1. Define internal energy, work, and heat. internal energy: Kinetic energy + potential energy Heat: energy that moves into or out of the system because of a temperature difference between the system and its surroundings. Work: energy exchange that results when a force F moves an object through a distance d; work (w) = F x d.

2. State the first law of thermodynamics and deduce the mathematical expression for the law of sign conventions. Or relate the ∆U and q. Statement: total energy of an isolated system remains constant though it may change from one form to another. - 1 st law related to law of conservation of energy Heat absorb by the system, (q) = +ve Heat evolve by the system, (q) = -ve Work done by the system (w) = -ve Work done on the system (w) = +ve A gas is enclosed in a vessel with a piston. Heat flows into the vessel from the surroundings, which are at a higher temperature. As the temperature of the gas increases, the gas expands, lifting the weight (doing work). The system in gains internal energy from the heat absorbed and loses internal energy via the work done. ∆U = q + w

3. Explain why the work done by the system as a result of expansion or contraction during a chemical reaction is P∆V. Or As hydrogen is evolved, work must be done by the system to push back the atmosphere. How can you calculate this work? The work done by the system in expanding of gas equals the force of gravity times the distance the piston moves. w = F. h

4. Enthalpy (H) or Heat of reaction(q) - The amount of heat released or absorbed during a reaction H = U + PV - ∆H = negative; exothermic reaction (system evolve energy and surroundings absorb energy) - ∆H = positive; endothermic reaction (system absorb energy and surroundings evolve energy) - - at once it was thought that spontaneous reaction must be exothermic where “melting of ice is an endothermic but spontaneous”. As a result 1 st law of thermodynamics cannot not predict whether a reaction is favorable

5. Show how heat of reaction at constant pressure (q p ) equal the change of enthalpy (∆H) Applied heat = Increase of internal energy + Work done by the system to increase the volume of gas Calculation:

6. State the 2 nd law of thermodynamics in terms of system plus surroundings. - The second law is expressed in terms of entropy. Entropy: - measure of randomness of a system - measure of how dispersed the energy of a system is among the different possible ways - Does not predict whether a reaction is favorable Statement: the total entropy of a system and its surroundings always increases for a spontaneous process. Entropy increases due to increase of 1.Number of molecules or ions 2.Kinetic energy of the molecules 3.temperature

7. Why is the entropy of liquid water higher than that of solid water (ice)? In an ice crystal- water molecules occupy regular fixed positions. Molecules can only vibrate or oscillate about these fixed positions In the liquid state- water molecules can rotate as well as vibrate internally and can move around somewhat. —the energy is dispersed over more available energy states in the liquid. Thus, the entropy of liquid water is expected to be higher than that of ice

8. “For a spontaneous process the change in entropy of the system is greater than the heat flow” - 2 nd law in terms of system Entropy is produced during a spontaneous process. Suppose a spontaneous process occurs within a system at a given temperature T—say, a chemical reaction in a flask. As the chemical reaction occurs, entropy is produced. At the same time, heat might flow into or out of the system as a result of the thermal contact. Heat flow is also a flow of entropy, because it is a dispersal of energy, either into the flask or outside of it. In general, the entropy change associated with a flow of heat q at an absolute temperature T can be shown to equal q/T. The net change of entropy : The quantity of entropy created during a spontaneous process is a positive quantity

9. Entropy Change for a Phase Transition Certain processes occur at equilibrium or, very close to equilibrium. Ice (0 C) Water (l) (0 C) If heat is slowly absorbed by the system, it remains very near equilibrium, but the ice melts. At equilibrium conditions, no significant amount of entropy is created. The entropy change only from the absorption of heat. Therefore, ∆S = entropy created + q/T (heat flow) ∆S = 0 + q/T = q/T

10. The entropy change S for a phase transition equals H/T, where H is the enthalpy change. Why is it that the entropy change for a system in which a chemical reaction occurs spontaneously does not equal H/T?