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Activated Complex: Definition, Theory & Examples

The activated complex is a fundamental concept in chemistry that plays a crucial role in understanding chemical reactions and reaction rates.

During the transition state of a reaction, the activated complex forms as reactants undergo a temporary configuration change before converting into products. This stage accounts for the energy barrier that must be overcome for a reaction to occur.

Catalysts facilitate this process by providing an alternative pathway with lower activation energy. By studying the properties and behavior of the activated complex, scientists can gain insights into reaction mechanisms and design more efficient catalysts.

Definition of the Activated Complex in Chemistry

The activated complex, also known as the transition state, is a crucial concept in chemistry. It refers to an intermediate state that occurs during a chemical reaction. This short-lived arrangement of atoms has high energy and instability.

Intermediate State with High Energy

The activated complex is like a temporary pitstop for molecules during a reaction. Imagine it as a roller coaster ride where the molecules reach the highest point before zooming down again.

In this case, the activated complex represents the highest energy point along the reaction pathway.

Determining Reaction Progression

The activated complex plays a vital role in determining whether a chemical reaction will proceed or not.

If the reactant molecules successfully overcome this energy barrier and form an activated complex, they have a higher chance of proceeding with the reaction and forming new products.

However, if they fail to reach this critical state, they may bounce off each other without any significant change.

Energy Barrier

Think of the activated complex as a hurdle that reactant molecules need to overcome to move from one side of the reaction equation to another.

The height of this energy barrier determines how fast or slow a reaction will occur. Reactions with lower energy barriers are typically faster because they require less energy for reactants to transform into an activated complex.

Activation Energy and its Significance

Activation energy plays a crucial role in chemical reactions. It is the minimum amount of energy required for a reaction to occur. This energy acts as a barrier that must be overcome by reactant molecules before they can form products.

Higher activation energy means slower reactions,

While lower activation energy means faster reactions.

The magnitude of activation energy directly influences the rate at which molecules collide and react. When the activation energy is high, it takes more time and effort for molecules to reach the necessary energy state to initiate a reaction.

As a result, reactions proceed at a slower pace. Conversely, when the activation energy is low, molecules can easily attain the required energy level, leading to faster reactions.

The impact on reaction rates and kinetics

Activation energy affects not only how fast or slow a reaction occurs but also its overall rate and kinetics.

Reactions with higher activation energies have fewer successful collisions between reactant particles because only a small fraction of them possess enough kinetic energy to surpass the activation barrier.

On the other hand, lower activation energies increase the number of successful collisions and enhance reaction rates.

Enzymes: Catalysts that lower activation energies

Enzymes are biological catalysts that play an essential role in speeding up chemical reactions within living organisms. They achieve this by reducing the activation energies required for these reactions to take place.

By providing an alternative pathway with lower activation barriers, enzymes enable more reactant molecules to reach their transition states and form products efficiently.

Theory behind the Activated Complex

The theory behind the activated complex is based on collision theory and transition state theory. These theories provide valuable insights into how chemical reactions occur at a molecular level.

Collision Theory

Collision theory explains that for a chemical reaction to occur, particles must collide with sufficient energy.

This energy, known as activation energy, allows molecules to overcome the repulsive forces between them and undergo a successful reaction. Essentially, it’s like two cars colliding with enough force to cause damage.

Transition State Theory

Transition state theory complements collision theory by describing what happens during a reaction. It states that reactants must pass through an unstable, high-energy state called the transition state or activated complex before forming products.

Think of it as walking on a tightrope; you need to balance perfectly in that brief moment before reaching your destination.

Both theories contribute to our understanding of the formation and behavior of activated complexes. They help us visualize how reactant molecules approach each other, collide, and rearrange their atoms to form new products.

To better grasp these concepts, let’s consider an example involving benzene (a common organic compound). By analyzing its molecular structure using diagrams and graphs, we can deduce how different factors influence its reactivity.

Activation Energy and the Activated Complex

To determine activation energy, scientists use various experimental methods such as temperature changes or catalysts.

By measuring reaction rates at different temperatures, they can calculate activation energies using mathematical equations like the Arrhenius equation or Eyring equation.

Experimental Methods for Determining Activation Energy

  1. Temperature Changes: Scientists observe how a reaction rate changes with temperature. They measure the rate of reaction at different temperatures and plot this data on a graph. From this graph, they can determine the activation energy by analyzing the slope of the line.

  2. Catalysts: Catalysts are substances that speed up chemical reactions without being consumed in the process. By introducing a catalyst into a reaction, scientists can lower the activation energy required for the reaction to occur.

    This allows them to study how changes in activation energy affect the overall reaction rate.

Mathematical Equations for Calculating Activation Energy

  1. Arrhenius Equation: The Arrhenius equation relates the rate constant of a reaction to its activation energy and temperature. It is given by:

k=Ae^{{\frac  {-E_{a}}{RT}}}

In this equation, k is the rate constant, A is the pre-exponential factor (related to collision frequency), Ea is the activation energy, R is the ideal gas constant, and T is the temperature in Kelvin.

  1. Eyring Equation: The Eyring equation provides a more detailed understanding of how reactant molecules transform into an activated complex. It takes into account factors such as enthalpy (ΔH) and entropy (ΔS) changes during a chemical reaction.

    κ = K(kT/h)exp(–ΔG ‡/kT)

Importance of Activation Energy

Understanding activation energy helps scientists design more efficient reactions by optimizing conditions such as temperature and catalyst usage.

For example, enzymes are biological catalysts that play crucial roles in biochemical reactions within living organisms. By studying their activation energies, scientists can develop drugs or therapies that target specific enzymes to modulate their activity.

Determining activation energy and studying the activated complex provide valuable insights into the mechanisms of chemical reactions. This knowledge contributes to advancements in fields such as medicine, materials science, and environmental research.

Calculating Activation Energy

To further understand the concept of activation energy and activated complex in chemistry, let’s dive into a worked example that will help us grasp the calculations involved.

Chemical Reaction and Exothermic Reactions

In this example, we will consider a chemical reaction that is exothermic, meaning it releases energy. The reaction proceeds through multiple steps, with each step having its own activation energy.

Example Scenario

Let’s imagine a scenario where we have an exothermic reaction occurring in two steps. The first step has an activation energy of 50 kJ/mol, while the second step has an activation energy of 30 kJ/mol.

Minimum Amount of Energy Required

To determine the overall activation energy required for this reaction to proceed, we need to find the minimum amount of energy needed among all the steps.

In this case, since we are dealing with an exothermic reaction, the minimum amount would be the lowest activation energy value among all the steps.

Calculating Activation Energy

In our example scenario, the lowest activation energy value is 30 kJ/mol from the second step. Therefore, the overall activation energy for this exothermic reaction would be 30 kJ/mol.

By calculating and understanding these values for each step in a chemical reaction, chemists can gain insights into how reactions occur and identify potential barriers or catalysts that may affect their progress.


We learned that the activated complex is a high-energy intermediate state that molecules must pass through during a chemical reaction. It is formed when reactant molecules collide with enough energy to overcome the activation energy barrier.

Understanding the concept of activation energy was crucial in comprehending the significance of the activated complex.

Activation energy determines whether a reaction will occur or not, as it represents the minimum amount of energy required for reactants to transform into products.

By studying and calculating activation energy, scientists can gain valuable insights into reaction rates and mechanisms.


What are some real-life examples of activated complexes?

The concept of an activated complex can be observed in various real-life scenarios. For example, when wood burns, it undergoes combustion through a series of chemical reactions involving an activated complex. Similarly, during photosynthesis, plants convert carbon dioxide and water into glucose with the help of sunlight and an activated complex.

How does temperature affect activation energy and the formation of an activated complex?

Increasing temperature generally leads to higher kinetic energy among molecules, resulting in more frequent collisions and higher chances for successful collisions that form an activated complex. As a result, higher temperatures decrease activation energy by providing more particles with sufficient energy to overcome it.

Can catalysts influence the formation of an activated complex?

Yes, catalysts can significantly impact the formation of an activated complex by providing alternative reaction pathways with lower activation energies. They achieve this by creating temporary bonds with reactant molecules or altering their orientation during collisions.

Is there any relationship between reaction rate and the concentration of reactants?

Yes, there is a direct relationship between the reaction rate and the concentration of reactants. Generally, an increase in reactant concentration leads to a higher likelihood of successful collisions, resulting in more frequent formation of activated complexes and faster reaction rates.

How can the concept of the activated complex be applied in industrial processes?

Understanding the concept of the activated complex is crucial for optimizing industrial processes. By studying activation energies and reaction mechanisms, scientists and engineers can develop efficient catalysts, control reaction rates, and design more sustainable chemical reactions that minimize energy consumption and waste production.