Chapter 11: Thermodynamics

Essential Definitions

System

• System is the part of universe under study
• Everything else is called surroundings
• Open systems transfer energy and matter to surroundings
• Closed system can transfer energy but not matter
• Isolated system has no energy or matter transfer

State Functions

• State functions describes an aspect of a chemical system
  • Pressure, volume, temperature, and moles are state functions in the ideal gas law
• Thermodynamics state functions can be internal energy (E), enthalpy (H), entropy (S), and Gibbs free energy (G)
• Changes in state function depend on the initial and final state of a system
• Not state functions are heat (q) and work (w) because they rely on the steps used to transform matter from initial to final

Standard State

• Extensive properties of matter change as the amount of sample changes
  • Such as ΔH, ΔS, ΔG, and ΔE
• To make them intensive properties, the temperature, pressute, mass, and physical state must be defined
• Standard state is 1 atm, 25 celsius, and 1 mol of compound present
  • This gives them a superscript of zero to note that they are at standard state
  • ΔH°, ΔS°, ΔG°, ΔE°

Exo- and Endo- Prefixes and Sign Conventions

• Exo-, or exothermic, means energy is being lost from the system to the surroundings
  • Negative sign for quantities
  • ΔH is negative if exothermic
• Endo-, or endothermic, means energy is being gained from surroundings
  • absorbs heat energy
  • ΔH is positive

Types of Energy

• Energy can be heat, light, chemical, nuclear, electrical, or mechanical
• Law of conservation of energy states energy cannot be created or destroyed. All forms of energy can be converted to heat energy
• Can be called kinetic, KE, or potential, PE, energy
  • Kinetic is the energy matter has because of its motion
  • potential energy is stored and released under certain conditions
  • gravitational energy is PE
  • electrostatic attraction between oppositely charge ions is PE

Measurement of Energy

Specific Heat

• Calorie is the amount of heat needed to raise the temperature of 1 gram of something by 1 degree celsius. 1 calorie = 4.184 joules
  • The specific heat of water is the energy needed to raise 1 gram of water by 1 degree celsius = 4.184 J/g*C
• Specific heat equation
  • q = sp.ht x g x ΔT
  • heat energy = (specific heat)(mass in grams)(temperature change)

First Law of Thermodynamics

• First law of thermodynamics states energy is always conserved
• ΔE = q + w
  • change in energy = heat + work
  • if q is positive, heat is added to the system
  • if w is positive, work is done on the system

Work

• Work is the force applied to an object as it moves a certain distance
  • work = force x distance moved
• Can also be found over a given area
  • work = pressure x area x volume change

Definition of qp, qv, ΔE, and ΔH

• First law of thermodynamics
  • ΔE = qp - PΔV
  • heat, q, have the subscipt, p, to indicate the pressure must be held constant
• When heat energy is measure in a calorimeter that does not allow the volume to change, PΔV is zero, resulting in ΔE = qv, where the volume is held constant
  • Calorimeters that do not let volume change are called bomb calorimeters
• Enthalpy change is ΔH where H is the heat content of the substance and the Δ is the difference
  • ΔH = Hproducts - Hreactants
• ΔE = ΔH - PΔV

Standard Enthalpy Changes and the Standard Heat of Reaction

• ΔH is an extensive property and to make it intensive, the amount of chemical reacting must be specified
• Standard heat of reaction, ΔH°, is the heat produced when the number of moles in the balanced equation reacts. in kJ.
• If a formation reaction occurs, ΔH°f, where only one mole of product is the equation, the units will be kJ/mol

Hess’s Law

• When converting coefficients in a balanced equation by multiplying, dividing, or any mathematic conversion, that same conversion must be done to the heat of reaction
• Summarized law
  • If the coefficients are multiplied by a constant, that ΔH°reaction is multiplied by the same constant
  • If two or more equations are added to create an overall reaction, the heats of the equations can be added to obtain the overall reaction heat
• When trying to use Hess’s Law, follow the principles:
  • Focus on complex molecules first
  • focus on atoms and molecules that occur in just one reaction
  • Focus on atom and molecules that are in the overall reaction
  • Focus on finding atoms and molecules to cancel out unneeded ones from equations

Formation Reactions and Heats of Formation

• Formation reactions are when reactants are in their standard state and there is only one mole of product. Examples:
  • Fe (s) + ½O2 (g) → FeO (s)
  • 2K (s) + ½H2 + 2O2 (g) + P (s) → Fe2O3 (s)
• Heat of formation for elements is always zero

Entropy and the Second Law of Thermodynamics

• Entropy is the number of different ways a system can arrange the particles
  • Increasing number of particles increases entropy
  • Increasing volume increases entropy
  • Changing state (liquid to gas) increasing entropy
• Ice has the least entropy while gaseous water has the most
• Entropy is assigned the symbol S with the units J/C
• Standard entropy is S° with the units J/C mol
• Entropy principles
  • Gas formation increases entropy greatly. A higher Δng means a greater entropy
  • If Δng = 0, solid to liquid phase is a high contributor to entropy increase
  • solid formation decreases entropy
  • increase in temperature increases entropy

Gibbs Free Energy, ΔG°

• Free energy change, ΔG°, is the max amount of energry available from any chemical reaction
  • Driven by entropy and enthalpy values
• ΔG° = ΔH° - TΔS°
• Second law of thermodynamics is any physical or chemical change must result in the increase of entropy in the universe

Free Energy and Equilibrium

• When not at standard state, ΔG is used not ΔG°
• ΔG = ΔG° + RT ln(Q)
  • Q is reaction quotient
  • R is 8.314 J/mol K
• ΔG° = -RT ln(K)
  • K is equilibrium constant
• If K > 1, reaction will move forward
• If K < 1, reaction will move reverse

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