collision theory
Flashcard 1
Front:
Learning Objective – Reaction Rates and Collision Theory
Back:
To measure the rate of a chemical reaction using appropriate methods and outline the conditions that need to be met for a reaction to occur.
Success Criteria:
• Define reaction rate as “amount of substance produced/used over time”
• Outline the three conditions for collision theory
Flashcard 2
Front:
Activate Prior Knowledge – Limestone Chips Experiment
Back:
You have two types of limestone chips: granules vs crushed. You add the limestone to 10 mL of 2 mol/L hydrochloric acid each.
Which reaction will be faster? Why?
How can we measure how fast each reaction goes?
Crushed limestone (5 g) vs Limestone granules (5 g).
Flashcard 3
Front:
Reaction Rates – Fast vs Slow Reactions
Back:
Chemical reactions are taking place all around us.
Some are very fast:
• E.g. In a car accident, the airbag reaction must happen extremely quickly to produce the gas that inflates the bag.
Some are very slow:
• E.g. Rusting of exposed metal on a scratched car happens over a long period.
Flashcard 4
Front:
Definition of Reaction Rate
Back:
The rate at how fast or slow a reaction occurs can be defined as:
• Rate of formation of products (how quickly products are formed) OR
• Rate of consumption of reactants (how quickly reactants are consumed).
Formula:
Rate of reaction=amount of substance used or producedtime taken\text{Rate of reaction} = \frac{\text{amount of substance used or produced}}{\text{time taken}}Rate of reaction=time takenamount of substance used or produced
Units may include mol/s, g/s, or mL/s depending on the reaction and data collection method.
Flashcard 5
Front:
Graphing Reaction Rates
Back:
Instead of calculating the rate, we can graph the amount of substance used/produced over time.
• The slope or gradient gives the reaction rate.
• A steeper gradient = greater reaction rate.
• A horizontal gradient indicates the reaction has stopped.
Flashcard 6
Front:
Instantaneous Rate of Reaction – Example Calculations
Back:
a)
Rate = amount produced / time taken = 16 / 30 = 0.53 mL/s
b)
Volume of H₂(g) collected = 32 – 24 = 8 mL
Rate = amount produced / time taken = 8 / 30 = 0.27 mL/s
Flashcard 7
Front:
Collision Theory – Three Conditions
Back:
Collision theory describes chemical reactions in terms of collisions between reacting particles.
For a reaction to occur:
Particles must collide.
The collision energy must be equal to or greater than the activation energy (Ea).
The particles must collide with the proper orientation.
Vocabulary:
Activation energy (Ea) = minimum collision energy required for a reaction to occur.
Flashcard 8
Front:
Potential Energy Profiles – Transition State
Back:
When particles approach each other:
• Repulsive forces between electron clouds cause them to slow down and lose kinetic energy, which becomes potential energy.
• If sufficient kinetic energy is available, they form a transition state (unstable, bonds breaking/forming).
• The minimum energy to reach this state = activation energy (Ea).
• Stronger bonds = higher Ea.
Vocabulary:
Transition state (activation complex): highest potential energy state for the reaction.
Flashcard 9
Front:
Potential Energy Profiles – Endothermic vs Exothermic
Back:
Endothermic:
• More bond breaking than forming.
• Net energy absorption.
• Products have higher energy than reactants.
Exothermic:
• More bond forming than breaking.
• Net energy release.
• Products have lower energy than reactants.
Remember:
Bond breaking = energy absorbed.
Bond forming = energy released.
Flashcard 10
Front:
Lesson Closure – Reaction Rate and Collision Theory
Back:
• Reaction rate = “amount of substance produced/used over time.”
• Three conditions for collision theory:
Collide
Energy ≥ Ea
Proper orientation
⚡ FLASHCARDS — Factors Affecting Reaction Rate (Part 1)
Flashcard 11
Front:
Learning Objective – Factors Affecting Reaction Rate
Back:
To explain how nature of reactants, surface area, temperature, concentration, catalysts, and pressure affect reaction rate, using collision theory.
Success Criteria:
• Describe and predict how nature, surface area, and concentration influence reaction rate.
• Apply collision theory to explain how these factors affect collision frequency and energy.
Flashcard 12
Front:
Six Factors Affecting Reaction Rate
Back:
Nature of reactants
Surface area of solid reactants
Concentration of reactants in a solution
Temperature
Gas pressure/volume
Use of catalyst
These change the number of collisions per unit time → affect reaction rate.
Flashcard 13
Front:
Factor 1 – Nature of Reactants
Back:
Refers to the inherent chemical and physical properties of substances involved.
These influence how easily reactants interact and form products.
Physical States:
• Gases and liquids react faster than solids (greater freedom of movement → more collisions).
Flashcard 14
Front:
Bonding Type and Strength; Polarity and Solubility
Back:
• Strong covalent bonds → slower reactions (more energy needed).
• Ionic compounds in solution → faster (ions are separated and free to interact).
• Polar reactants dissolve better in polar solvents → faster reactions.
• Non-polar reactants in polar solvents → slower (poor solubility).
Flashcard 15
Front:
Factor 2 – Surface Area (State of Subdivision)
Back:
• Reactions occur at the surface boundary between phases.
• Increasing surface area exposes more particles to collision.
• More collisions = higher chance of successful collisions = faster rate.
Flashcard 16
Front:
Methods to Increase Surface Area
Back:
• Crushing/grinding → smaller particles.
• Dissolving → particles spread out.
• Thin films/sheets → more exposed area.
• Aerosols → droplets with large surface area.
• Stirring/agitation → maximizes interaction (though not surface area).
Flashcard 17
Front:
Factor 3 – Concentration of Reactants in Solution
Back:
Increasing concentration → faster reaction (more frequent collisions).
Decreasing concentration → slower reaction.
If concentration increases → collisions increase → higher chance of successful collisions → faster reaction.
Flashcard 18
Front:
Independent Practice – Examples
Back:
1⃣ Sodium chloride (ionic) reacts faster in water than methane (covalent).
2⃣ Powdered CaCO₃ reacts faster with HCl than marble chips (more surface area).
3⃣ 2.0 M HCl reacts faster with Mg ribbon than 0.5 M (higher concentration → more collisions).
🔥 FLASHCARDS — Factors Affecting Reaction Rate (Part 2)
Flashcard 19
Front:
Factor 4 – Temperature
Back:
Temperature measures average kinetic energy (E = ½mv²).
• Higher temperature = higher kinetic energy = faster moving particles.
• More frequent and energetic collisions → higher reaction rate.
Flashcard 20
Front:
Temperature – Collision Energy
Back:
• Average collision energy increases.
• More particles have energy ≥ activation energy (Ea).
• Higher proportion of successful collisions → faster rate.
Flashcard 21
Front:
Factor 5 – Gas Pressure
Back:
• Increasing pressure (by reducing volume or adding gas) increases gas molecule concentration.
• More collisions per unit time → faster rate.
• Doubling pressure often doubles reaction rate.
Example:
Unopened Coke can (higher pressure) is fizzier than an opened one (lower pressure).
Flashcard 22
Front:
Factor 6 – Catalyst
Back:
A catalyst increases reaction rate but is not consumed.
• Provides an alternate pathway with lower activation energy.
• More collisions have enough energy to react → faster rate.
• Does NOT reduce Ea itself; adding more catalyst beyond saturation doesn’t infinitely speed reactions.
Flashcard 23
Front:
Types of Catalysts
Back:
• Homogeneous: same phase as reactants (e.g. H₂SO₄ in esterification).
• Heterogeneous: different phase (e.g. Pt in catalytic converters; Fe in Haber process).
• Auto-catalyst: product catalyzes its own reaction (e.g. MnO₄⁻ + oxalic acid).
Transition metals (V, Mn, Pt, Pd, Au, Rh) often act as catalysts.
Flashcard 24
Front:
Enzymes – Biological Catalysts
Back:
• Enzymes are highly specific – only correct substrate fits in active site (“lock and key” model).
• Weak intermolecular forces hold substrate; enzyme changes shape to strain bonds.
• Products leave, enzyme returns to original form.
Enzymes = high specificity. Inorganic catalysts = not specific.
Flashcard 25
Front:
Nano-catalysts
Back:
Operate at nanoscale (1–100 nm).
• Very large surface area-to-volume ratio → more active sites.
• Surface atoms more reactive due to unsatisfied bonds and quantum effects.
• Can be engineered for specific shapes/sizes/reactions.
Examples:
Pt nanoparticles (fuel cells), Fe₂O₃ (pollutant breakdown), TiO₂ (water splitting, air purification).
Flashcard 26
Front:
Independent Practice – Factors 4–6
Back:
1⃣ Higher temperature → faster rate (more particles have energy ≥ Ea).
2⃣ Higher gas pressure → faster rate (more frequent collisions).
3⃣ Catalyst → increases rate by lowering activation energy pathway, not used up.
Flashcard 27
Front:
Lesson Closure – Factors Affecting Reaction Rate
Back:
• Changes in temperature, pressure, and catalysts affect rate of reaction.
• Collision theory explains rate changes in terms of collision frequency and energy