Final answer:
In chemistry, ΔH < 0 indicates an exothermic reaction, and equilibrium shifts are explained by Le Chatelier's Principle based on temperature and pressure changes. Phase diagrams can be used to determine the state of water under specified conditions.
Step-by-step explanation:
Understanding Chemical Reactions and Phase Changes
When analyzing the impact of certain conditions on a chemical reaction or a system at equilibrium, two important concepts in chemistry come into play: enthalpy change (ΔH) and Le Chatelier's Principle. For a reaction where ΔH is less than zero (ΔH < 0), it suggests that the reaction is exothermic, releasing heat to the surroundings. In such cases, an increase in temperature would generally shift the equilibrium towards the reactants, as the system seeks to absorb excess heat, according to Le Chatelier's Principle.
Le Chatelier's Principle also explains how a system at equilibrium responds to changes in concentration, pressure, and temperature. For instance, an increase in pressure typically shifts the equilibrium towards the side with fewer moles of gas, while a decrease in pressure favors the side with more moles of gas. When a reaction component like HI is removed, the equilibrium will shift towards producing more HI to counteract this change.
Using a phase diagram is a practical way to determine the state of water under specific conditions of pressure and temperature. For example, at a pressure of 50 kPa and a temperature of -10 °C, water would be in the solid state (ice), while at a pressure of 50 kPa and a temperature of 50 °C, it would exist as a liquid. In a gaseous state, we might find water at 25 kPa and 200 °C, indicating the dependence of the physical state on both temperature and pressure.