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Introduction Measuring Reaction Rate Factors Influencing Reaction Speed Collision Theory Rate Laws and Order of Reaction Integrated Rate Equations Experimental Methods to Measure Rate Temperature and the Arrhenius Equation Key Concepts
Some reactions take place almost instantly, like fireworks, others, like rusting take place in days. How fast do chemical reactions take place? Understanding the answer is important for many things, such as, designing new drugs, minimising the impact of chemicals we burn, and more.
The rate of a chemical reaction is defined as the change in concentration of a reactant or product during a specific unit of time. It gives an idea of how fast the reactants are being consumed or products being formed.
To know with how much speed a reaction is taking place, scientists measure:
A classic example of the chemical reaction is, dissolving magnesium in hydrochloric acid, which releases hydrogen gas. The reaction rate can be established or measured by the amount of hydrogen gas given off in the period of a second.
There is a chemical reaction that takes place when iodine and propanone are mixed with acid. The reaction produces a color change, and the speed of the reaction can be measured by the amount of time it takes for the color of the solution to change.
Rate = Change in concentration / Time taken
Here is a breakdown of the factors that influence the speed of a reaction.
A reaction occurs between particles, so when a concentration is increased, more particles are present, increasing the likelihood of more collisions, speeding the reaction.
Example: Producing hydrogen gas with magnesium increases when the acid concentration is increased.
When there is an increase in temperature, molecules and particles are moving faster and more collisions occur. As a result, the reaction occurs faster.
Rule of thumb: Usually, a 10 degree increase in temperature increases reaction rate by around 2x.
When there is increased surface area, there are more opportunities for collisions, speeding the reaction.
Example: Powdered calcium reacts faster than when it is in a solid block.
One of the many properties of catalysts is that they increase the rate of a reaction while also not getting consumed in the reaction. They also provide a different pathway for the reaction to occur, and in turn, lower the amount of activation energy.
Example: Enzymes, and platinum in catalytic converters.
Higher pressure decreases the distance between gas molecules, leading to more collisions, therefore speeding the reaction.
Collision theory is an explanation for different reaction rates. The theory states that in order for a reaction to occur, gas molecules must collide, and the energy of the collisions must be strong enough to overcome the activation energy barriers.
Not all collisions result in a reaction. Only collisions with energy that is equal to or above the activation energy barrier will result in a reaction.
Activation energy (Ea): The energy required for the reaction to start.
When the temperature is increased, more particles will have the energy equal to or greater than Ea, causing the reaction to happen more quickly.
Catalysts: decrease the activation energy, leading to more successful collisions and faster reactions.
When describing a reaction, the rate of a reaction is typically dependent on the amount of reactants involved. This is shown in a rate law as:
Rate = k[A]ᵐ[B]ⁿ
Where:
Example: For reaction A + B → C with rate = k[A][B]²
The reaction order is determined by experiment and not just from the balanced equation.
Integrated rate laws allow the prediction of the concentration at a given time. The most common cases are first and second order reactions.
First order reactions:
When a graph of ln[A] is made vs time, it yields a straight line and slope = -k
Second order reactions:
A graph of 1/[A] vs time also yields a straight line and slope = k
Half-life, (t₁/₂): the amount of time for half the reactants to react.
For first order reactions, t₁/₂ = 0.693/k
This is independent of the initial concentration.
It is possible to measure the volume of gas that is produced in a given amount of time.
Example: Reaction of HCl with magnesium (Mg) which produces H₂ gas
Mass seems to decrease because the reaction produces a gas. The change in mass can be recorded using a balance.
During a reaction, a colored reactant is measured, and the color (absorbance) is measured over time.
Example: Reaction of I₂ (iodine) with thiosulfate
Take samples at different time intervals, then titrate to determine concentration.
Example: Reaction of sodium thiosulfate with acid
Arrhenius equation:
where:
The Arrhenius equation will help experimentally determine the activation energy and how fast the reactions will occur at various temperatures.
While understanding reaction rates is critical for examination purposes, it also allows Chemists to:
Understand that the study of reaction rate allows the visualisation, prediction and control over the invisible motion of atoms in a chemical reaction.