Question

Hydrogenation reactions (for example, the process of converting \(\mathrm{C}=\mathrm{C}\) bonds to \(\mathrm{C}-\mathrm{C}\) bonds in food industry) are facilitated by the use of a transition metal catalyst, such as Ni or \(\mathrm{Pt}\). The initial step is the adsorption, or binding, of hydrogen gas onto the metal surface. Predict the signs of \(\Delta H, \Delta S,\) and \(\Delta G\) when hydrogen gas is adsorbed onto the surface of Ni metal.

Short answer

In the adsorption of hydrogen gas onto the surface of Nickel metal, the signs of \(\Delta H\), \(\Delta S\), and \(\Delta G\) are all predicted to be negative.

Step-by-step solution

Understanding Enthalpy (ΔH)

Enthalpy (\( \Delta H \)) is the measure of heat content in a system. In the context of this exercise, when hydrogen gas molecules bind to the metal surface, they likely release energy, which implies that the system will lose heat. This would make \( \Delta H \) negative because heat is released from the system.

Understanding Entropy (ΔS)

Entropy (\( \Delta S \)) measures the degree of disorder or randomness in a system. In the current scenario, gas molecules are being adsorbed onto the surface of a solid (i.e., from a more disordered state to a less disordered state). This implies a decrease in randomness or disorder, and so \( \Delta S \) will be negative.

Understanding Gibbs Free Energy (ΔG)

Gibbs free energy (\( \Delta G \)) is given by the equation \( \Delta G = \Delta H - T\Delta S \). Here, \( \Delta H \) and \( \Delta S \) are both negative. Considering that temperature (T) is always positive (measured in Kelvin), the negative sign for \( \Delta S \) becomes positive when multiplied by the positive T. This makes the second term (T\Delta S) positive. If we assume the released heat (\( \Delta H \)) to be larger than the T\Delta S term, then \(\Delta G \) overall will be negative.

Key concepts

Catalyst in Chemical Reactions

Catalysts play a pivotal role in enhancing the rate of chemical reactions without undergoing any permanent chemical change themselves. Think of them as facilitators that allow a process to occur more efficiently. In the context of hydrogenation reactions, catalysts like nickel (Ni) or platinum (Pt) provide a surface for hydrogen molecules to adsorb and interact with other reactants. This lowers the energy barrier for the reaction to proceed, much like smoothing out a road for vehicles to pass more easily.

For students learning about chemistry, it's vital to understand the role of a catalyst in reactions. It doesn't participate directly or change the final products of the reaction; it merely speeds up its rate. By providing a more effective pathway for the reaction to occur, a catalyst can lead to significant increases in reaction efficiency, which is crucial in industrial applications like the food industry, where the conversion of double bonds into single bonds in unsaturated fats is a common process.

Enthalpy (ΔH) in Reactions

Enthalpy, denoted as \( \Delta H \), is a term in chemistry that relates to the total heat content of a system. It's a reflection of the energy absorbed or released during a reaction. To make this idea more digestible, consider it as the reaction's heat signature. In an exothermic reaction such as the hydrogenation we are discussing, where \( \mathrm{C}=\mathrm{C} \) bonds are converted to \( \mathrm{C}-\mathrm{C} \) bonds, energy is released when the hydrogen gas binds to the catalyst's surface. This would result in a negative \( \Delta H \), indicating that the reaction releases heat to its surroundings.

Understanding \( \Delta H \) is crucial for students as it helps predict reaction behavior. When \( \Delta H \) is negative, it implies that the reaction may occur spontaneously since the system is losing heat energy. This concept is a cornerstone in thermochemistry and pivotal to energy considerations in reaction design.

Entropy (ΔS) in Chemistry

Entropy, symbolized as \( \Delta S \), is a measure of the randomness or disorder within a system. In the classroom, entropy can be likened to the messiness of a teenager's room – the more disorganized it is, the higher the entropy. For hydrogenation reactions, the adsorption of hydrogen gas onto a metal surface represents a decrease in entropy because the movement of the gas molecules becomes restricted as they bind to the solid surface. Hence, \( \Delta S \) is negative in this scenario.

In terms of educational application, grasping the concept of entropy is essential for understanding the direction in which a reaction might proceed. A negative \( \Delta S \) typically suggests a system becoming more ordered, which can be counterintuitive to students who might expect chemical reactions to move towards chaos. Emphasizing this point helps demystify why certain reactions require specific conditions to proceed.

Gibbs Free Energy (ΔG)

Gibbs free energy, expressed as \( \Delta G \), is a thermodynamic quantity that predicts the direction of a chemical reaction and its spontaneity. The equation that relates enthalpy, entropy, and Gibbs free energy is \( \Delta G = \Delta H - T\Delta S \). As explained earlier, both \( \Delta H \) and \( \Delta S \) can be negative in the hydrogenation process, leading to interesting implications for this equation.

For learning purposes, it's important to note that reactions will spontaneously proceed when \( \Delta G \) is negative. In our exercise, if the enthalpic contribution (\( \Delta H \)) overshadows the product of temperature and entropy change (\( T\Delta S \)), the overall Gibbs free energy will be negative, indicating spontaneity. By understanding this relationship, students can predict the feasibility of chemical reactions under various conditions and comprehend the interplay between energy, entropy, and temperature in affecting chemical processes.