Question

Nitrogen gas reacts with oxygen gas to make \(\mathrm{NO}(\mathrm{g})\) while absorbing \(180 \mathrm{~kJ}\). Write a properly balanced thermochemical equation for this process.

Short answer

\(N_2(g) + O_2(g) \rightarrow 2NO(g)\; \Delta H = +180 \text{ kJ}\)

Step-by-step solution

Write the Unbalanced Equation

Start by writing the chemical equation for the reaction between nitrogen gas (\(N_2\)) and oxygen gas (\(O_2\)) to form nitrogen monoxide (\(NO\)). \[ N_2(g) + O_2(g) \rightarrow NO(g) \]

Balance the Equation

Balance the chemical equation. Since \(N_2\) provides two nitrogen atoms and \(NO\) has one nitrogen atom, you will need two \(NO\) molecules. Similarly, balance the oxygen atoms. \[ N_2(g) + O_2(g) \rightarrow 2NO(g) \]

Add the Thermochemical Data

Include the energy absorbed in the reaction as part of the equation. Since the problem states that 180 kJ is absorbed, this indicates an endothermic reaction. \[ N_2(g) + O_2(g) \rightarrow 2NO(g) \quad \Delta H = +180 \text{ kJ} \]

Key concepts

Balancing Chemical Equations

Balancing chemical equations is an essential skill in chemistry, as it ensures that mass and charge are conserved in a chemical reaction. The Law of Conservation of Mass tells us that in any chemical reaction, the number of atoms for each element must be the same on both sides of the equation. This is because atoms are neither created nor destroyed in any chemical process. Let's take a closer look at how to balance a reaction properly.

In the original exercise, we're given the unbalanced chemical equation where nitrogen gas (\(\mathrm{N}_2\)) combines with oxygen gas (\(\mathrm{O}_2\)) to form nitrogen monoxide (\(\mathrm{NO}\)). Initially, the equation is:\[\mathrm{N}_2(g) + \mathrm{O}_2(g) \rightarrow \mathrm{NO}(g)\]

To balance this, note that nitrogen (\(\mathrm{N}_2\)) in the reactants contains two nitrogen atoms, while nitrogen monoxide (\(\mathrm{NO}\)) in the products side has one nitrogen atom. This mismatch means we need two \(\mathrm{NO}\) molecules to provide two nitrogen atoms:\[\mathrm{N}_2(g) + \mathrm{O}_2(g) \rightarrow 2\,\mathrm{NO}(g)\]

Now for the oxygen, we have two atoms in \(\mathrm{O}_2\) and each \(\mathrm{NO}\) has one, needing two \(\mathrm{NO}\) to match. Balancing chemical equations is like solving a puzzle, where the pieces must fit perfectly without any atom left behind!

Endothermic Reactions

Endothermic reactions are chemical reactions that absorb energy from their surroundings. This typically takes the form of heat absorption, making the surroundings feel cooler. In an endothermic reaction, the products have higher enthalpy (heat content) compared to the reactants.

In the context of the exercise, when nitrogen and oxygen gases react to form nitrogen monoxide, the process absorbs energy. The absorbed energy is indicated by the positive \(\Delta H\) in the thermochemical equation. This particular reaction absorbs \(+180\,\text{kJ}\) of heat:\[\mathrm{N}_2(g) + \mathrm{O}_2(g) \rightarrow 2\,\mathrm{NO}(g)\quad \Delta H = +180\,\text{kJ}\]

These reactions are crucial in many areas of science and industry. Understanding endothermic processes helps you know why certain reactions require an initial input of energy to proceed. Common examples include photosynthesis and the evaporation of water.

Reaction Enthalpy

Reaction enthalpy, represented by the symbol \(\Delta H\), denotes the change in heat energy during a chemical reaction at constant pressure. It tells you whether a reaction releases or absorbs energy. A positive \(\Delta H\) indicates that the reaction is endothermic, signifying an absorption of heat, while a negative \(\Delta H\) signifies an exothermic reaction, which releases heat.

In the example given, the reaction of nitrogen and oxygen gases to form nitrogen monoxide absorbs \(180\,\text{kJ}\) of energy, making the \(\Delta H\) value positive:\[\Delta H = +180\,\text{kJ}\]
This shows the amount of energy required for the reaction to occur. Reaction enthalpy is vital in determining the energy efficiency of chemical processes. It helps chemists predict how much energy is required or released, aiding in the design of industrial processes, where energy management is critical.