Transcript Thermochemistry

Thermochemistry
AP Chem
Ch. 6
Thermochemistry
Thermochemistry – the study of heat
changes that accompany chemical
reactions and phase changes
Universe = System + Surroundings
 Endothermic Reaction – one in which
energy (heat) is absorbed into the system
 Exothermic Reaction – one in which energy
(heat) is released from the system
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Energy

Energy – the ability to do work or
produce heat
 the SI unit for heat is the Joule (J)
 Heat CANNOT be measured
directly; it is calculated
 “q” = quantity of heat
In the lab…
Calorimeter – an
insulated device
used for measuring
the temperature
change during a
chemical process
Useful Conversion Factors
1 cal = 4.184 J
1 Cal = 1000 cal
101.3 J = 1 L•atm
1 J = 1 kg•m2 / s2
Specific Heat
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Specific Heat – energy required to raise
the temperature of one gram of any
substance one degree Celsius.
Each substance has its own specific heat
 The higher the specific heat, the longer
it will take to raise the temp.
 water has a very high specific heat
4.184 J/g°C
q = m c ΔT
q = heat absorbed or released
c = specific heat value
m = mass or moles of substance
ΔT = change in temperature
Use this when there is NO change in state
**Make sure units will cancel (usually “c” will
determine what units to use for q, m, DT)**
Heat Lost = Heat Gained

When 2 substances at different temperatures
come into contact with each other, heat is
transferred from the warmer to the cooler
substance until both substances are at the
same temperature
Enthalpy
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Enthalpy (H) – measure of the amount of
energy aborbed/released by a reaction
Changes in enthalpy (ΔH rxn )can be
calculated for specific chemical reactions
Exothermic (lose heat) = -ΔHrxn
Endothermic (gains heat) = +ΔHrxn
Enthalpy (cont’d)
ΔH rxn = H products – H reactants
 DH = q only at constant pressure
 ΔH° = standard enthalpy change
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° = occurs under standard conditions

standard conditions = 1atm, 298K (25°C)
Thermochemical Equations
A balanced chemical equation that
includes all reactants, products, and
the change in energy (ΔH)
 Use energy in stoichiometric
calculations

Thermochemical Equations
Can be written 2 ways:
a. 4Fe + 3O2 2Fe2O3 + 1625kJ
or
4Fe + 3O2 2Fe2O3
ΔH = -1625kJ
b. NH4NO3 +27kJ  NH4+ + NO3or
NH4NO3  NH4+ + NO3-
ΔH = 27kJ
Bond Energy
Bond energy = energy required to
BREAK a bond
 Forming bonds = energy is released
 Breaking bonds = energy is absorbed

DHo = bond energies broken – bond energies formed
Heat in Changes of State
 Molar
Heat Values (table 11.5 on
handout)
 ΔH vap = - ΔH cond
 ΔH fus = - ΔH solid
Heat in Changes of State
q = mDH
m = mole or mass
DH = molar heat value
**units for m depend on units of DH**
Heating/Cooling Curve
Changes in Temp & State
Use Heating/Cooling Curve
 Use q=mcDt and q=mDH to calculate
total heat

Heat of Formation


Heat of formation-change in enthalpy that takes
place when a compound is formed from it’s
elements
ΔH°f = standard heat of formation = the change in
enthalpy that accompanies the formation of 1 mole
of compound with all substances in their standard
states


Table 11.6 in handout
all in standard states: 1atm, 298K (25°C)
Heat of Formation
ΔH°rxn = Σ ΔH°f (products) - Σ ΔH°f (reactants)
**ΔH°f of a pure element in it’s standard
state is 0.0 kJ**
Hess’s Law
Add two or more thermochemical equations
together to obtain a desired thermochemical
equation
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Multiply rxn by a
coefficient
Divide rxn by a
coefficient
Reverse rxn
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Multiply DH by same
coefficient
Divide DH by same
coefficient
Change sign of DH