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THE VISUAL LEARNING STUDIO
See the idea. Understand the connection.
35 visual explanations across 29 chapters, with topic maps throughout the library.
The chapter index will grow as additional notes are prepared. Chapters shown here are a learning collection, not a claim that the entire exam syllabus is complete.
SEE IT. CONNECT IT. EXPLAIN IT.
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Study the expanded notes and the credited detailed textbook lessons. The current chapter index covers core topics; it is not a claim of exhaustive coverage of every examination syllabus. Check the official syllabus and primary references.
Equilibrium — lesson overview
Dynamic equilibrium
Chemical equilibrium in a closed system is dynamic: forward and reverse reactions continue while their rates are equal at equilibrium.
Constant macroscopic composition does not mean all reactions have stopped or all reactant and product concentrations are equal.
The equilibrium constant relates equilibrium activities, approximated by suitable concentration or pressure expressions in introductory settings, at a stated temperature.
Response to changes
Changing conditions can disturb equilibrium and favour a response described qualitatively by Le Chatelier’s principle within its domain.
Temperature changes can alter the equilibrium constant. A catalyst helps the system approach equilibrium faster without changing its equilibrium constant at the same temperature.
Pressure effects depend on the reaction and gaseous stoichiometry; increasing pressure is not a universal method for increasing product yield.
Acid–base foundations
For dilute aqueous solutions in the usual introductory approximation, pH = −log₁₀[H⁺]. The more general definition uses activity.
A lower pH indicates greater hydrogen-ion activity. Acid strength, concentration and total amount are different quantities.
For example, an idealised [H⁺] of 10⁻³ mol L⁻¹ gives pH 3; increasing it tenfold lowers pH by one unit.
