Post on 31-Dec-2015
Quick Review of Concepts
We have been introduced to heat producing (exothermic) reactions and heat using (endothermic) reactions
Heat is a measure of the transfer of energy from a system to the surroundings and from the surroundings to a system
Enthalpy
The heat given off or absorbed during a reaction (change in heat of a system) is called the change in enthalpy (ΔH) when the pressure of the system in kept constant
DH = H (products) – H (reactants)
Calorimetry
We measure the transfer of heat (at a constant pressure) of a chemical reaction by a technique called calorimetry
In calorimetry: The heat released by the system is equal to
the heat absorbed by its surroundings The heat absorbed by the system is equal
to the heat released by its surroundings The total heat of the system and the
surroundings remains constant First Law of Thermodynamics
Calorimetry
We use an insulated device called a calorimeter to measure this heat transfer
A typical device is a “coffee cup calorimeter” Reaction is open to the atmosphere
Therefore, constant pressure q = ∆H at constant pressure
Calorimetry
To measure ΔH for a reaction using a coffee-cup calorimeter:1.Dissolve the reacting
chemicals in known volumes of water
2.Measure the initial temperatures of the solution
3.Mix the solutions4.Measure the final temperature
of the mixed solution
Calorimetry
The heat generated by the reactants is absorbed by the water
We know the mass of the water, mwater
We know the change in temperature, ∆Twater
We also know that water has a specific heat of Cwater = 4.18 J/°C-g.
We can calculate the heat of reaction by:qsys = ∆H = −qsurr
qrxn = ∆H = −qwater
qrxn = ∆H = -(mwater × Cwater × ∆Twater)
Practice Problem
• When 25.0 mL of water containing 0.025 mol of HCl at 25.0°C is added to 25.0 mL of water containing 0.025 mol of NaOH at 25.0°C in a coffee cup calorimeter, a reaction occurs. Calculate ∆H (in kJ) of one mole of water during this reaction if the highest temperature observed is 32.0°C.
• Note:
• Assume the densities of the solution are 1.00 g/mL
• Assume all heat was absorbed by water
Summary
Heat is a measure of the transfer of energy from a system to the surroundings and from the surroundings to a system
The change in heat of a system is called the change in enthalpy (ΔH) when the pressure of the system in kept constant
We measure the transfer of heat (at a constant pressure) of a chemical reaction by a technique called calorimetry
We use an insulated device called a calorimeter to measure this heat transfer
H2O (s) H2O (l) DH = + 6.01 kJ
Is DH negative or positive?
System absorbs heat
Endothermic
DH > 0
6.3
Thermochemical Equations
Is DH negative or positive?
System gives off heat
Exothermic
DH < 0
Thermochemical Equations
CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l) DH = -890.4 kJ
Thermochemical Stoichiometry Enthalpy is STOICHIOMETRIC!
Recall that the stoichiometric coefficients always refer to the number of moles of a substance
You would read the above equation as:“6.01 kJ are absorbed for every 1 mole of ice that melts at 00C and 1 atm”
“890.4 kJ are released for every 1 mole of methane that is combusted at 250C and 1 atm”
H2O (s) H2O (l) DH = 6.01 kJ
CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l) DH = -890.4 kJ
6.3
How much heat is evolved when 266 g of white phosphorus (P4) burn in air?
P4 (s) + 5O2 (g) P4O10 (s) DH = -3013 kJ
266 g P4
1 mol P4
123.9 g P4
x 3013 kJ1 mol P4
x = 6470 kJ
Practice Problem