Rate: pressure, surface area (AQA GCSE Chemistry): Revision Notes
Rate: pressure, surface area
How surface area affects reaction rate
When we break solid reactants into smaller pieces, we change how fast they react. This happens because of the surface area to volume ratio.
The surface area to volume ratio is a key concept in understanding reaction rates. When you increase surface area while keeping volume constant, more of the solid becomes available for reactions to occur.
Understanding surface area to volume ratio
Think of a big cube compared to lots of small cubes with the same total volume:
Worked Example: Comparing Cube Surface Areas
Big cube example:
- Each face of the cube can react with other chemicals
- A 2×2×2 cube has 6 faces, each with area 4
- Total surface area = units²
- Volume = units³
- Surface area to volume ratio =
Small cubes example:
- Eight 1×1×1 cubes have the same total volume (8 units³)
- Each small cube has 6 faces with area 1 each
- Total surface area = units²
- Surface area to volume ratio =
Why this matters for reactions
When you break a solid into smaller pieces, several important changes occur that affect the reaction rate:
- More surface gets exposed to other reactants
- Reactant particles can collide with more places on the solid
- Collisions happen more often
- The reaction goes faster
This principle explains why powdered medicines work faster than tablets, and why we crush things to make them react quicker.
How pressure affects reaction rate
Pressure changes how fast gas reactions happen by affecting molecular behaviour and collision frequency.
What happens when pressure increases
When you increase pressure on gases, the molecular dynamics change significantly:
- Gas molecules get pushed closer together
- There's less space between molecules
- Molecules bump into each other more frequently
- More collisions mean faster reactions
Think of it like a crowded room - people bump into each other more often when the room is packed.
Critical Concept: Collision Frequency
The key to understanding pressure effects is collision frequency. When gas molecules are compressed into a smaller space, they don't just get closer - they actually collide with each other much more often, which directly increases the reaction rate.
The pressure rule
In many gas reactions, there's a direct mathematical relationship:
Pressure-Rate Relationship:
- Double the pressure = Double the collision frequency
- This often means Double the reaction rate
For many simple gas reactions: Rate ∝ Pressure
You can visualise this with simple diagrams - high pressure containers have molecules packed much closer together than low pressure ones.
Worked example: marble chips and acid
Worked Example: Marble Chips and Hydrochloric Acid
The experiment:
- Marble chips (calcium carbonate) react with hydrochloric acid
- Gas (carbon dioxide) is produced and measured each minute
- Results get plotted on a rate vs. time graph
What the graphs show:
- Smaller marble chips produce gas faster at the start (steeper curve)
- Larger marble chips produce gas more slowly (gentler curve)
- Both reactions produce the same total amount of gas eventually
Why this happens:
- Smaller chips have more surface area exposed
- More acid can attack the marble simultaneously
- The reaction starts faster but still stops when one reactant runs out
Practice thinking
Practice Problem: Gas Concentration Effects
Try to work out: if you put three times as many gas molecules in the same space, how much faster might the reaction go?
Step-by-step reasoning:
- Three times more molecules in same space = three times higher concentration
- Higher concentration = more frequent collisions
- More collisions = faster reaction rate
Answer: The collision frequency increases proportionally, so the reaction could go roughly three times faster (though real reactions can be more complex due to other factors).
Key Points to Remember:
- Breaking solids into smaller pieces increases surface area - this makes reactions faster
- Higher pressure pushes gas molecules closer together - this increases collision frequency
- More frequent collisions mean faster reaction rates
- Surface area matters for solids, pressure matters for gases
- The total amount of product stays the same - you just get it faster or slower