Lecture 9.1 — Factors Affecting Reaction Rates
What the rate of a reaction means: the speed of conversion of reactants to products; examples include rusting (slow), baking (moderate), explosions (very fast).
Core question: why do some reactions occur faster than others, and how can we make them faster when needed?
Activation energy (E_a): the minimum energy that reacting particles must have to overcome the energy barrier and react; depends on the molecular structure and physical characteristics of the molecules.
Two basic determinants of the rate of reaction:
Frequency of collisions: more particles lead to more collisions; as particle number decreases, collision frequency decreases and so does the rate.
Energy of collisions: collisions must have enough energy to overcome E_a to form products.
Collision theory analogy: finding a partner with the right energy and orientation is like dating; you need both the right energy (activation energy) and the correct orientation to form a bond.
Activation energy concept will be discussed more when temperature and temperature effects on rate are covered; E_a governs whether a collision leads to reaction.
Relationship between collision frequency and energy: both factors determine whether a collision will result in a successful reaction.
How to Change the Rate of a Reaction
Temperature: increasing temperature gives particles more energy and makes them move faster, increasing both collision frequency and the energy of collisions.
Concentration: higher concentration means more particles in the same space, leading to more collisions and a greater chance of proper orientation.
Pressure (gases): increasing pressure brings particles closer together, increasing collision frequency without necessarily changing temperature.
Surface area (solids): increasing surface area exposes more sites for reaction, effectively increasing accessible collisions.
Physical state and accessibility: smaller pieces of a solid have more surface accessible to solvent or other reactants, increasing the rate.
Temperature and Collision Dynamics
Cold vs hot: colder particles move slower with less energy; higher temperature speeds up particles and increases collision frequency and energy.
Warmer conditions lead to more frequent collisions and more energetic collisions, raising the rate.
Beaker demonstrations illustrate that as temperature rises, the reaction proceeds faster due to increased collision rate and energy.
Practical rule: for many reactions, a rise of about 10 degrees Celsius can double the rate (not universal, but common).
You do not need to boil to speed up a reaction: a modest temperature increase is often sufficient.
Concentration Effects
Higher concentration of a dissolved reactant increases the rate because there are more particles to collide with each other.
More collisions increase the chance of a collision with the correct orientation that leads to bond formation.
Whether in solution or gas phase, increasing concentration generally accelerates the rate via more frequent collisions.
Physical States and Dissolution
When dissolving solids (e.g., sugar in water), smaller pieces dissolve faster than a single large chunk.
Reason: smaller pieces have more surface area exposed to the solvent, allowing more collision opportunities and faster access to reactants.
In general, smaller, more mobile particles move faster, increasing collision frequency and energy transfer.
For gases, increasing pressure (while keeping temperature constant) increases collision frequency due to reduced volume.
Gas Pressure Details
Raising pressure in a gas increases collision frequency without changing the temperature (energy per molecule remains tied to temperature).
At low pressure, particles are spread out, leading to longer times between collisions; at high pressure, collisions occur more frequently, speeding up the reaction.
Surface Area and Solid Reactions
Reactions on solids often occur at the surface; increasing the surface area (by breaking into smaller pieces) increases the rate.
The greater the surface area, the greater the chance of collisions at the reactive interface.
What is a Catalyst?
A catalyst speeds up a reaction without being consumed in the overall process.
It provides a different mechanism with a lower activation energy, allowing more molecules to react with the energy they already have.
Analogy: helping an elderly person climb stairs by giving them a hand — the end result is the same, but the process is faster; the helper (catalyst) is not used up and can assist others.
Catalysts lower the activation energy, enabling more molecules to reach the transition state and form products.
Catalysts do not change the stoichiometry or the final amount of product that can be formed; they simply speed up the rate.
The activation energy can be lowered by catalysts, enabling more reactant molecules to surpass the barrier with available energy.
Catalysts in Industry and Everyday Life
In industry, catalysts save money and energy by enabling reactions at lower temperatures or with faster rates.
They also help reduce pollution by allowing reactions to proceed with less energy input and fewer unwanted byproducts.
Everyday catalysts include: nickel in margarine production; iron in the Haber process for ammonia synthesis; platinum in car catalytic converters, which reduce pollutants by converting some emissions (e.g., CO) into less harmful substances (e.g., CO2 and N2).
Catalysis and Biochemistry
Catalysts are essential in biochemistry: enzymes act as biological catalysts, accelerating metabolic reactions in the body.
Most physiological processes rely on enzymes to proceed at essential rates; the human body has numerous enzymes working concurrently to support thought, movement, and metabolism.
Understanding catalysis is critical for medicine, nursing, and allied health fields because many bodily processes are enzyme-catalyzed.
Product Formation and Stoichiometry
The amount of product formed is determined by the stoichiometry of the reaction and the starting amounts of reactants.
Catalysts affect the rate, not the ultimate quantity of product under given starting conditions.
Quick Recap and Formulas
Key idea: Reaction rate depends on collision frequency and collision energy, governed by temperature, concentration, pressure, surface area, and catalysis.
Arrhenius relationship (standard formula):
\small k = A \, e^{-\frac{E_a}{RT}}
Conceptual takeaway: Lower activation energy (via a catalyst) increases the fraction of collisions that lead to product formation, speeding up the reaction.
Real-world relevance: Catalysts enable industrial scales of production, improve energy efficiency, and reduce environmental impact; enzymes enable life-sustaining metabolism.