Chapter 6: Thermochemistry
Thermodynamics – the study of energy and its transformation
Thermochemistry – chemical reactions and energy changes
Energy = the capacity to do work or to transfer heat
1. Kinetic Energy (motion) = ½(mass)(velocity)2
2. Potential Energy = energy of position or composition
KE = ½mv2
Heat – involves the transfer of energy between two objects due to a temperature
difference
Temperature – random motion of particles
Work – force acting over a distance; converts one form of energy to another form
of energy
Work = force × distance
W = fd
Units of Energy
1. Joule (J) SI unit
2. calorie (cal)
3. British Thermal unit (BTU)
1 cal = 4.184 J
1 BTU = 1.05 kJ
Exothermic – reaction that evolves (gives off) heat
Endothermic – reaction that absorbs heat from the surroundings
State Function/State Property – a property of a system that is determined by specifying
its conditions
-depends only on its initial and final state
-independent of the path by which the change occurs
Enthalpy (H) – sum of the internal energy and the product of pressure and volume
Entropy (S) –thermodynamic function that describes the randomness and disorder
of molecules based on the number of different arrangements available
to them in a given system or reaction
Gibbs Free Energy (G) – enthalpy of the system minus product of the temperature
times the entropy of the system
System – limited part of the universe you are interested in (where attention focused)
Surroundings – everything not in the system
First Law of Thermodynamics – law of conservation of energy; energy of the universe
is constant
1
Internal energy – total energy of the system (what you are working on)
- the sum of the PE and KE for particles in the system
- change in internal energy represented by ΔE
- reactants Ei, products Ef
- ΔE number is the magnitude of change, sign is direction of change
A system can exchange energy with its surroundings in two ways
1. heat (q) = heat is given off by system (-); absorbed (+)
2. work (W) = work done by system on surrounding (-) (gas expansion)
work done on system by surroundings (+) (gas contracts)
ΔE = q + W
P-V Work = pressure – volume work (expansion and contraction of gas); expanding gas
works on surroundings by pushing back the atmosphere (negative sign for work)
W = -PΔV
1 L×atm = 101.3 J
liquid
gas
gas
liquid
means q (ΔH) = + ; w = means q (ΔH) = - ; w = +
if volume is constant in a container, ΔV = 0 so E = qv (only in a bomb calorimeter)
if pressure is constant, qp = ΔH (enthalpy)
this is everyday lab conditions
ΔHrxn = Heat of reaction = ΣHproducts – ΣHreactants
If ΔHrxn = (+), heat is absorbed (endothermic)
If ΔHrxn = (-), heat is released (exothermic)
ΔHrxn may not be /mole or /gram; ΔHcomb, ΔHvap, ΔHfus, ΔHf
Enthalpy – (a state function)
1. Extensive property – dependent on amounts of substance
2. Enthalpy change for a reaction is equal in magnitude by opposite in sign for the
reverse reaction
3. Enthalpy change for a reaction depends on the state of the reactants and
products
2
Calorimetry = measurement of heat flow
-ΔH is measured experimentally by measuring the heat flow of a raction at
constant pressure
Heat Capacity – energy required to raise the temperature of a substance by 1°C
Specific Heat Capacity – the amount of heat required to raise 1 gram of substance by 1°C
Molar Heat Capacity – a specific heat capacity given per mole of substance instead of
per gram of substance
heat (q) = mass × specific heat (c) × change in temp
q = m × c × ΔT
***At constant pressure, qp = ΔH***
Bomb calorimetry – used to study combustion reactions (measure ΔH)
qv = ΔE
- sealed vessel with a smaller container inside called a “bomb”
- carried out at constant volume
- ΔE = q + W, W = -PΔV; in rigid container ΔV = 0 so W = 0
- heat released represents internal energy change (ΔU), not enthalpy
Hess’s Law – if a reaction is carried out in a series of steps, ΔH for the reaction will be
equal to the sum of the enthalpy changes for each step
Enthalpy Changes
1. If a reaction is reversed, the sign of the ΔH is reversed
2. If the coefficients in an equation are multiplied by an integer, the value
of ΔH is multiplied by the same integer
Heats (enthalpy) of Formation (ΔHf) – the formation of one mole of a compound from
the elements that compose it
ΔHf° = standard (enthalpy) heat of formation
= degree symbol indicates that conditions for the reaction are standard
ΔHf° = 1 atm and 298 K = standard heat of formation
Standard state for compound = 1 atm, condensed state in pure
liquid or solid, solution 1 M
Standard state for element = 1 atm, 25°C, form in native state
ΔHf° of any element in usual (stable) form = 0
ΔH° = ΣnΔHproducts – ΣnΔHreactants
Ways to calculate ΔH:
1.
2.
3.
4.
ΔH = q = cmΔT
ΔH = ΣΔHprod – ΣΔHreact
Hess’s Law
ΔH = E – W = E + PΔV
3
Present Sources of Energy
-historically used energy sources (pre-Industrial Revolution) were animal fats (esp. whale oil)
-first oil well in 1859 in Pennsylvania (kerosene)
Petroleum – thick, dark liquid; mostly hydrocarbons (carbons and hydrogens mainly)
Natural Gas – mainly methane (CH4), but also ethane, propane, and butane; associated
with petroleum deposits
- very plentiful, but hard to obtain
- fracking (hydraulic fracturing)
Coal – plant matter buried and subjected to high pressure and heat over long periods of
time
- 23% of our energy
- expensive and dangerous to mine
- many pollutants put off
New Energy Sources
Coal Conversion – convert to a gaseous form (coal gasification)
- use coal slurries instead of solid coal (similar to residual oil)
Hydrogen – only product of H2 combustion is water (no CO2)
- high cost for production, storage, transport
- virtually no H2 exists as free gas (get ours from treatment of natural gas)
- all reactions to get H2 are extremely high in energy
Other Energy Alternatives
- oil shale = kerogen (not fluid, can’t be pumped); must heat rock to 250°C
- ethanol = fermentation of crops (not good for low temperatures)
- methanol = used in race cars; from natural gas and coal; low energy
- seed oil = sunflower oil used in some to replace diesel
4