Ever wondered why some reactions happen lightning-fast whilst others take...
Chemistry477 views·Updated 22 Jul 2026·2 pagesUnderstanding Chemical Reactions: Rate and Extent
Evelyn Ridley@ev_alice
1
of 2
Measuring and Controlling Reaction Rates
Understanding reaction rates is like being a chemical detective - you're measuring how quickly reactants disappear or products appear. The maths is straightforward: rate equals the quantity of substance divided by time, giving you units like g/s or cm³/s.
There are three main ways to track reactions in the lab. You can measure decreasing mass (like when gas escapes), increasing gas volume, or decreasing light passing through a solution (the classic disappearing cross experiment). Each method gives you real data about how fast your reaction is proceeding.
The secret behind all reactions lies in particle collisions. Think of it like a game of molecular snooker - particles must crash into each other with enough energy (called activation energy) to actually react. Without enough energy, they just bounce off each other uselessly.
Five key factors control how fast reactions happen: temperature, concentration, pressure (for gases), surface area, and catalysts. Master these, and you can predict and control virtually any chemical reaction.
Quick Check: Remember that smaller pieces of solid reactants mean larger surface area, which means more collisions and faster reactions!
Catalysts are absolute game-changers in chemistry. They speed up reactions without getting used up themselves, often by lowering the activation energy barrier. Industries love them because they save time, money, and energy - platinum in car exhausts and iron in ammonia production are perfect examples.
2
of 2
Reversible Reactions and Chemical Equilibrium
Some reactions are like chemical see-saws - they can go forwards and backwards depending on conditions. Reversible reactions use the ⇌ symbol and can completely change direction, like litmus paper switching from red to blue and back again.
When a reversible reaction reaches equilibrium, both forward and reverse reactions keep happening at exactly the same rate. The concentrations stay constant, but don't assume this means equal amounts - equilibrium can favour either reactants or products.
Le Chatelier's Principle is your secret weapon for predicting what happens when you change conditions. If you increase temperature, pressure, or concentration, the equilibrium shifts to counteract that change. It's like the reaction is trying to restore balance.
Temperature changes affect equilibrium based on whether reactions are endothermic or exothermic. For pressure changes in gaseous reactions, equilibrium shifts towards the side with fewer gas molecules. Concentration changes push the equilibrium towards whichever side you're removing substances from.
Exam Tip: Always count the number of gas molecules on each side of the equation when predicting pressure effects - this is a common exam question!
The classic example of hydrated copper sulfate demonstrates energy conservation perfectly. The energy absorbed when heating the blue crystals to white powder equals exactly the energy released when adding water back - no energy is ever lost or gained overall.
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Chemistry477 views·Updated 22 Jul 2026·2 pagesUnderstanding Chemical Reactions: Rate and Extent
Evelyn Ridley@ev_alice
Ever wondered why some reactions happen lightning-fast whilst others take ages? This topic explores how to measure and control reaction rates, plus how some reactions can actually reverse themselves. You'll discover the secrets behind making reactions speed up or slow...
1
of 2
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Measuring and Controlling Reaction Rates
Understanding reaction rates is like being a chemical detective - you're measuring how quickly reactants disappear or products appear. The maths is straightforward: rate equals the quantity of substance divided by time, giving you units like g/s or cm³/s.
There are three main ways to track reactions in the lab. You can measure decreasing mass (like when gas escapes), increasing gas volume, or decreasing light passing through a solution (the classic disappearing cross experiment). Each method gives you real data about how fast your reaction is proceeding.
The secret behind all reactions lies in particle collisions. Think of it like a game of molecular snooker - particles must crash into each other with enough energy (called activation energy) to actually react. Without enough energy, they just bounce off each other uselessly.
Five key factors control how fast reactions happen: temperature, concentration, pressure (for gases), surface area, and catalysts. Master these, and you can predict and control virtually any chemical reaction.
Quick Check: Remember that smaller pieces of solid reactants mean larger surface area, which means more collisions and faster reactions!
Catalysts are absolute game-changers in chemistry. They speed up reactions without getting used up themselves, often by lowering the activation energy barrier. Industries love them because they save time, money, and energy - platinum in car exhausts and iron in ammonia production are perfect examples.
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Reversible Reactions and Chemical Equilibrium
Some reactions are like chemical see-saws - they can go forwards and backwards depending on conditions. Reversible reactions use the ⇌ symbol and can completely change direction, like litmus paper switching from red to blue and back again.
When a reversible reaction reaches equilibrium, both forward and reverse reactions keep happening at exactly the same rate. The concentrations stay constant, but don't assume this means equal amounts - equilibrium can favour either reactants or products.
Le Chatelier's Principle is your secret weapon for predicting what happens when you change conditions. If you increase temperature, pressure, or concentration, the equilibrium shifts to counteract that change. It's like the reaction is trying to restore balance.
Temperature changes affect equilibrium based on whether reactions are endothermic or exothermic. For pressure changes in gaseous reactions, equilibrium shifts towards the side with fewer gas molecules. Concentration changes push the equilibrium towards whichever side you're removing substances from.
Exam Tip: Always count the number of gas molecules on each side of the equation when predicting pressure effects - this is a common exam question!
The classic example of hydrated copper sulfate demonstrates energy conservation perfectly. The energy absorbed when heating the blue crystals to white powder equals exactly the energy released when adding water back - no energy is ever lost or gained overall.
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