Common ion effect: adding a common ion decreases solubility
Electrochemistry
- Nernst equation: \( E = E^\circ - \frac{0.0592}{n} \log Q \) (at 25°C)
- Concentration cells: \( E = -\frac{0.0592}{n} \log\left(\frac{[dilute]}{[conc]}\right) \)
- ΔG° = –nFE° (F = 96,485 C/mol)
Thermodynamics
- ΔG° = –RT ln K (R = 8.314 J mol⁻¹ K⁻¹)
- van’t Hoff equation: \( \ln K = -\frac{\Delta H^\circ}{R} \cdot \frac{1}{T} + \frac{\Delta
S^\circ}{R} \)
---
100 Practice Questions with Answers & Rationales
Beer’s Law & Spectrophotometry (1–15)
1. Explain Beer's Law and identify each component in the equation A = ε d c.
- A: absorbance (unitless)
- ε: molar absorptivity (L mol⁻¹ cm⁻¹)
- d: path length of the cuvette (cm)
Course Overview
CHM2046L is the General Chemistry II Laboratory course at the University of Florida. This guide
is designed to help you prepare for the final exam by reviewing key concepts, providing practice
questions, and summarizing important experiments.
---
Core Lab Experiments
1. Beer’s Law & Spectrophotometry – Calibration curves, absorption spectroscopy, and
concentration determination.
2. Chemical Kinetics – Reaction rates, rate laws, pseudo‑first‑order conditions, temperature
dependence.
3. Chemical Equilibrium – The iron thiocyanate system, equilibrium constant (Kc) determination.
4. Acids, Bases, & Buffers – pH measurements, buffer preparation, weak acid titrations.
5. Solubility Equilibria – Ksp determination, common ion effect.
6. Electrochemistry – Voltaic cells, the Nernst equation, concentration cells.
7. Thermodynamics – ∆G°, ∆H°, ∆S° relationships, van’t Hoff plots.
8. Transition Metal Complexes – Ligand field theory, color, and absorption.
---
Key Formulas & Concepts
,Beer‑Lambert Law
\[
A = \varepsilon c d
\]
- A = absorbance (unitless)
- ε = molar absorptivity (L mol⁻¹ cm⁻¹)
- c = concentration (mol/L)
- d = path length (cm)
Calibration curve: \( A = m \cdot c + b \), with \( m = \varepsilon d \). The y‑intercept should be
near zero; otherwise, there is systematic error.
Kinetics
- Rate law: rate = \( k [A]^m [B]^n \)
- Pseudo‑first‑order: when one reactant is in large excess, rate = \( k_{obs} [B]^n \)
- First‑order integrated rate law: \(\ln[A]_t = -kt + \ln[A]_0\)
- Second‑order integrated rate law: \(\frac{1}{[A]_t} = kt + \frac{1}{[A]_0}\)
Equilibrium
- Equilibrium constant: \( K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b} \) for \( aA + bB \rightleftharpoons
cC + dD \)
- Reaction quotient (Q): same form as Kc but with initial concentrations
- ICE tables: used to solve for equilibrium concentrations
Acids, Bases & Buffers
- Henderson‑Hasselbalch equation: \( pH = pK_a + \log\left(\frac{[base]}{[acid]}\right) \)
- At half‑equivalence point: \( pH = pK_a \)
, - Buffer capacity: greatest when \([base]/[acid] = 1\)
Solubility Equilibria
- Ksp expression: \( K_{sp} = [M^{n+}][X^{m-}] \) (for \( M_aX_b \))
- Common ion effect: adding a common ion decreases solubility
Electrochemistry
- Nernst equation: \( E = E^\circ - \frac{0.0592}{n} \log Q \) (at 25°C)
- Concentration cells: \( E = -\frac{0.0592}{n} \log\left(\frac{[dilute]}{[conc]}\right) \)
- ΔG° = –nFE° (F = 96,485 C/mol)
Thermodynamics
- ΔG° = –RT ln K (R = 8.314 J mol⁻¹ K⁻¹)
- van’t Hoff equation: \( \ln K = -\frac{\Delta H^\circ}{R} \cdot \frac{1}{T} + \frac{\Delta
S^\circ}{R} \)
---
100 Practice Questions with Answers & Rationales
Beer’s Law & Spectrophotometry (1–15)
1. Explain Beer's Law and identify each component in the equation A = ε d c.
- A: absorbance (unitless)
- ε: molar absorptivity (L mol⁻¹ cm⁻¹)
- d: path length of the cuvette (cm)
