Welcome to Rates of Reaction!
Have you ever wondered why milk goes sour in a few days on the kitchen counter, but lasts much longer in the fridge? Or why a firework explodes in a fraction of a second, while an iron gate takes years to rust? It all comes down to the rate of reaction (how fast or slow a chemical reaction happens).
In this chapter of C2 Chemistry, you will learn how to measure reaction rates, how to interpret graphs, and how particles behave using collision theory. Don't worry if this seems tricky at first — once you master a few core ideas, you will find these questions very predictable and rewarding in your exams!
1. What is the Rate of Reaction?
The rate of reaction is a measure of how quickly reactants are turned into products in a chemical reaction.
To find the rate, we measure either:
• How quickly a reactant is used up, OR
• How quickly a product is formed over time.
The Rate Equations
\(\text{Mean Rate of Reaction} = \frac{\text{Quantity of reactant used}}{\text{Time taken}}\)
\(\text{Mean Rate of Reaction} = \frac{\text{Quantity of product formed}}{\text{Time taken}}\)
Units of Rate
The units depend on what you measure and how you measure it:
• If measuring mass in grams (\(g\)) and time in seconds (\(s\)): \(g/s\) (or \(g\text{ }s^{-1}\))
• If measuring gas volume in cubic centimetres (\(cm^3\)) and time in seconds (\(s\)): \(cm^3/s\) (or \(cm^3\text{ }s^{-1}\))
Quick Summary: Rate tells us how much product is made (or reactant lost) per second.
2. Measuring Rates in the Laboratory
In your C2 Chemistry course, there are three main experimental methods you need to know to measure reaction rates:
Method 1: Measuring the Volume of Gas Produced
This method is used when a reaction produces a gas (for example, magnesium reacting with hydrochloric acid to produce hydrogen gas: \(Mg + 2HCl \rightarrow MgCl_2 + H_2\)).
• Equipment: Conical flask, delivery tube, and a gas syringe (or an upturned measuring cylinder filled with water).
• How it works: As the reaction takes place, the gas collects in the syringe. You record the volume of gas at regular time intervals (e.g., every \(10\text{ s}\)).
• Advantage: Very accurate and allows you to plot a complete graph of volume against time.
Method 2: Measuring the Loss in Mass
This method works well when a heavy gas like carbon dioxide (\(CO_2\)) is produced (for example, marble chips reacting with acid: \(CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2\)).
• Equipment: Conical flask, cotton wool plug, and an accurate digital balance.
• How it works: As the reaction proceeds, the gas escapes through the cotton wool, causing the total mass on the balance to decrease. You record the mass at regular time intervals.
• Why the cotton wool? It allows the gas to escape freely while preventing any acid spray or liquid splashes from escaping, which would cause an inaccurate reading.
Method 3: Measuring the Formation of a Precipitate (The "Disappearing Cross")
This method is used when a reaction produces an insoluble solid (a precipitate) that turns the solution cloudy (turbid).
• Reaction: Sodium thiosulfate and hydrochloric acid produce solid sulfur: \(Na_2S_2O_3 + 2HCl \rightarrow 2NaCl + SO_2 + S + H_2O\).
• How it works: Place the conical flask on top of a piece of paper with a black cross (\(X\)) drawn on it. Mix the chemicals and start the stopwatch. Look down through the solution from above. Stop the timer the exact moment the cross is no longer visible.
• Calculating Rate: \(\text{Rate} = \frac{1}{\text{time}}\) (units: \(s^{-1}\)).
• Limitation: Deciding when the cross completely disappears is subjective and can vary from person to person.
Key Takeaway: Choose your method based on what is made: gas syringe for any gas, mass balance for heavy gases (\(CO_2\)), and the disappearing cross for reactions that form a precipitate.
3. Interpreting Rate Graphs
A typical graph of Volume of Gas Produced vs Time or Mass of Product vs Time always has a characteristic curve with three clear stages:
1. At the start (Steepest gradient): The reaction is at its fastest. There is a high concentration of reactant particles, meaning frequent collisions occur.
2. In the middle (Curve becomes less steep): The reaction slows down. Reactant particles are being used up, so there are fewer particles available to collide.
3. At the end (Horizontal / Flat line): The reaction has stopped (rate = \(0\)). One or both of the reactants have been completely used up (this reactant is called the limiting reactant).
Finding the Rate from a Graph
• Mean (Average) Rate: Divide the total volume of gas by the total time taken.
• Rate at a Specific Time: Draw a straight line touching the curve at that exact point (a tangent), then calculate the gradient of your tangent line:
\(\text{Gradient} = \frac{\text{Change in } y}{\text{Change in } x} = \text{Rate at that point}\)
Common Mistake to Avoid: Don't assume a flat line means the reaction is fast! A horizontal line on a product-time graph means the volume is no longer changing — the reaction has completely finished.
4. Collision Theory: How Reactions Happen
For a chemical reaction to occur, particles must collide with each other. However, simply bumping into each other is not enough!
According to Collision Theory, a reaction will only happen if:
1. The particles collide with each other.
2. The particles collide with sufficient energy (equal to or greater than the activation energy).
What is Activation Energy?
Activation Energy (\(E_a\)) is the minimum amount of energy that colliding particles must possess for a collision to be successful and lead to a reaction.
Analogy: Think of activation energy like pushing a boulder over a hill. If you don't push it hard enough to reach the peak, it just rolls back down (no reaction). Once you give it enough energy to get over the top, the reaction happens!
Key Takeaway: To speed up a reaction, you must either make collisions happen more frequently, or ensure that a greater fraction of collisions has energy \(\ge E_a\) (or both!).
5. Factors Affecting the Rate of Reaction
There are four major factors you can change to alter the rate of a reaction:
1. Surface Area of Solid Reactants
• What happens: Breaking a large solid lump into smaller pieces (or a fine powder) exposes many more reactant particles to the surrounding solution or gas.
• Collision Theory explanation: A larger surface area means more exposed particles, which leads to more frequent collisions per second.
• Result: The rate of reaction increases.
• Did you know? Custard powder or flour in a mill can create an explosive fireball if sparked, because the fine dust has an enormous surface area exposed to oxygen!
2. Concentration of Solutions
• What happens: Increasing the concentration means putting more reactant particles into the exact same volume of liquid.
• Collision Theory explanation: Because the particles are packed closer together, they bump into each other more often, resulting in more frequent collisions per unit time.
• Result: The rate of reaction increases.
• Note for gases: Increasing the pressure of a gas does the same thing — it pushes gas molecules closer together, increasing collision frequency.
3. Temperature
Increasing temperature has a powerful effect because it changes two things at once:
• Point A (Speed): Particles gain thermal energy and convert it to kinetic energy, so they move faster. This causes more frequent collisions.
• Point B (Energy - Most Important!): A much larger proportion of the colliding particles now have energy equal to or greater than the activation energy (\(\ge E_a\)). Therefore, a higher percentage of collisions are successful.
• Result: The rate of reaction increases dramatically. (In many reactions, a \(10\text{ }^\circ\text{C}\) rise roughly doubles the rate!).
4. Use of a Catalyst
• Definition: A catalyst is a substance that speeds up the rate of a chemical reaction without being used up or chemically changed at the end of the reaction.
• How it works: A catalyst provides an alternative reaction pathway that has a lower activation energy (\(E_a\)).
• Collision Theory explanation: Because the energy barrier is lower, a greater proportion of particle collisions have enough energy to react successfully.
• Example: Manganese(IV) oxide (\(MnO_2\)) catalyzes the decomposition of hydrogen peroxide into water and oxygen: \(2H_2O_2 \rightarrow 2H_2O + O_2\).
• Biological Catalysts: Enzymes are biological catalysts made of protein that speed up metabolic reactions inside living organisms.
Key Takeaway:
• Surface Area \(\rightarrow\) More exposed particles \(\rightarrow\) More frequent collisions.
• Concentration / Pressure \(\rightarrow\) Particles closer together \(\rightarrow\) More frequent collisions.
• Temperature \(\rightarrow\) Faster particles AND more collisions have energy \(\ge E_a\).
• Catalyst \(\rightarrow\) Lowers activation energy by offering an alternative pathway.
6. Summary & Exam Success Tips
Golden Exam Phrase:
Whenever you explain an increase in rate, always use the word "frequent" or the phrase "per unit time" / "per second".
Say: "There are more frequent collisions" or "more collisions per second."
Avoid: "There are more collisions" (this alone loses marks because it doesn't mention time!).
Quick Review Box:
• Rate = \(\frac{\text{Amount}}{\text{Time}}\)
• Gradient of tangent = Rate at that specific instant.
• Activation energy (\(E_a\)) = Minimum energy required for a successful collision.
• Catalysts lower \(E_a\) and remain unchanged at the end.