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30 Dec 2025
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Comprehensive Higher Chemistry Unit One Notes
Catherine Closs
@catieeliza
1 / 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.
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.
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.
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.
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.
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).
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).
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.
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.
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.
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Thomas R
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Just amazing. Let's me revise 10x better, this app is a quick 10/10. I highly recommend it to anyone. I can watch and search for notes. I can save them in the subject folder. I can revise it any time when I come back. If you haven't tried this app, you're really missing out.
Basil
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This app has made me feel so much more confident in my exam prep, not only through boosting my own self confidence through the features that allow you to connect with others and feel less alone, but also through the way the app itself is centred around making you feel better. It is easy to navigate, fun to use, and helpful to anyone struggling in absolutely any way.
David K
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Sudenaz Ocak
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I know a lot of apps use fake accounts to boost their reviews but this app deserves it all. Originally I was getting 4 in my English exams and this time I got a grade 7. I didn’t even know about this app three days until the exam and it has helped A LOT. Please actually trust me and use it as I’m sure you too will see developments.
Xander S
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Elisha
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This apps acc the goat. I find revision so boring but this app makes it so easy to organize it all and then you can ask the freeeee ai to test yourself so good and you can easily upload your own stuff. highly recommend as someone taking mocks now
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Comprehensive 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.
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.
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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.
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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.
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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.
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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.
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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).
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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).
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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.
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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.
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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.
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Stefan S
iOS user
This app is really great. There are so many study notes and help [...]. My problem subject is French, for example, and the app has so many options for help. Thanks to this app, I have improved my French. I would recommend it to anyone.
Samantha Klich
Android user
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.
Anna
iOS user
Best app on earth! no words because it’s too good
Thomas R
iOS user
Just amazing. Let's me revise 10x better, this app is a quick 10/10. I highly recommend it to anyone. I can watch and search for notes. I can save them in the subject folder. I can revise it any time when I come back. If you haven't tried this app, you're really missing out.
Basil
Android user
This app has made me feel so much more confident in my exam prep, not only through boosting my own self confidence through the features that allow you to connect with others and feel less alone, but also through the way the app itself is centred around making you feel better. It is easy to navigate, fun to use, and helpful to anyone struggling in absolutely any way.
David K
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The app's just great! All I have to do is enter the topic in the search bar and I get the response real fast. I don't have to watch 10 YouTube videos to understand something, so I'm saving my time. Highly recommended!
Sudenaz Ocak
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In school I was really bad at maths but thanks to the app, I am doing better now. I am so grateful that you made the app.
Greenlight Bonnie
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very reliable app to help and grow your ideas of Maths, English and other related topics in your works. please use this app if your struggling in areas, this app is key for that. wish I'd of done a review before. and it's also free so don't worry about that.
Rohan U
Android user
I know a lot of apps use fake accounts to boost their reviews but this app deserves it all. Originally I was getting 4 in my English exams and this time I got a grade 7. I didn’t even know about this app three days until the exam and it has helped A LOT. Please actually trust me and use it as I’m sure you too will see developments.
Xander S
iOS user
THE QUIZES AND FLASHCARDS ARE SO USEFUL AND I LOVE THE SCHOOLGPT. IT ALSO IS LITREALLY LIKE CHATGPT BUT SMARTER!! HELPED ME WITH MY MASCARA PROBLEMS TOO!! AS WELL AS MY REAL SUBJECTS ! DUHHH 😍😁😲🤑💗✨🎀😮
Elisha
iOS user
This apps acc the goat. I find revision so boring but this app makes it so easy to organize it all and then you can ask the freeeee ai to test yourself so good and you can easily upload your own stuff. highly recommend as someone taking mocks now
Paul T
iOS user
The app is very easy to use and well designed. I have found everything I was looking for so far and have been able to learn a lot from the presentations! I will definitely use the app for a class assignment! And of course it also helps a lot as an inspiration.
Stefan S
iOS user
This app is really great. There are so many study notes and help [...]. My problem subject is French, for example, and the app has so many options for help. Thanks to this app, I have improved my French. I would recommend it to anyone.
Samantha Klich
Android user
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.
Anna
iOS user
Best app on earth! no words because it’s too good
Thomas R
iOS user
Just amazing. Let's me revise 10x better, this app is a quick 10/10. I highly recommend it to anyone. I can watch and search for notes. I can save them in the subject folder. I can revise it any time when I come back. If you haven't tried this app, you're really missing out.
Basil
Android user
This app has made me feel so much more confident in my exam prep, not only through boosting my own self confidence through the features that allow you to connect with others and feel less alone, but also through the way the app itself is centred around making you feel better. It is easy to navigate, fun to use, and helpful to anyone struggling in absolutely any way.
David K
iOS user
The app's just great! All I have to do is enter the topic in the search bar and I get the response real fast. I don't have to watch 10 YouTube videos to understand something, so I'm saving my time. Highly recommended!
Sudenaz Ocak
Android user
In school I was really bad at maths but thanks to the app, I am doing better now. I am so grateful that you made the app.
Greenlight Bonnie
Android user
very reliable app to help and grow your ideas of Maths, English and other related topics in your works. please use this app if your struggling in areas, this app is key for that. wish I'd of done a review before. and it's also free so don't worry about that.
Rohan U
Android user
I know a lot of apps use fake accounts to boost their reviews but this app deserves it all. Originally I was getting 4 in my English exams and this time I got a grade 7. I didn’t even know about this app three days until the exam and it has helped A LOT. Please actually trust me and use it as I’m sure you too will see developments.
Xander S
iOS user
THE QUIZES AND FLASHCARDS ARE SO USEFUL AND I LOVE THE SCHOOLGPT. IT ALSO IS LITREALLY LIKE CHATGPT BUT SMARTER!! HELPED ME WITH MY MASCARA PROBLEMS TOO!! AS WELL AS MY REAL SUBJECTS ! DUHHH 😍😁😲🤑💗✨🎀😮
Elisha
iOS user
This apps acc the goat. I find revision so boring but this app makes it so easy to organize it all and then you can ask the freeeee ai to test yourself so good and you can easily upload your own stuff. highly recommend as someone taking mocks now
Paul T
iOS user