First and secold law of thermodynamics

• Energy is conserved during chemical reactions.
• The first law states that energy cannot be created or destroyed, only transferred from one form to another.
• Specific heat capacity refers to the amount of heat needed to increase the temperature of a substance by a specific amount per unit mass.
• Heat capacity is defined as the amount of heat required to raise the temperature of an object by a certain number of degrees Celsius.
• Thermodynamics is the study of relationships involving heat, mechanical work, and other aspects of energy and energy transfer.
• Thermodynamic processes are processes that involve changes in the state of thermodynamic systems.
• A thermodynamic system can interact with its surroundings or environment in at least two ways, one of which is heat transfer.
• The first law of thermodynamics is an extension of the principle of conservation of energy.
• The principle includes energy exchange by heat transfer and by the performance of mechanical work and introduces the concept of the internal energy of the system.
• In thermodynamics, the quantity of heat, Q, added to the system and the work, W, done by the system are used to describe the energy relations in any thermodynamic process.
• Both Q and W may be positive, negative or zero.
• The relation between volume and temperature changes for an infinitesimal adiabatic process in an ideal gas is dU = - dW.
• For an ideal monoatomic gas, γ = 1.67.
• The work done during an adiabatic process is W = - nCV(T1 – T2).
• The molar heat capacities of gases at low pressure are shown in Table 6-1.
• The gas in the constant-pressure and the constant-volume process is brought at the same temperature change so dU is the same for both processes, then nCpdT = nCVdT + nRdT.
• For an ideal diatomic gas, γ = 1.41.
• The ratio of heat capacities is defined as γ = Cp/CV.
• For an ideal gas the internal energy change in any process is given by ΔU = nCVdT, whether the volume is constant or not.
• In thermodynamics, heat flows into the system when Q = (+), and out of the system when Q = (–).
• Work is done by the system against its surroundings when W = (+), and on the system by the surroundings when W = (–).
• The work done by the system during a volume change is represented by the area under the curve of pressure versus volume between the limits V1 and V2.
• If the pressure remains constant while the volume changes from V1 to V2, the work done by the system is W = p(V2 – V1).
• In any system in which the volume is constant, the system does no work on the surroundings.
• When both heat transfer and work occur, the total change in internal energy is U2 – U1 = ΔU = Q – W.
• For a cyclic process, the final state is the same as the initial state, and so the total internal energy change must be zero; then U2 = U1 and Q = W.
• For any process taking place in an isolated system, W = Q = 0.
• The work done by the system during a transition between two states depends on the path chosen.
• Equation (6-5) is the First Law of Thermodynamics, a generalization of the principle of conservation of energy to include energy transfer through heat as well as mechanical work.
• In an isolated system, there is no work done on its surroundings and no heat flow to or from its surroundings.
• When a thermodynamic system changes from initial state to a final state, it passes through a series of intermediate states known as a path.
• When a quantity of heat Q is added to a system which does no work, its internal energy increases by an amount equal to Q; that is, ΔU = Q.
• A process that eventually returns a system to its initial state is called a cyclic process.
• The change in internal energy is independent of path, unlike Q and W which depend on the initial and final states and the path leading from one state to the other.
• During a change of state of the system, the internal energy may change from an initial value U1 to a final value U2.
• Internal energy U of a system is defined as the sum of the total kinetic energy of all of its constituent particles and all the total potential energy of interaction among these particles.
• In an isolated system, the internal energy is constant.
• When the system does work W by expanding against its surroundings and no heat is added during the process, energy leaves the system and the internal energy decreases; then ΔU = -W.
• In a constant-volume process, if an infinitesimal quantity of heat dQ flows into the gas, its temperature increases by dT, therefore dQ = nCVdT.
• The first law in differential form is given as dU = dQ – dW.