This section covers how fast reactions happen (rates) and the factors that affect them, including a key concept: reversible reactions and chemical equilibrium.
Rate = amount of product formed / time (or amount of reactant used / time). Measured by: volume of gas collected, mass loss, colour change, precipitate formation. Faster rate shown by steeper gradient on graph. Factors: temperature, concentration, surface area, catalysts.
Particles must collide with sufficient energy (≥ activation energy) and correct orientation to react. Increasing temperature: particles move faster, more frequent collisions AND more particles exceed activation energy. Increasing concentration: more particles per unit volume → more frequent collisions. Increasing surface area (smaller pieces): more exposed particles → more collisions. Catalysts: lower activation energy → more particles can react.
Reversible reactions: can go forwards and backwards (⇌). Dynamic equilibrium: rate of forward reaction = rate of backward reaction, concentrations remain constant. Le Chatelier\'s principle (Supplement): if conditions change, equilibrium shifts to oppose the change. Increase temperature → shifts towards endothermic direction. Increase concentration → shifts away from added substance. Pressure increase → shifts to side with fewer gas moles. Catalysts do NOT change equilibrium position (speed up both directions equally).
A catalyst provides an alternative pathway with lower activation energy, speeding up the reaction. However, it lowers the activation energy equally for both the forward and backward reactions. Since both rates increase by the same factor, the equilibrium position is unchanged — equilibrium is reached faster, but the final ratio of products to reactants is the same. This is why industrial processes use catalysts to reach equilibrium quickly, not to shift it.
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