Key terms
Thermodynamics – study of energy and its interconversions
Energy – capacity to do work or produce energy
Law of conservation of energy – energy can be neither
created nor destroyed but only changed in form.
Potential energy (PE) – energy due to position or
Kinetic energy (KE) – energy due to motion
Heat – transfer of energy between objects of different
Work – force acting over a distance
Pathway - the way the energy is transferred
State function/state property – depends on current
(present) state (ex = energy)
Heat transfer
Exothermic = energy given off; heat flows out of the system
Endothermic = energy absorbed; heat flows into the system
Where does the energy come from?
– Difference in potential energies between reactants and
– Energy gained by the surroundings must be equal to the
energy lost by the system (reaction)
– Exothermic reaction - lowers potential energy; reactants
higher than products
– Endothermic reaction – increase potential energy;
reactants lower than products
– See energy diagrams
Energy Diagrams
Exothermic Reaction
Endothermic Reaction
• Feels warm
• Change in energy
negative - ∆E < 0
• Feels cool
• Change in energy positive ∆E > 0
Internal energy in a system – E = sum of
potential and kinetic energies of all the particles
in the system
∆E = q + w
∆E = change in systems internal energy
q = heat
w = work (system)
Or if we look at surroundings
∆E = q – w’
w’ = work (surroundings view)
Units for E = Joules - ?
Enthalpy = H (a state property)
H = E + PV
E = internal energy of system
P = pressure of system
V = volume of system
∆H = qp
At constant pressure, the change in enthalpy
(heat of reaction) is equal to energy flow as heat.
∆H = Hproducts – Hreactants
∆H positive = endothermic
∆H negative = exothermic
Calorimetry Key Terms
Calorimetry – science of measuring heat
Calorimeter – device used to determine heat of
reaction; p.255 for diagram
Constant pressure calorimetry – atmospheric
pressure constant
used for determining change in enthalpy
(heats of reaction) ∆H = qp = q
Constant volume calorimetry – use a bomb
∆E = qv = q
∆E = ∆T * heat capacity of calorimeter
Specific Heat
Heat capacity – energy needed to raise
temperature 1 degree of substance
heat absorbed
increase in temp
J/K or J/°C; depends on amount of substance
Specific heat capacity – heat capacity per grams of
substance; J/K*g or J/°C*g
chart on p.254
Molar heat capacity – heat capacity per mole of
substance; J/K*mol or J/°C*mol
q = s x m x ∆T
q = C x m x ∆T
q= heat (energy) ; J
s or C = specific heat (from chart);
m = mass; g
∆T = change in temperature (Tfinal –Tinitial); °C
Hess’s Law
• Change in enthalpy is same whether occurs in
one step or series of steps
• Get a net reaction if occurs in series of steps
Reaction 1
Reaction 2
Net reaction
∆H = H1 + H2
• If reverse the reaction, sign of ∆H reversed
Standard enthalpy of formation
• Change in enthalpy that accompanies
formation of 1 mole of a compound from its
elements in its standard states; ∆H°f
– Under standard conditions/standard states
Enthalpy of formation of compounds = chart
Enthalpy of formation of an element = 0
Apply Hess’s Law to figure out overall ∆H°f
∆H°f = ∑np∆H°f (products) - ∑nr∆H°f (reactants)
Multiply the enthalpies by the coefficients in balanced
Sources of Energy
• Fossil fuels – lead to greenhouse effects
– Coal
– Petroleum
– Natural gas
• New energy sources - costly
– Coal conversion
– Hyodrogen
– Alternate fuels
• Oil shale
• Ethanol
• kerogen