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Thermodynamics

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Thermodynamics

  1. 1. THERMODYNAMICS Mrs. Praveen Garg VITS College, Satna
  2. 2. INTRODUCTION • Thermodynamics is the branch of physics that deals with the relationships between heat and temperature and their relation to energy and work. • Thermo – Heat , Dynamics – Motion or move • In particular, it describes how thermal energy is converted to and from other forms of energy and how it affects matter.
  3. 3. Thermodynamics Heat, work, different form of energy, heat changes Heat flow from high temperature to low temperature Example: Gas Pressure, Enthalpy Temperature, Entropy Volume, Gibbs energy Thermodynamics
  4. 4. Branches of thermodynamics  Classical thermodynamics  Statistical mechanic  Chemical thermodynamics  Equilibrium thermodynamics
  5. 5. Classical Thermodynamics: • In classical thermodynamics, the behavior of whole matter is analyzed, such as temperature and pressure. It helps to calculate properties and to predict the characteristics of the matter that is undergoing the process. Statistical Thermodynamics: • In statistical thermodynamics, study the different properties of each and every molecule and to characterize the behavior of a each molecules.
  6. 6. Macroscopic analysis Microscopic analysis
  7. 7. Equilibrium Thermodynamics: • Equilibrium thermodynamics is the study of transformations of energy and matter as they approach the state of equilibrium. Chemical Thermodynamics: • Chemical thermodynamics is the study of how work and heat relate to each other both in chemical reactions and in changes of states.
  8. 8. Basic Concepts of Thermodynamics – Thermodynamic Terms  Systems: A system is a specific portion of matter or substances, which have to be study.  Surrounding: Except system everything in universe is called surrounding.  System boundary: It separate system and surrounding. The system boundary may be real or imaginary, fixed or deformable. System + Surrounding = Universe
  9. 9. Boundary Diathermic Adiabatic Across heat change easily (Conducting) heat can not flow (Insulating) (∆q=0) Heat flow High temp. Low temp.
  10. 10. There are three types of system as: • Open System – In an open system, the mass and energy both may be transferred between the system and surroundings. • Closed System – Across the boundary of the closed system, the transfer of energy takes place but the transfer of mass doesn’t take place. Energy can exchange in the form of heat and work. Refrigerator, compression of gas in the piston-cylinder assembly are examples of closed systems. • Isolated System – An isolated system cannot exchange both energy and mass with its surroundings. The universe is considered an isolated system.
  11. 11. Real Imaginary Mass change – form of vapour Due to open boundary Conducting-- Heat change Mass can not change due to closed boundary Real, adiabatic (Insulated) ∆q=0
  12. 12. SYSTEM (variable/parameter/properties) Extensive  Its depend on amount of matter, size of system or sample.  Example: • Mass (water) • Volume (water) • No. of moles (solution) • Heat capacity, • Energy- Heat, Work, Enthalpy, Entropy, Internal Energy, Gibbs free energy etc. Intensive  Does not depend on size of sample/amount of matter.  Example: • Density (Glass of water) • Specific heat capacity • Temperature (Whole or small part of room) • Pressure, Viscosity (constant) • Concentration (sugar solution) • Molarity, Normality, Mole fraction, Molar entropy, molar enthalpy, pH etc.
  13. 13. Extensive (X) Extensive (Y) = Entensive (Z) Example: Molarity = Molar entropy= entropy/mole No. of moles Volume
  14. 14. State variable/ State function/State parameter • Condition or state of system is defined as variables. • Example: Volume Gas Pressure Variables Temperature State • Variables of state: pressure(P), temperature(T), volume(V), no. of moles, enthalpy(H), entropy(S), internal energy (E), Gibbs free energy (G) etc. • Function of state can be change due to change in variable. • These changes in variables depend on initial and final condition of state and does not depend on path.
  15. 15. State function: ∆T = Tf – Ti ∆P = Pf – Pi ∆V = Vf – Vi • Changes in temperature, pressure, volume depend on initial (i) and final (f) condition. • But Heat and Work depend only in path function, it does not depend on initial and final condition. • In the case of any cyclic process, change in state variables is always zero (no change). ∆T = 0 ∆P = 0 ∆V = 0 But: ∆q ≠ 0 ∆w ≠ 0
  16. 16. Types of thermodynamic process  Isothermal process: Occur in constant/same temperature. ∆T = 0 If ideal gas, ∆T = 0, ∆E = 0 Internal energy (E = K.E. ∝ T)  Isobaric process: Occur in constant pressure.  Isochoric process: Occur in constant volume (closed vessel).  Adiabatic process: In this process, heat can not change. ∆q = 0 Note: These all process can be reversible or irreversible.
  17. 17. Reversible • Which occur in infinite • Infinite time to finish • Small steps • Slow process • Each step in equilibrium exist always. Irreversible • At once • Finite time to finish • Fast process • Equilibrium depend on initial and final
  18. 18. 1) Zeroth law of thermodynamics: When two systems are each in thermal equilibrium with a third system, the first two systems are in thermal equilibrium with each other. 2) First law of thermodynamics: It is the law of conservation of energy. The change in a system’s internal energy is equal to the difference between heat added to the system from its surroundings and work done by the system on its surroundings. • LAWS OF THERMODYNAMICS
  19. 19. 3) Second law of thermodynamics: • Heat does not flow spontaneously from a colder region to a hotter region, or, equivalently, heat at a given temperature cannot be converted entirely into work. • Consequently, the entropy of a closed system, or heat energy per unit temperature, increases over time toward some maximum value. • Thus, all closed systems tend toward an equilibrium state in which entropy is at a maximum and no energy is available to do useful work.
  20. 20. 4) Third law of thermodynamics: • The entropy of a perfect crystal of an element in its most stable form tends to zero as the temperature approaches absolute zero. • This allows an absolute scale for entropy to be established that, from a statistical point of view, determines the degree of randomness or disorder in a system.
  21. 21. • State in which no spontaneous change (macroscopic changes) are observed with respect of time. • It may be three types: • Thermal equilibrium- equal temperature • Mechenical equilibrium- equal force/pressure • Chemical equilibrium- equal chemical force EQUILIBRIUM
  22. 22. “Systems that are in thermal equilibrium exist at the same temperature”. • When two or more bodies at different temperatures are brought into contact then after some time they attain a common temperature and they are said to exist in thermal equilibrium. Thermal equilibrium
  23. 23. • When a body ‘A’ is in thermal equilibrium with another body ‘b’, and also separately in thermal equilibrium with a body ‘C’, then body ‘B’ and ‘C’ will also be in thermal equilibrium with each other. This statement defines the zeroth law of thermodynamics. The law is based on temperature measurement. • Zeroth law of thermodynamics states that, if the systems are in thermal equilibrium, no heat flow will take place. • This law is mostly used to compare temperatures of different objects. ZEROTH LAW OF THERMODYNAMICS
  24. 24. • The most common application of the zeroth law of thermodynamics can be seen in thermometers. • Another example of the zeroth law of thermodynamics is when you have two glasses of water. One glass will have hot water and the other will contain cold water. Now if we leave them in the table for a few hours they will attain thermal equilibrium with the temperature of the room. Application of Zeroth law
  25. 25. First Law of Thermodynamics • The First Law of Thermodynamics states that heat is a form of energy, and thermodynamic processes are of conservation of energy. • This means that heat energy cannot be created or destroyed. It can be transferred from one location to another and converted to other forms of energy. • The First Law says that the internal energy of a system has to be equal to the work that is being done on the system, plus or minus the heat that flows in or out of the system. "So, it’s a restatement of conservation of energy." • This is expressed mathematically as: ΔU = Q – W where ΔU is the change in the internal energy, Q is the heat added to the system, and W is the work done by the system.
  26. 26. First law of thermodynamics: ΔEuniv = ΔEsys + ΔEsurr = 0 The changes of in the internal energy of a system is equal to the sum of the heat gained or lost by the system and the work done by or on the system. ΔEsys = q - w (closed system) Example: Beaker of water on hot plate When hot plate is on, temp and internal energy of the system increase, ΔE is positive. When hot plate is off, temp and internal energy decrease, ΔE is negative.
  27. 27. Application of First law • Light bulbs transform electrical energy into light energy (radiant energy). • One pool ball hits another, transferring kinetic energy and making the second ball move. • Plants convert the energy of sunlight (radiant energy) into chemical energy stored in organic molecules.
  28. 28. Limitations of First Law of Thermodynamics • The limitation of the first law of thermodynamics is that it does not say anything about the direction of flow of heat. • It does not say anything whether the process is a spontaneous process or not. • The reverse process is not possible. In actual practice, the heat doesn’t convert completely into work.
  29. 29. HEAT • Heat is energy transferred between substances or systems due to a temperature difference between them. • Heat is conserved as a form of energy. • Heat is the transfer of thermal energy between systems, while work is the transfer of mechanical energy between two systems. • Heat is transfered via solid material (conduction), liquids and gases (convection), and electromagnetical waves (radiation). • Heat describes the transfer of thermal energy between molecules within a system and is measured in Joules. Heat measures how energy moves or flows. • Temperature describes the average kinetic energy of molecules within a material or system and is measured in Celsius (°C), Kelvin(K), Fahrenheit (°F).
  30. 30. INTERNAL ENERGY • Internal energy, in thermodynamics, the property or state function that defines the energy of a substance. • The value of the energy depends upon the state of the substance and not upon the nature of the processes by which it attained that state. • According to first law of thermodynamics, system undergoes a change of state only due to work is involved, the work is equal to the change in internal energy. • The law also implies that if both heat and work are involved in the change of state of a system, then the change in internal energy is equal to the heat supplied to the system minus the work done by the system. • The internal energy as a sum of terms that can be interpreted as kinetic energy, potential energy, and chemical energy.
  31. 31. ENTHALPY • Enthalpy, the sum of the internal energy and the product of the pressure and volume of a thermodynamic system. • Enthalpy is an energy-like property or state function—it has the dimensions of energy (and is thus measured in units of joules or ergs). • Its value is determined entirely by the temperature, pressure, and composition of the system. • The enthalpy (H) equals the sum of the internal energy (E) and the product of the pressure (P) and volume (V) of the system: H = E + PV • The change in internal energy is equal to the heat transferred to, less the work done by the system. If the only work done is a change of volume at constant pressure, the enthalpy change is exactly equal to the heat transferred to the system.
  32. 32. ENTROPY • Entropy, the measure of a system’s thermal energy per unit temperature. • Entropy is a function of the state of the system, so the change in entropy of a system is determined by its initial and final states. • The amount of entropy is also a measure of the molecular disorder, or randomness, of a system. • The entropy of a substance can be obtained by measuring the heat required to raise the temperature a given amount, using a reversible process. • Enthalpy is the measure of total heat present in the thermodynamic system where the pressure is constant while Entropy is the measure of disorder in a thermodynamic system. where Q is the heat content and T is the temperature. ΔS = ΔQ /T
  33. 33. • Zero entropy means perfect knowledge of a state ; no motion, no temperature, no uncertainty. • Entropy is not conserved because it can be created. But once it is created and dissipated in surroundings the action cannot be reversed, so therefore entropy always keep on increasing. • As we know laws of conservation of mass and energy say that both the quantities are uncreatable and undestroyable.
  34. 34. • This is irreversible process. • The second law of thermodynamics states that the heat energy cannot transfer from a body at a lower temperature to a body at a higher temperature without the addition of energy. • The second law of thermodynamics states that the entropy of an isolated system is always increasing • The entropy of the system is measured in terms of the changes the system has undergone from the previous state to the final state. • Thus the entropy is always measured as the change in entropy of the system denoted by ∆S. • The universe is an isolated system. This means the entropy of the universe will always increase. The entropy change of the universe is the sum of the entropy change of the system and the entropy change of the surroundings. Second Law of Thermodynamics
  35. 35. Causes of increase in entropy of the closed system are: • In a closed system, the mass of the system remains constant but it can exchange the heat with surroundings. • Any change in the heat content of the system leads to disturbance in the system, which tends to increase the entropy of the system. • Due to internal changes in the movements of the molecules of the system, there is disturbance inside the system. • This causes irreversibilities inside the system and an increase in its entropy.
  36. 36. • Exergonic Reaction: It is spontaneous reaction that release energy. The entire reaction is catabolic. The release of energy (Gibbs free energy) is negative. • Endergonic Reaction: It is anabolic reaction that consume energy. It has a positive ΔG because energy is required to break bonds. • The free energy gained or lost in a reaction can be calculated: ΔG = ΔH – TΔS Where, G= Gibbs free energy H= enthalpy S= entropy T= temperature
  37. 37. Application of Second law • This law is applicable to all types of heat engine cycles including Otto, Diesel, etc. for all types of working fluids used in the engines. This law has led to the progress of present day vehicles. • Another application is refrigerators and heat pumps based on the Reversed Carnot Cycle. If you want to move heat from a body at a lower temperature to a body at a higher temperature, then you have to supply external work. • Removing heat from the food items in the refrigerator and throwing it away to the higher temperature atmosphere doesn’t happen automatically. We need to supply external work via the compressor to make this happen in the refrigerator. • Air conditioner is another example.
  38. 38. • The third law of thermodynamics states that the entropy of a perfect crystal at a temperature of zero Kelvin (absolute zero) is equal to zero. • At zero kelvin the system must be in a state with the minimum energy. • If the crystal has only one minimum energy state, third law holds true. • Entropy is related to the number of microstates, and with only one microstates available at zero kelvin, the entropy exactly zero. Third Law of Thermodynamics
  39. 39. • It helps in the calculation of the absolute entropy of a substance at any temperature ‘T’. • These determinations are based on the heat capacity measurements of the substance. Application of the third law
  40. 40. THANK YOU

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