Question

Consider a gas-phase reaction. (a) Cooling the mixture of reactants can slow and even stop the chemical reaction. Explain why this is so. (b) Increasing the number of reactant molecules in the flask makes the reaction go faster. Explain why this is so.

Short answer

(a) Cooling the mixture of reactants reduces their kinetic energy, leading to lower energy collisions between reactant particles. As a result, fewer collisions possess the necessary activation energy to initiate a chemical reaction, slowing down the reaction rate or stopping it completely. (b) Increasing the number of reactant molecules in the flask enhances the frequency of favorable collisions involving the reactants, according to the collision theory. This leads to a faster reaction rate due to more reactant particles being present in proximity, increasing the chance of successful collisions and product formation.

Step-by-step solution

a) Explanation of the effect of cooling on the reaction

: To explain the effect of cooling on the reaction, we need to consider the activation energy and the distribution of molecular energies in the gas phase. Activation energy is the minimum amount of energy required for a reaction to occur between specific reactants. In a gas phase reaction, reactant particles collide with each other, and if the collision has enough energy (above the activation energy), the reaction takes place. Otherwise, the molecules simply collide and bounce off one another without any reaction occurring. When a mixture of reactant molecules is cooled, reduced kinetic energy results in lower energy collisions between reactant particles. Consequently, fewer collisions possess the necessary energy (activation energy) to initiate a chemical reaction. As the temperature decreases, the number of effective collisions also decreases, ultimately leading to a slower reaction rate or even a complete halt in the reaction process.

b) Explanation of the effect of increasing reactants in the flask

: To explain the effect of increasing the number of reactant molecules in the flask, we need to consider the collision theory of chemical reactions. Collision theory states that the number of successful molecular collisions directly affects the rate of a chemical reaction. More specifically, the more reactant molecules present in a system, the higher the likelihood of successful collisions, and thus the faster the reaction. By increasing the number of reactant molecules in the flask, we are indirectly enhancing the frequency of favorable collisions involving the reactants, leading to a faster reaction rate. This happens because more reactant particles are present in proximity, increasing the chance of successful collisions and the formation of product molecules.

Key concepts

Activation Energy

Activation energy is a fundamental concept in chemical kinetics that represents the minimum energy required for reactants to undergo a chemical transformation.
Think of activation energy as a barrier; only those particles with energy equal to or greater than this barrier can proceed to form products.
In the context of a gas-phase reaction, reactants must collide with sufficient energy to overcome this barrier and transition into products.

• High activation energy means fewer particles have enough energy to react, leading to a slower reaction rate.
• Lower activation energy allows more particles to surpass the barrier, resulting in a faster reaction.

When the reaction mixture is cooled, the kinetic energy of particles decreases, leading to fewer successful collisions.
This means fewer particles can overcome the activation energy barrier, and the reaction rate slows down or stops.

Collision Theory

Collision theory is integral to understanding why some reactions occur faster than others. It suggests that molecules must collide to react, but not all collisions result in a reaction.
For a successful collision, molecules need to have:

• Sufficient energy: The collision's energy must surpass a minimum threshold, known as the activation energy.
• Correct orientation: Molecules must be oriented properly to break and form the necessary bonds.

When we apply collision theory to a gas-phase reaction, increasing the temperature or concentration of reactants can increase the number of successful collisions, thus speeding up the reaction.
In essence, collision theory helps us understand how changes in physical conditions can influence reaction rates.

Reaction Rate

The reaction rate measures how quickly a chemical reaction proceeds. It is influenced by factors such as concentration, temperature, and the presence of catalysts.
A key factor in the reaction rate of gas-phase reactions is how often reactant molecules collide with the proper energy and orientation.
Factors affecting reaction rate include:

• Concentration: More molecules mean more collisions and a faster reaction rate.
• Temperature: Higher temperatures increase kinetic energy, most collision energy, and reaction speed.
• Catalysts: Lower activation energy, enhancing reaction rates without being consumed.

By understanding the determinants of reaction rate, we can manipulate conditions to accelerate or slow reactions, which is particularly useful in industrial applications.

Gas-Phase Reaction

Gas-phase reactions occur entirely within the gaseous state and involve the collision of reactant gas molecules. Unlike reactions in the liquid or solid phase, gas-phase reactions are notably influenced by temperature and concentration.
In these reactions, kinetic energy plays a critical role. As gases are composed of particles moving rapidly and freely, changes in temperature can significantly impact their speed and energy.
Key aspects of gas-phase reactions include:

• High mobility: Particles move quickly, increasing the likelihood of collisions.
• Dependence on pressure and volume: As gases expand or are confined, the frequency of collisions can change.

Gas-phase reactions are crucial in various fields, including atmospheric chemistry and industrial processes. Understanding their dynamics allows for better control and optimization of reaction conditions.