Ever wondered why some chemical reactions happen in seconds whilst... Show more
Chemistry116 views·Updated May 31, 2026·25 pagesComprehensive Higher Chemistry Unit One Notes
Catherine Closs@catieeliza
1 / 10
1
of 10
Controlling Reaction Rates
Five key factors control how fast chemical reactions happen: catalyst, concentration, particle size, pressure, and temperature. These factors all work by affecting how particles collide with each other.
Collision theory explains that particles must crash into each other to react. However, not every collision leads to a reaction - particles need enough energy to overcome the activation energy (EA), which is the minimum energy barrier that must be crossed.
For successful collisions, you need two things: particles with kinetic energy greater than the activation energy (EK > EA), and the correct collision geometry where particles hit each other at the right angle.
Key Point: Think of activation energy like a hill - particles need enough speed to get over it before they can react on the other side.
2
of 10
Temperature, Concentration, and Pressure Effects
Temperature is basically a measure of how fast particles are moving. When you heat something up, particles gain more kinetic energy and move faster. This means more particles have enough energy to overcome the activation energy barrier, leading to more successful collisions and faster reactions.
Concentration affects reaction rates because more particles in the same space means more chances for collisions. It's like a crowded room versus an empty one - you're more likely to bump into someone when there are more people around.
Pressure works similarly by squashing particles closer together. When particles are packed tighter, they collide more frequently, increasing the chances of successful reactions.
Memory Tip: Higher temperature, concentration, and pressure all increase collision frequency or energy - both speed up reactions.
3
of 10
Particle Size and Rate Calculations
Smaller particles react faster because they have a larger surface area exposed to other reactants. Think of sugar cubes versus granulated sugar dissolving in tea - the smaller granules dissolve much quicker because more surface is in contact with the water.
For rate calculations, you'll use the formula: rate = 1/t, where t is time. If a reaction takes 30 seconds, the rate is 1/30 = 0.033 s⁻¹. You can also rearrange this to find time: time = 1/rate.
The units change depending on whether time is measured in seconds, minutes, or hours. Always check your units match the question requirements.
Exam Tip: Most questions ask you to calculate time using rate = 1/t, so practise rearranging this formula until it becomes automatic.
4
of 10
Rate Experiments
A classic rate experiment involves decomposing hydrogen peroxide (H₂O₂) using different catalysts. You measure the volume of oxygen gas produced over time to work out reaction rates.
The experimental setup uses a measuring cylinder filled with water, connected to a test tube containing hydrogen peroxide and catalyst via a delivery tube. As oxygen is produced, it displaces water in the measuring cylinder.
From the data collected, you can calculate the average rate using: Average rate = ΔQ/ΔT, where ΔQ is the change in volume and ΔT is the change in time. For example, if 87 cm³ of gas is produced in 120 seconds, the average rate is 87/120 = 0.725 cm³s⁻¹.
Lab Safety: Some catalysts like MnO₂ and PbO make reactions too fast to measure safely - that's why liver is often used as a gentler biological catalyst.
5
of 10
Energy Distribution Curves
Energy distribution curves show how energy is spread among particles in a reaction mixture. These bell-shaped curves help visualise why temperature changes affect reaction rates so dramatically.
The area under the curve to the right of the activation energy line shows how many particles have enough energy to react successfully. When temperature increases, the whole curve shifts right, meaning more particles cross the energy threshold.
At higher temperatures (T₂), significantly more particles have kinetic energy greater than the activation energy. This explains why even small temperature increases can dramatically speed up reactions.
Visual Learning: These curves make it crystal clear why heating up reactions works so well - you're literally giving more particles the energy they need to react.
6
of 10
Catalysts and Energy Changes
Catalysts are reaction accelerators that speed up reactions without being consumed. They work by lowering the activation energy and providing an alternative pathway for reactants to follow.
Enthalpy (ΔH) measures the energy difference between reactants and products. Exothermic reactions release energy, so products have less energy than reactants, giving a negative ΔH value.
Potential energy diagrams show these energy changes visually. The peak represents the activated complex - an unstable intermediate formed during the reaction. The activation energy is the difference between reactants and this peak.
Remember: Catalysts lower the mountain (activation energy) but don't change the final destination (enthalpy change).
7
of 10
Exothermic vs Endothermic Reactions
Exothermic reactions release energy to the surroundings - they feel hot. The products sit at a lower energy level than the reactants, creating a negative enthalpy change. Think of burning fuel or hand warmers.
Endothermic reactions absorb energy from surroundings - they feel cold. Products have higher energy than reactants, giving a positive enthalpy change. Examples include photosynthesis or instant cold packs.
On energy diagrams, you can spot the difference immediately: exothermic reactions slope downward from reactants to products, while endothermic reactions slope upward.
Memory Hook: Exothermic = Exit (energy leaves), Endothermic = Enter (energy enters the reaction).
8
of 10
Enthalpy Calculations
Enthalpy calculations use the formula ΔH = cmΔT, where c is specific heat capacity (4.18 for water), m is mass in kg, and ΔT is temperature change.
Enthalpy of combustion measures energy released when one mole of substance burns completely in oxygen. You'll often calculate this from experimental data involving heating water.
The process involves: calculating energy transferred to water using ΔH = cmΔT, finding moles of fuel burned using moles = mass/formula mass, then dividing energy by moles to get enthalpy per mole.
Unit Check: Convert cm³ to kg by dividing by 1000, and remember enthalpy of combustion values are always negative for energy released.
9
of 10
Working Through Enthalpy Examples
Let's work through a typical problem: 3g ethanol burns, heating 100cm³ water from 20.1°C to 32.3°C. First, calculate energy transferred: ΔH = 4.18 × 0.1 × 12.2 = 5.1 kJ.
Next, find moles of ethanol burned. Ethanol (C₂H₅OH) has formula mass 46, so moles = 3/46 = 0.065 moles.
Finally, calculate enthalpy per mole: 5.1/0.065 = -78.2 kJ/mol (negative because energy is released).
Step-by-Step Success: Always work through these calculations methodically - energy first, then moles, then divide to get enthalpy per mole.
10
of 10
Advanced Enthalpy Problems
Sometimes you'll need to work backwards from enthalpy of combustion to find required fuel mass. If you know ethanol's enthalpy of combustion is -1367 kJ/mol, you can calculate how much you need to heat specific amounts of water.
For heating 100cm³ water by 10°C: energy needed = 4.18 × 0.1 × 10 = 4.18 kJ. Since 1 mole releases 1367 kJ, you need 4.18/1367 = 3 × 10⁻³ moles.
Convert moles to mass using mass = moles × formula mass. For ethanol: mass = (3 × 10⁻³) × 46 = 0.138g.
Real-World Connection: These calculations help determine fuel requirements for heating systems, camping stoves, and industrial processes.
We thought you’d never ask...
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Chemistry116 views·Updated May 31, 2026·25 pagesComprehensive Higher Chemistry Unit One Notes
Catherine Closs@catieeliza
Ever wondered why some chemical reactions happen in seconds whilst others take hours? Understanding reaction rates is crucial for everything from cooking food to manufacturing medicines, and it's easier to grasp than you might think.
1
of 10Sign up to see the content. It's free!
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Controlling Reaction Rates
Five key factors control how fast chemical reactions happen: catalyst, concentration, particle size, pressure, and temperature. These factors all work by affecting how particles collide with each other.
Collision theory explains that particles must crash into each other to react. However, not every collision leads to a reaction - particles need enough energy to overcome the activation energy (EA), which is the minimum energy barrier that must be crossed.
For successful collisions, you need two things: particles with kinetic energy greater than the activation energy (EK > EA), and the correct collision geometry where particles hit each other at the right angle.
Key Point: Think of activation energy like a hill - particles need enough speed to get over it before they can react on the other side.
2
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Temperature, Concentration, and Pressure Effects
Temperature is basically a measure of how fast particles are moving. When you heat something up, particles gain more kinetic energy and move faster. This means more particles have enough energy to overcome the activation energy barrier, leading to more successful collisions and faster reactions.
Concentration affects reaction rates because more particles in the same space means more chances for collisions. It's like a crowded room versus an empty one - you're more likely to bump into someone when there are more people around.
Pressure works similarly by squashing particles closer together. When particles are packed tighter, they collide more frequently, increasing the chances of successful reactions.
Memory Tip: Higher temperature, concentration, and pressure all increase collision frequency or energy - both speed up reactions.
3
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Particle Size and Rate Calculations
Smaller particles react faster because they have a larger surface area exposed to other reactants. Think of sugar cubes versus granulated sugar dissolving in tea - the smaller granules dissolve much quicker because more surface is in contact with the water.
For rate calculations, you'll use the formula: rate = 1/t, where t is time. If a reaction takes 30 seconds, the rate is 1/30 = 0.033 s⁻¹. You can also rearrange this to find time: time = 1/rate.
The units change depending on whether time is measured in seconds, minutes, or hours. Always check your units match the question requirements.
Exam Tip: Most questions ask you to calculate time using rate = 1/t, so practise rearranging this formula until it becomes automatic.
4
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Rate Experiments
A classic rate experiment involves decomposing hydrogen peroxide (H₂O₂) using different catalysts. You measure the volume of oxygen gas produced over time to work out reaction rates.
The experimental setup uses a measuring cylinder filled with water, connected to a test tube containing hydrogen peroxide and catalyst via a delivery tube. As oxygen is produced, it displaces water in the measuring cylinder.
From the data collected, you can calculate the average rate using: Average rate = ΔQ/ΔT, where ΔQ is the change in volume and ΔT is the change in time. For example, if 87 cm³ of gas is produced in 120 seconds, the average rate is 87/120 = 0.725 cm³s⁻¹.
Lab Safety: Some catalysts like MnO₂ and PbO make reactions too fast to measure safely - that's why liver is often used as a gentler biological catalyst.
5
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Energy Distribution Curves
Energy distribution curves show how energy is spread among particles in a reaction mixture. These bell-shaped curves help visualise why temperature changes affect reaction rates so dramatically.
The area under the curve to the right of the activation energy line shows how many particles have enough energy to react successfully. When temperature increases, the whole curve shifts right, meaning more particles cross the energy threshold.
At higher temperatures (T₂), significantly more particles have kinetic energy greater than the activation energy. This explains why even small temperature increases can dramatically speed up reactions.
Visual Learning: These curves make it crystal clear why heating up reactions works so well - you're literally giving more particles the energy they need to react.
6
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Catalysts and Energy Changes
Catalysts are reaction accelerators that speed up reactions without being consumed. They work by lowering the activation energy and providing an alternative pathway for reactants to follow.
Enthalpy (ΔH) measures the energy difference between reactants and products. Exothermic reactions release energy, so products have less energy than reactants, giving a negative ΔH value.
Potential energy diagrams show these energy changes visually. The peak represents the activated complex - an unstable intermediate formed during the reaction. The activation energy is the difference between reactants and this peak.
Remember: Catalysts lower the mountain (activation energy) but don't change the final destination (enthalpy change).
7
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Exothermic vs Endothermic Reactions
Exothermic reactions release energy to the surroundings - they feel hot. The products sit at a lower energy level than the reactants, creating a negative enthalpy change. Think of burning fuel or hand warmers.
Endothermic reactions absorb energy from surroundings - they feel cold. Products have higher energy than reactants, giving a positive enthalpy change. Examples include photosynthesis or instant cold packs.
On energy diagrams, you can spot the difference immediately: exothermic reactions slope downward from reactants to products, while endothermic reactions slope upward.
Memory Hook: Exothermic = Exit (energy leaves), Endothermic = Enter (energy enters the reaction).
8
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Enthalpy Calculations
Enthalpy calculations use the formula ΔH = cmΔT, where c is specific heat capacity (4.18 for water), m is mass in kg, and ΔT is temperature change.
Enthalpy of combustion measures energy released when one mole of substance burns completely in oxygen. You'll often calculate this from experimental data involving heating water.
The process involves: calculating energy transferred to water using ΔH = cmΔT, finding moles of fuel burned using moles = mass/formula mass, then dividing energy by moles to get enthalpy per mole.
Unit Check: Convert cm³ to kg by dividing by 1000, and remember enthalpy of combustion values are always negative for energy released.
9
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Working Through Enthalpy Examples
Let's work through a typical problem: 3g ethanol burns, heating 100cm³ water from 20.1°C to 32.3°C. First, calculate energy transferred: ΔH = 4.18 × 0.1 × 12.2 = 5.1 kJ.
Next, find moles of ethanol burned. Ethanol (C₂H₅OH) has formula mass 46, so moles = 3/46 = 0.065 moles.
Finally, calculate enthalpy per mole: 5.1/0.065 = -78.2 kJ/mol (negative because energy is released).
Step-by-Step Success: Always work through these calculations methodically - energy first, then moles, then divide to get enthalpy per mole.
10
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Advanced Enthalpy Problems
Sometimes you'll need to work backwards from enthalpy of combustion to find required fuel mass. If you know ethanol's enthalpy of combustion is -1367 kJ/mol, you can calculate how much you need to heat specific amounts of water.
For heating 100cm³ water by 10°C: energy needed = 4.18 × 0.1 × 10 = 4.18 kJ. Since 1 mole releases 1367 kJ, you need 4.18/1367 = 3 × 10⁻³ moles.
Convert moles to mass using mass = moles × formula mass. For ethanol: mass = (3 × 10⁻³) × 46 = 0.138g.
Real-World Connection: These calculations help determine fuel requirements for heating systems, camping stoves, and industrial processes.
We thought you’d never ask...
What is the Knowunity AI companion?
Our AI Companion is a student-focused AI tool that offers more than just answers. Built on millions of Knowunity resources, it provides relevant information, personalised study plans, quizzes, and content directly in the chat, adapting to your individual learning journey.
Where can I download the Knowunity app?
You can download the app from Google Play Store and Apple App Store.
Is Knowunity really free of charge?
That's right! Enjoy free access to study content, connect with fellow students, and get instant help – all at your fingertips.
Similar content
Most popular content: Enthalpy Change (δh)
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Explore the various types of enthalpy changes in A-Level Chemistry, including reaction, combustion, formation, and lattice energies. This summary covers key concepts such as Hess's Law, standard conditions, and the factors affecting ionic size. Ideal for WJEC Unit 3 students seeking to grasp the fundamentals of thermodynamics in chemistry.
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Wow, I am really amazed. I just tried the app because I've seen it advertised many times and was absolutely stunned. This app is THE HELP you want for school and above all, it offers so many things, such as workouts and fact sheets, which have been VERY helpful to me personally.
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