Thermochemistry

• Types of Energy Relevant to Chemistry

  • potential
  • chemical
  • nuclear
  • kinetic
  • thermal
  • radiant
  • electrical
• Potential energy: energy associated with an objects position or composition
• Chemical energy: energy stored in intramolecular and intermolecular bonds; released in chemical reactions
• nuclear energy: energy stored in the collection of neutrons and protons in the atom's nucleus
• Kinetic energy: energy associated with the random kinetic motion of atom and molecules
• Radiant Energy: energy in the form of radiation
• Electrical Energy: energy associated with the flow of charges
• System

  • specific part of the universe of interest in a specific application
  • surroundings: things around a system
• Internal energy

  • the capacity of a system to do work
  • E = sum of all energies
  • ∆Esys = Ef - Ei
  • absolute internal energy is not measurable
  • can measure changes in the energy of a system
  • system = chemical reaction
• Conservation of Energy

  • energy of the universe is constant
  • energy can't be created or destroyed
  • Esystem lost = E surroundings gain
  • ∆Esys + ∆Esurroundings = ∆Euniverse = 0
  • ∆Esys = -∆Esurroundings
• Chemical Energy

  • reactants --> products: energy released to the surroundings (heat/work)
  • ∆Esys = Eproducts - Ereactions < 0
  • ∆Esurroundings = -∆Esys > 0
  • when ∆Esys is negative the system loses energy as heat/work to the surroundings
  • reactants --> products (upwards): energy from surroundings to system
  • ∆Esys = Eproducts - Ereactions > 0
  • ∆Esurroundings = -∆Esys < 0
• Work and Energy

  • work: by or on the system is energy transfer that results in macroscopic changes in the system
  • w = F * D: (Joules) = Newtons * meters
  • ∆Esys = q + w: heat transfer + work done
• Mechanical work

  • w = F*∆h: pressure cylinder-system
  • w = -P∆V
  • 1 L*atm = 101.325 J
  • Closed-system:
  • when cylinder expands ∆V is (+) (Vf - Vi)
  • work must be (-) because ∆Esys is (-), gas is losing energy
• Heat ≠ Temperature

  • temperature: a measure of the random molecular motion in an object or substance
  • Ke avg = 3/2 RT
  • heat: energy transfer between 2 objects often driven by a temperature difference
  • q = m*C*∆T
  • w = -P∆V
• Quantifying Heat

  • heat (q) --> system ∆Temp
  • heat capacity: (C) = amount of heat (q) required to raise the temp of entire system by 1 C
  • specific heat capacity (Cs) - the amount of heat (q) needed to raise the temperature of 1 gram of the substance by 1 C
  • Molar heat capacity (Cm) = same, but for 1 mol of substance
  • q = mC∆T = mCs∆T = nCm∆T
  • ∆T = Tfinal - Tinitial
• endothermic: ∆H and q (+)
• exothermic: ∆H and q (-)
• Heat (q) +/-

  • positive: gains from surroundings
  • negative: loses heat to surrounding
• work (w) +/-

  • positive: surroundings do work on system
  • negative: system does work on surroundings
• Change in Energy ∆E +/-

  • positive: energy surroundings flow into system
  • negative: energy flows out of the system into surroundings
• State functions

  • depends on the present state
  • energy is state function
  • heat and work are not state functions
  • ∆E = Ef - Ei
  • pressure, volume, and temperature are state functions
  • ∆H = Hfinal - Hinitial
  • ∆H = H (products) - H (reactants)
• constant volume

  • ∆E = q (volume)
  • chemical reaction at constant volume, ∆V = 0
  • Change in internal energy = heat exchanged when V is constant
• Constant-pressure Calorimetry

  • coffee cup calorimeter measure ∆Hrxn = q (rxn) at a constant pressure
  • heat flows from reactants to solution or vice versa
  • heat of reaction has opposite sign from heat of solution
  • q (rxn) = -q (solution)
  • physical measurement: temperature change of solution ∆T
  • ∆H = q (pressure)
• Calorimetry

  • is the measurement of heat of a reaction (q)
  • physical measurement is change in temperature (∆T)
  • intangible quantity is heat (q)
• Bomb Calorimeter

  • measures ∆Erxn = q (rxn) at a constant volume
  • heat of reaction has opposite sign form heat of calorimeter q (cal) = -q (rxn)
  • physical measurement: temperature change of calorimeter
• properties of enthalpies of reaction

  1. enthalpies of reaction (∆Hrxn) are extensive: depends on the amount of reactants
  2. ∆Hxn for backwards reaction changes the sign of ∆Hrxn for forwards reaction (∆H2 = - ∆H1)
  3. Hess's Law
• Hess's Law

  • ∆Hrxn is the sum of enthalpy change of each step of a stepwise reaction (∆Hrxn = ∆H1 + ∆H2 + ∆H3)
  • state function
  • H is not path-dependent
  • only depends on initial and final states
• Standard state of elements

  • pure substance in its most stable state (Ex. O2, noble gases)
  • pressure of 1 atm
  • temperature of 25 C
• Standard states of gases: pressure of pure gas at 1 atm
• Standard states of liquids or solids

  • pure substance in its most stable state
  • pressure of 1 atm
  • temperature of 25 C
• standard states of solutions

  • concentration of exactly 1M (moles/Liter)
• Standard enthalpy

  • all reactants and products are at standard stage
  • pure compounds: 1 mol forms from constituent elements in their standard state
  • pure elements: the constituent element of an element is itself Hf = 0 (Ex. for O2 standard enthalpy is 0)
• Bond energy

  • energy associated with breaking a specific chemical bond
  • endothermic process (positive ∆H)
  • ∆Hbond > 0
• Bond Enthalpy Values

  • in reaction, reactant bonds are broken and new bonds are formed in product
  • bond energy values are averages of bond types in all known molecules
  • form (-)
  • break (+)
• bond breaking and formation

  • bonded + (break) = +∆H
  • unbonded + bonded = -∆Hbond
• Strong bond and Broken bond: large and positive ∆H
• Weak bond and Broken bond: Small and positive ∆H
• Strong and formed bond: Large and negative ∆H
• Weak and Formed bond: small and negative ∆H