Thermodynamics

• Chemical energy stored by molecules can be released as heat during chemical reactions when a fuel like methane, cooking gas, or coal burns in air
• Chemical energy may also be used to do mechanical work when a fuel burns in an engine or to provide electrical energy through a galvanic cell like a dry cell
• Various forms of energy are interrelated and can be transformed from one form to another under certain conditions
• The study of energy transformations forms the subject matter of thermodynamics
• The laws of thermodynamics deal with energy changes of macroscopic systems involving a large number of molecules rather than microscopic systems containing a few molecules
• Thermodynamics is based on the initial and final states of a system undergoing a change
• The laws of thermodynamics apply when a system is in equilibrium or moves from one equilibrium state to another
• Macroscopic properties like pressure and temperature do not change with time for a system in an equilibrium state
• In thermodynamics, a system refers to the part of the universe where observations are made, while the surroundings include everything other than the system
• The system and surroundings together constitute the universe
• For practical purposes, the surroundings are that portion of the remaining universe which can interact with the system
• Types of systems:

  • Open System: Exchange of energy and matter between system and surroundings
  • Closed System: No exchange of matter, but exchange of energy is possible
  • Isolated System: No exchange of energy or matter between the system and the surroundings
• The state of a system is described by its measurable or macroscopic properties like pressure, volume, and temperature
• Variables like pressure, volume, and temperature are called state variables or state functions because their values depend only on the state of the system
• The state of a thermodynamic system is described by its measurable or macroscopic properties
• The internal energy of a system, U, represents the total energy of the system and may change when heat passes into or out of the system, work is done on or by the system, or matter enters or leaves the system
• An adiabatic system does not allow heat exchange between the system and surroundings through its boundary
• An adiabatic process is a process in which there is no transfer of heat between the system and surroundings
• In an adiabatic process, the wall separating the system and the surroundings is called the adiabatic wall
• J. P. Joule conducted experiments between 1840-1850 showing that a given amount of work done on a system, regardless of the path, produces the same change of state, measured by the change in temperature
• Internal energy, U, of a system is a state function, where the adiabatic work required to bring about a change of state is equal to the difference between the value of U in one state and that in another state
• Other familiar state functions include V (volume), p (pressure), and T (temperature)
• Heat, q, is transferred from the surroundings to the system or vice versa without the expenditure of work
• In chemical thermodynamics, q is positive when heat is transferred from the surroundings to the system, increasing the internal energy of the system; q is negative when heat is transferred from the system to the surroundings, decreasing the internal energy of the system
• The change in internal energy of a system when no heat is absorbed but work is done on the system is represented by ∆U = wad, where the wall is adiabatic
• The change in internal energy of a system when no work is done but q amount of heat is taken out from the system and given to the surroundings is represented by ∆U = -q, where the wall is thermally conducting
• The change in internal energy of a system when w amount of work is done by the system and q amount of heat is supplied to the system is represented by ∆U = q - w, indicating a closed system
• In single-celled organisms, substances can easily enter the cell due to a short distance, while in multicellular organisms, the distance is larger due to a higher surface area to volume ratio
• Multicellular organisms require specialised exchange surfaces for efficient gas exchange of carbon dioxide and oxygen due to their higher surface area to volume ratio
• When heat is absorbed by the system at constant pressure, changes in enthalpy are measured
• For 1 mol of water vaporized at 1 bar pressure and 100°C:

  • The change H2O (l) → H2O (g) can be calculated using the formula ∆U = ∆H – ∆ngRT
  • Heat capacity is a measure of heat transferred to a system, proportional to the temperature rise
• Heat capacity (C) is directly proportional to the amount of substance and is the quantity of heat needed to raise the temperature of one mole by one degree Celsius
• Extensive properties like mass, volume, internal energy, and heat capacity depend on the quantity of matter, while intensive properties like temperature, density, and pressure do not
• The relationship between heat capacities at constant volume (CV) and constant pressure (CP) for an ideal gas is given by CP – CV = R
• Calorimetry is an experimental technique used to measure energy changes associated with chemical or physical processes
• ∆U measurements are done at constant volume using a bomb calorimeter, while ∆H measurements are done at constant pressure in a calorimeter
• Enthalpy change (∆rH) of a reaction is given by the sum of enthalpies of products minus the sum of enthalpies of reactants
• Enthalpy change in a reaction is calculated using the formula:

  ΔH = (∑Hproducts) - (∑Hreactants)

  Where:

  • ∑ represents summation
  • ai and bi are the stoichiometric coefficients of the products and reactants respectively in the balanced chemical equation
• Standard Enthalpy of Reactions:

  • Enthalpy of a reaction depends on the conditions under which it is carried out
  • Standard enthalpy of reaction is the enthalpy change when all participating substances are in their standard states
• Enthalpy Changes during Phase Transformations:

  • Phase transformations involve energy changes
  • Melting, vaporization, and sublimation are examples of phase transformations that require or release heat