Question

Would you expect a reaction in which all reactants are in the gas phase to go faster, slower, or remain at the same rate if the temperature is increased while everything else is held constant? Explain your answer.

Short answer

In conclusion, we would expect a reaction with all reactants in the gas phase to go faster when the temperature is increased, while everything else is held constant. This is due to the increased kinetic energy of the molecules, leading to more effective collisions that can overcome the activation energy barrier, as explained by the Arrhenius equation and collision theory.

Step-by-step solution

Collision Theory

Collision theory states that for a reaction to occur, reactant molecules must collide with proper orientation and with sufficient energy to overcome the energy barrier (activation energy).

The Arrhenius Equation

The Arrhenius equation relates the rate constant (k) of a reaction to its activation energy (Ea) and temperature (T): \[k = Ae^{\frac{-Ea}{RT}}\] where: - k is the rate constant - A is the frequency factor or pre-exponential factor - Ea is the activation energy - R is the gas constant - T is the temperature in Kelvin

Effect of Temperature on Reaction Rate

According to the Arrhenius equation, when the temperature increases, the exponent \(\frac{-Ea}{RT}\) becomes less negative, resulting in an increase in the rate constant (k). This in turn leads to an increased reaction rate. This can be explained through collision theory as well. When the temperature increases, the molecules have more kinetic energy, which leads to a higher chance of collisions with the required energy to overcome the activation energy barrier. So more molecules can collide with proper orientation and sufficient energy, which increases the reaction rate.

Conclusion

If the temperature is increased while everything else is held constant, we would expect a reaction with all reactants in the gas phase to proceed at a faster rate. This is due to the increased movement and kinetic energy of the molecules, leading to more effective collisions that are energetic enough to overcome the activation energy barrier.

Key concepts

Collision Theory

Collision theory provides vital insights into how chemical reactions occur. The theory suggests that for a reaction to take place, reactant molecules need to collide. However, not just any collision will do. These molecules must hit each other with the right orientation and with enough energy to surpass a specific energy barrier, known as the activation energy. This means not every collision leads to a reaction.

• Proper orientation: Molecules need to align themselves precisely.
• Sufficient energy: They must have enough kinetic energy to drive the reaction forward.

Hence, the number and nature of collisions directly impact the reaction rate. In simple terms, more frequent and effective collisions usually result in a faster reaction.

Arrhenius Equation

The Arrhenius equation is a fundamental formula used to understand and calculate the rate of a chemical reaction. It expresses how the rate constant, denoted as \( k \), relates to temperature \( T \), activation energy \( Ea \), and a constant \( A \), known as the frequency factor. The equation is given by:
\[ k = Ae^{\frac{-Ea}{RT}} \]

• \( k \): Rate constant, signifies how fast the reaction proceeds.
• \( Ea \): Activation energy, the energy barrier that must be overcome.
• \( A \): Frequency factor, indicating the likelihood of a collision resulting in a reaction.
• \( R \): Universal gas constant.

This relationship shows us that as the temperature increases, the exponential part becomes a smaller negative number due to the division by \( T \). This results in a larger \( k \), meaning a faster reaction rate.

Activation Energy

Activation energy \( Ea \) is a crucial concept in understanding reaction rates. It is the minimum energy threshold that reactants must overcome to transform into products. It acts like a hill that molecules need to climb to start reacting.

• Higher \( Ea \): Fewer molecules will have enough energy to react, slowing down the reaction.
• Lower \( Ea \): More molecules can participate in the reaction, leading to a faster rate.

In practical terms, reactions with low activation energy are generally faster because the energy barrier is easy to overcome, while those with high activation energy often require additional conditions, like increased temperature, to proceed at a noticeable rate.

Temperature Effect on Reactions

Temperature is a key player in changing reaction rates. A common observation in chemistry is that an increase in temperature usually speeds up a reaction. Here's why: when the temperature rises, molecules move faster and possess more kinetic energy.

• Increased kinetic energy: Molecules collide more frequently and with greater force.
• More effective collisions: High-energy collisions are more likely to surpass the activation energy barrier.
• Faster reactions: With more energetic collisions, the reaction rate generally increases.

The increase in collision frequency and energy aligns perfectly with both collision theory and the Arrhenius equation. As these theories suggest, a rise in temperature means more molecules have the energy required to react, hence speeding up the overall reaction process.