Question

For each of the following processes, indicate whether the signs of \(\Delta S\) and \(\Delta H\) are expected to be positive, negative, or about zero. (a) A solid sublimes. (b) The temperature of a sample of \(\mathrm{Co}(s)\) is lowered from \(60^{\circ} \mathrm{C}\) to \(25^{\circ} \mathrm{C}\) (c) Ethyl alcohol evaporates from a beaker. (d) \(\mathrm{A}\) diatomic molecule dissociates into atoms. (e) A piece of charcoal is combusted to form \(\mathrm{CO}_{2}(g)\) and \(\mathrm{H}_{2} \mathrm{O}(g)\).

Short answer

(a) ΔS: positive, ΔH: positive (b) ΔS: negative, ΔH: negative (c) ΔS: positive, ΔH: positive (d) ΔS: positive, ΔH: positive (e) ΔS: positive, ΔH: negative

Step-by-step solution

(a) Solid sublimation

For a solid to sublime, it has to convert from a solid-state into a gas state. In the solid state, the particles are highly ordered, and in the gas state, the particles are randomly distributed. Therefore, the process results in increased entropy, so ΔS is positive. Since energy is absorbed for the solid particles to break away and become gaseous particles, the process is endothermic; thus, ΔH is positive.

(b) Temperature decrease of Co(s)

The temperature of a sample of solid cobalt is decreased from 60°C to 25°C. When the temperature of a substance decreases, its particles have less kinetic energy, and the substance becomes more ordered, leading to decreased entropy. As a result, ΔS is negative. Since losing temperature implies losing heat energy, the process is exothermic, and ΔH is negative.

(c) Evaporation of ethanol

When ethanol evaporates, it changes from a liquid state to a gaseous state. Going from liquid to gas means increased randomness for the particles, resulting in increased entropy. Thus, ΔS is positive. The process also requires energy to overcome intermolecular forces and turn the liquid into a gas, so it is endothermic, and ΔH is positive.

(d) Diatomic molecule dissociation

In this process, a diatomic molecule dissociates into individual atoms. Since the number of particles in the system has increased, the randomness and disorder have also increased. Consequently, ΔS is positive. Breaking apart molecules requires energy input to overcome the bonds, making the process endothermic. Therefore, ΔH is positive.

(e) Charcoal combustion

During the combustion of charcoal, solid carbon reacts with oxygen to form CO2(g) and H2O(g). Since the process involves the formation of gaseous products from a solid reactant, the overall randomness and disorder increase, resulting in a positive ΔS. Combustion reactions are exothermic as they release heat energy, so ΔH is negative.

Key concepts

Entropy Change (ΔS)

Understanding entropy change (ΔS) is akin to examining the level of disorder or randomness within a system. Fundamentally, entropy is a measure of the number of specific ways in which a thermodynamic system may be arranged, often leading to a state of confusion among students because it's an abstract concept. However, when objects like a block of ice melt or vaporize — think of ethyl alcohol evaporating from your beaker during an experiment in the lab — the molecules spread out and the disorder increases. This is why we have a positive ΔS, signifying increased entropy.

Conversely, when a sample of cobalt cools down, it exhibits a decrease in entropy; the molecular motion slows, leading to a more ordered state. Here, ΔS is negative, indicative of lower entropy. Students find it helpful to link entropy to the degree of predictability in a system. An orderly, tightly packed ice cube has low entropy because it's predictable, whereas a vapor cloud has high entropy due to its unpredictable nature.

Enthalpy Change (ΔH)

When discussing the concept of enthalpy change (ΔH), it's like talking about the heat content of a system at constant pressure. It's a measure usually concerned with heat changes during chemical reactions or phase transitions. For a solid to sublime, or a liquid to evaporate, heat must be absorbed; the system essentially pulls in energy, which leads to an increase in enthalpy — a positive ΔH, as observed when a solid sublimes or ethyl alcohol vaporizes.

It's important to emphasize to students that the reverse, an exothermic reaction such as charcoal combustion, which gives off heat, is marked by a negative ΔH. The concept is akin to your bank account where depositing money is analogous to endothermic reactions (positive ΔH), and paying bills resemble exothermic reactions (negative ΔH). Simple everyday analogies can bridge the gap between the elusive nature of thermodynamic concepts and the concrete reality students are familiar with.

Phase Transitions

Phase transitions are the transformations from one state of matter to another, like ice melting into water or water vaporizing into steam. These transitions are pivotal events where both entropy and enthalpy play a role. Each phase transition comes with its own set of characteristics regarding energy and particle arrangement, which can be challenging for students to visualize.

Take sublimation, for instance, a process where a solid turns directly into a gas without passing through the liquid phase. This fascinating transformation entails a substantial gain in entropy as the ordered solid lattice becomes a disorderly gas, and endothermic enthalpy change as energy is required for the phase change. Providing students with an everyday example, such as dry ice subliming into carbon dioxide gas, helps to concretize these abstract processes.

Chemical Kinetics

When moving from the theoretical to the practical—specifically the speed at which chemical reactions occur—we delve into the field of chemical kinetics. It deals with understanding how different variables such as concentration, temperature and catalysts affect the rate of a reaction. The dissociation of a diatomic molecule into atoms is a compelling example, as the speed at which this process occurs is heavily influenced by the aforementioned factors.

Students are frequently intrigued by how reactions that might seemingly occur instantaneously can actually take thousands of years without a catalyst. The understanding of chemical kinetics is not only crucial for mastering chemistry but also for appreciating the complex balances at play in our everyday lives, from the baking of bread to the production of medicine. Digesting kinetics concepts paves the way for grasping how entropy, enthalpy, and phase transitions can affect the practical outcome of chemical processes.