Presentation to support laboratory experiment on Enthalpy.

1. Enthalpy of Reaction
General Chemistry Labs

2. Theory
• Change in enthalpy, ΔH, is the
measurement of the energy flow in a
reaction
• Chemical reactions absorb or release
energy, in the form of heat
• Reactions
– Exothermic (heat is released, Δ H < 0)
– Endothermic (heat is absorbed, Δ H > 0)

3. Calorimetry
• The absorption or release of heat can be
measured by monitoring the change in the
temperature of the surroundings.
• In order to determine the total amount of
energy needed for the reaction, the heat
used or released from the reaction must be
contained within the reaction solution.

4. Calorimeter
• A calorimeter is used to prevent the escape
of heat.
• In this lab, the calorimeter being used is a
Styrofoam cup that sets in a glass beaker
and has a cardboard lid.

5. Determination of heat (q) from Δ T
• The relationship between the ΔT and heat is dependent
on the properties of the solution
• Some solutions will increase in temperature more readily
than others even through they are absorbing the same
amount of heat.
• The specific heat for a substance is a constant that
describes how readily the temperature will increase when
heat is absorbed.
• s = specific heat of water, units of
𝐽
(𝑔∗°𝐶)

6. q = msΔT
• q = heat, units of Joules
• m = mass, units of grams
• s = specific heat of water, units of
𝐽
(𝑔∗°𝐶)
• Δ T = change in temperature, units of °C

7. q and ΔH
• q is the amount of heat lost or gained in the
total reaction.
– The units are Joules and it is an extensive
property,
– not dependent on the amount of the
substance.
• Δ H is an intensive property that is typically
expressed as the amount of heat per mole,
or kJ/mol.

8. Relationship of q and Δ H
• Therefore,
q
# moles
= Δ H
• WATCH the units!!!
• Δ H will be expressed as a negative or positive
dependent on the direction of heat flow.

9. Exothermic or Endothermic
• The positive or negative sign shows the direction of the
heat flow in a reaction
• Negative: heat produced by the reaction, and heats
surrounding areas
– Exothermic
• Positive: heat is consumed by the reaction, and therefore
cools the surrounding areas.
– Endothermic
• The (+) or (-) signs are assigned based on the direction of
heat flow not by the mathematical equation,
Δ H =
𝑞
𝑚𝑜𝑙
=
𝑚𝑠Δ 𝑇
𝑚𝑜𝑙

10. Calculating Enthalpy of Reaction
• Indirect Method
ΔHrxn = SnΔHf(products) - SmΔHf(reactants)
• Δ Hf for most common substances are
found in tables

11. Hess’s Law
– The enthalpy of a reaction does not depend
on number of steps involved
– If the enthalpy of a set of reactions is known,
• combining the enthalpy of these reactions then
enthalpy of a reaction of interest can be obtained
• Rearrange given reactions to get the overall
reactions. Whatever we do to the reaction we also
do to the enthalpy

12. Reaction Δ H (kJ/mole)
C(gr) + O2(g) → CO2(g), Δ H0
rxn1 = -393.5
S(rh) + O2(g) → SO2(g), Δ H0
rxn2= -296.4
CS2(l) + 3O2(g) → CO2(g) + 2SO2(g) Δ H0
rxn3= -1073.6
Determine the enthalpy of the reaction
shown using the three reactions given.
C(graphite) + 2 S(rhombic) → CS2(l)
Hess’s Law : Example

13. C(gr) + 2S(rh)  CS2(l)
C(gr) + O2(g) → CO2(g)
S(rh) + O2(g) → SO2(g)
CS2(l) + 3O2(g) → CO2(g) + 2SO2(g)
Therefore, rewrite equations
C(gr) + O2(g) → CO2(g)
2 S(rh) + 2 O2(g) → 2 SO2(g)
2 SO2(g) + CO2(g) → CS2(l) + 3 O2(g)
2 x
-1 x

14. C(gr) + 2S(rh)  CS2(l)
Compounds in equal quantities on both sides of
the reaction arrow cancel out
C(gr) + O2(g) → CO2(g)
2S(rh) + 2O2(g) → 2SO2(g)
2SO2(g) + CO2(g) → CS2(l) + 3O2(g)
C(gr) + 2S(rh) → CS2(l)
ΔH0
rxn = Δ H0
rxn1 + 2 Δ H0
rxn2 - Δ H0
rxn3
= 87.3 kJ/mol

15. Procedure
Put cardboard square on top and place
temperature probe through hole in cardboard.
Calorimeter

16. Finding Enthalpy
• Mix reactants
• Measure temperature change
• Calculate heat change
• Calculate change in enthalpy (ΔH)

17. Calculating Change in
Temperature
• Include the graphs of temperature
vs. time data for all three reactions
• Δ T
– Initial temperature: Right before
solutions are mixed
– Final temperature: Maximum
temperature in data
• Let experiment continue until the
temperature is no longer increasing

18. Experimental Results
• Enthalpy values for reactions 1, 2,
and 3
• Two experimental values for
reaction 3
– Value from running the reaction
– Using Hess’ law from reactions 1 & 2
• Accepted values are from pre-lab

19. Data Analysis
• Calculate Δ H0 for each reaction, in
terms of kJ/mol for one reactant
– Use the volume and molarity of one
reactant to calculate the number of
moles that were actually used in the
experiment.
– Each of the 3 reactions will have
different Δ H0 values

20. Data Analysis
• Calculate % error in Δ H0 for reaction
3 for both experimental values of Δ H0
– Do not average the two enthalpy values
for reaction 3
• The accepted value of Δ H0 for
reaction 3 was obtained in the pre-lab
exercises

21. Practice Question #1
• Consider two metals A and B.
Each metal has a mass of 250 g
and an initial temperature of 25°C.
The specific heat of B is larger
than that of A.
• Under the same heating
conditions, which metal would take
longer to reach a temperature of
26°C?

22. Practice Question #2
• A 15.0-g sample of nickel metal is
heated to 100.00 oC and dropped
into 55.0 g of water, initially at
23.0oC.
• Assuming that all the heat lost by the
nickel is absorbed by the water,
calculate the final temperature of the
nickel and water.
• (The specific heat of nickel is 0.444
J/oC g)

23. Practice Question #3
Given the following data:
2O3(g)  3O2(g) Δ H0 = -427kJ
O2(g)  2O(g) Δ H0 = +495kJ
NO(g) + O3(g)  NO2(g) + O2(g) Δ H0 = -199kJ
Calculate Δ H for the reaction
NO(g) + O(g)  NO2(g)