Question

One of the components of polluted air is NO. It is formed in the high- temperature environment of internal combustion engines by the following reaction: $$ \mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}(g) \quad \Delta H=180 \mathrm{kJ} $$ Why are high temperatures needed to convert \(\mathrm{N}_{2}\) and \(\mathrm{O}_{2}\) to NO?

Short answer

High temperatures are needed to convert N2 and O2 into NO since the reaction is endothermic, requiring energy input to progress. The strong triple bonds in N2 and double bonds in O2 have high bond dissociation energies, necessitating a significant amount of energy to break them. By providing a high-temperature environment, the required energy is supplied, allowing the bonds in the reactants to be broken and the formation of NO molecules. The high temperature also ensures that reactant molecules have enough kinetic energy to overcome the activation energy required for the reaction to proceed.

Step-by-step solution

Understanding the Reaction

First, let us analyze the given reaction: \[ \text{N}_2(g) + \text{O}_2(g) \rightarrow 2 \text{NO}(g) \quad \Delta H = 180 \text{ kJ} \] This reaction represents the formation of nitrogen monoxide (NO) from nitrogen (N2) and oxygen (O2) molecules. Notice that the reaction is endothermic, as indicated by the positive enthalpy change (\(\Delta H = 180 \text{ kJ}\)). This means that the reaction requires energy to proceed.

Bond Strengths and Dissociation Energies

Now let's focus on the bond strengths and dissociation energies involved in this reaction. Both N2 and O2 molecules have very strong bonds, with triple bonds in N2 (\(\text{N} \equiv \text{N}\)) and double bonds in O2 (\(\text{O} = \text{O}\)). These strong bonds result in high bond dissociation energies, meaning that a significant amount of energy is required to break these bonds. The energy needed to break one mole of N2 triple bonds is approximately 945 kJ/mol, while the energy needed to break one mole of O2 double bonds is around 498 kJ/mol. On the other hand, the bond dissociation energy for NO is around 631 kJ/mol. This means that the energy required to break the N2 and O2 bonds is greater than the energy released when forming NO bonds.

Overcoming the Energy Barrier

Since the reaction is endothermic and high bond dissociation energies are involved, a large amount of energy must be provided to overcome the energy barrier for the reaction to proceed. By providing a high-temperature environment, the necessary energy is supplied to break the N2 and O2 bonds and allow the formation of NO molecules. The high temperature provides the kinetic energy needed by the reactant molecules to collide with sufficient energy to overcome the activation energy required for the reaction to proceed.

Conclusion

In conclusion, high temperatures are required to convert N2 and O2 into NO because the reaction is endothermic, and strong bonds in the N2 and O2 molecules need to be broken for the reaction to occur. The high-temperature environment supplies the necessary energy to overcome the energy barrier, allowing the strong N2 and O2 bonds to be broken and the formation of NO molecules to take place.

Key concepts

Bond Dissociation Energy

Bond dissociation energy is a key concept in understanding chemical reactions, especially endothermic ones like the formation of NO from \(\text{N}_2\) and \(\text{O}_2\).
When we talk about bond dissociation energy, we mean the energy required to break a bond between two atoms in a molecule.
In the reaction forming nitrogen monoxide (NO), we deal with breaking very strong bonds.

• The triple bond in \(\text{N}_2\) is exceptionally strong and requires about 945 kJ/mol to break.
• The double bond in \(\text{O}_2\) needs approximately 498 kJ/mol to dissociate.

These energies tell us how much heat energy is needed to split these molecules.
The greater the bond dissociation energy, the more heat is required, making it challenging for the reaction to occur spontaneously at low temperatures.
In our reaction of forming NO, we also release energy when new bonds are formed, but because \(\Delta H = 180 \text{ kJ}\), we know that the energy needed to break the bonds is greater than the energy released by the new NO bonds.
This contributes to the overall energy demand, signifying why breaking these initial bonds is critical.

Activation Energy

In chemistry, activation energy is the minimum quantity of energy that the reacting species must possess in order to undergo a specified chemical reaction.
This energy barrier must be overcome for a reaction to proceed. Without sufficient energy, even reactants with strong tendencies to form new products cannot react quickly or at all. For the reaction \(\text{N}_2(g) + \text{O}_2(g) \rightarrow 2 \text{NO}(g)\), there is a significant activation energy barrier due to the strong covalent bonds of the \(\text{N}_2\) and \(\text{O}_2\).

• The molecules must collide with enough force to overcome these high initial bond energies.
• This means providing energy usually in the form of heat or through catalysis to initiate the reaction.

The positive \(\Delta H\) of 180 kJ indicates that the products have higher energy states than the reactants, requiring a significant energy input.
Thus, overcoming the activation energy is an essential step, necessitating external energy in this endothermic process.

High Temperature Reaction

High temperature reactions play a vital role in overcoming the intrinsic energy barriers of endothermic reactions.
For \(\text{N}_2\) and \(\text{O}_2\) to form \(\text{NO}\), high temperatures are essential. This condition provides the necessary energy input to overcome bond dissociation and activation energy hurdles.

• With heat, reactant molecules gain kinetic energy, increasing collision frequency and force between molecules.
• These energetic collisions enable bonds to break and new bonds to form, thus permitting the reaction to proceed.

In engines, the combustion of fuel generates such high temperatures, making endothermic processes like forming \(\text{NO}\) feasible.
As a result, reactions that are not spontaneously favorable at room temperature become viable, explaining why certain pollutants form in high-temperature conditions of an engine.
This underscores the relationship between temperature, energy requirements, and reaction feasibility in endothermic reactions.