Table of contents

1. Chemical Measurements1h 50m
2. Tools of the Trade1h 17m
3. Experimental Error1h 52m
4 & 5. Statistics, Quality Assurance and Calibration Methods1h 57m
6. Chemical Equilibrium3h 41m
7. Activity and the Systematic Treatment of Equilibrium1h 0m
8. Monoprotic Acid-Base Equilibria1h 53m
9. Polyprotic Acid-Base Equilibria2h 17m
10. Acid-Base Titrations2h 37m
11. EDTA Titrations1h 34m
12. Advanced Topics in Equilibrium1h 16m
13. Fundamentals of Electrochemistry2h 19m
14. Electrodes and Potentiometry41m
15. Redox Titrations1h 14m
16. Electroanalytical Techniques57m
17. Fundamentals of Spectrophotometry50m

6. Chemical Equilibrium

Chemical Thermodynamics: Gibbs Free Energy

Analytical Chemistry

6. Chemical Equilibrium

Chemical Thermodynamics: Gibbs Free Energy

Previous problemNext problem

Struggling with Analytical Chemistry?

Join thousands of students who trust us to help them ace their exams!Watch the first video

Multiple Choice

Which of the following species has the greatest reducing power based on standard Gibbs free energy of formation (ΔG°f) values?

Li (s)

Cu2+ (aq)

Ag+ (aq)

Fe2+ (aq)

Previous problemNext problem

Verified step by step guidance

Step 1: Understand the concept of reducing power. Reducing power refers to the ability of a species to donate electrons in a redox reaction. A species with greater reducing power will have a more negative standard Gibbs free energy of formation (ΔG°f) for its oxidation reaction.

Step 2: Recall the relationship between Gibbs free energy and the standard electrode potential (E°). The equation ΔG° = -nFE° connects Gibbs free energy (ΔG°) to the standard electrode potential (E°), where n is the number of electrons transferred, F is the Faraday constant, and E° is the standard electrode potential. A more negative ΔG° corresponds to a higher E° for reduction reactions.

Step 3: Compare the standard Gibbs free energy of formation (ΔG°f) values for the given species. For each species, determine the ΔG°f value associated with its oxidation reaction (e.g., Li(s) → Li⁺ + e⁻). The species with the most negative ΔG°f value will have the greatest reducing power.

Step 4: Analyze the chemical nature of the species. Metals like Li(s) are typically strong reducing agents because they readily lose electrons to form cations. Compare this behavior with the other species (Cu²⁺, Ag⁺, Fe²⁺), which are ions and generally act as oxidizing agents rather than reducing agents.

Step 5: Conclude that Li(s) has the greatest reducing power based on its chemical properties and the expected ΔG°f values. This conclusion aligns with the fact that alkali metals like lithium are known for their strong reducing capabilities.