Thermochemistry
 The study of energy changes in chemical reactions
Nature of Energy
Two types of energy
1) Kinetic energy- energy of motion; atoms and molecules possess KE
KE = ½ mv2
2) Potential energy- stored energy; found within the bonds of substances
**********potential energy can convert to kinetic energy***********
Energy is the ability to do work or produce heat
 Work -moving an object against a force
 Heat- energy transferred between object b/c of a temperature
difference
o Always moves from hot to cold
State Function
 A property of a substance that depends only on its present state
 Independent of the pathway taken or how you get from point A to
point B
 Examples: internal energy, enthalpy, temperature, pressure
 Non examples: distance traveled, work, heat
Think of the Universe being broken down into:
System
part you are concerned
with
Example: a reaction in a beaker
Surroundings
the rest
System: reactants(chemicals)
Surroundings: the beaker and everything beyond it
Types of Reactions
Endothermic Reactions
 System absorbs E form the surroundings
 Temperature of the surroundings lowers
 products are generally less stable (weaker bonds) than reactants
 ΔH is positive (+)
 phase changes of melting, vaporization and sublimation are examples
Exothermic Reactions
 System releases E into the surroundings
 Temperature of the surroundings raises
 products are generally more stable (stronger bonds) than reactants
 ΔH is negative (-)
 phase changes of freezing, condensation are examples
Every Energy (E) measurement has 3 parts:
1. A unit : Joules (J) or calories (cal))
 SI unit is the joule; 1 J = 1 kgm2/s2
 Usually express in kilojoules
 a calorie is equal to amount of heat needed to raise the
temperature of 1 gram of water by 1C; not a very common unit
used anymore
 Conversion factor between joules and calorie
4.184J= l cal
*********1 nutritional Calorie = 1000 cal or 1 kcal*********
2. A magnitude
3. A sign to tell direction
Three Laws of Thermodynamics
1. Law of Conservation of Energy (1st law of thermo)
 energy is can neither be created nor destroyed but converted from 1
form to another
 energy lost by the system is gained by the surroundings and vice versa
∆E=q+w
∆E =internal energy- sum of all PE and KE of system ( Ef – Ei)
q = heat (q is positive in endothermic reactions: heat added and negative in
exothermic reactions: heat removed)
w = work (w is positive if work is done on the system and negative if
the system does work)
W = - P∆V
units of liter-atm ( L· atm)
1L· atm = 101.3 J
Example 1
If a gas expands from 46 L to 64 L at a constant pressure of 15 atm. How
much work is done?
Example 2
What is the internal energy if 50 J of heat is added and 20 J of work are done
on the system
Enthalpy (H) is used to quantify the heat flow into or out of a system in a
process that occurs at constant pressure
ΔHreaction = H products - Hreactants
 ΔHrxn = (+)  endothermic
ΔHrxn = (-)  exothermic
 Usually expressed as kilojoules per mole (kJ/mol)
Thermochemical Equations
 Used to calculate the enthalpy released or absorbed in a chemical
reaction

The stoichiometric coefficients always refer to the number of moles of a

substance
ΔHrxn = 6.01 kJ/mol
Example: H2O (s) → H2O (l)
ΔH = 6.01 kJ

If you reverse a reaction, the sign of ΔH changes
Example: H2O (l) → H2O (s)
ΔH = -6.01 kJ

If you multiply both sides of the equation by a factor n, then ΔH must
change by the same factor n.
Example: 2H2O (s) → 2H2O (l) ΔH = 2 mol x 6.01 kJ/mol
• The physical states of all reactants and products must be specified in
thermochemical equations.
Example: H2O (s) → H2O (l)
ΔH = 6.01 kJ
H2O (l) → H2O (g)
ΔH = 44.0 kJ
Example 3
How much heat is evolved when 266 g of white phosphorus (P4) burn in
air?
Example 4
How much heat is evolved when 4.03 g of hydrogen is reacted with an
excess of oxygen?
Standard Enthalpies of Formation (Δ°Hf)
 Is the change in enthalpy that accompanies the formation of 1 mole of
a compound from its elements in their standard states
 standard states for an element is1 atm and 25°C ( solids and liquids in
normal state)
 There is a table of standard heats of formation Hf for reference in
the back of your book
 The standard enthalpy of formation of any element in its most stable
form is zero.
Examples: ΔHf 0 (O2) = 0
ΔHf 0 (O3) = 142 kJ/mol
ΔHf 0 (Cgraphite) = 0
ΔHf 0 (Cdiamond) = 1.90 kJ/mol
To Calculate Standard Enthalpy
 need to be able to write the a formation equation showing the
formation of 1 mole of a compound from its elements in their standard
states
Example 5
What is the equation for the formation of NO2?
Example 6
Write the equation for the formation of methanol, CH3OH
The standard enthalpy change for any reaction can be found using this very
important equation
ΣΔ°Hf products - ΣΔ°Hf reactants = °Hreaction
Example 7
Benzene (C6H6) burns in air to produce carbon dioxide and liquid
water. How much heat is released per mole of benzene combusted?
The standard enthalpy of formation of benzene is 49.04 kJ/mol.
2C6H6 (l) + 15O2 (g) → 12CO2 (g) + 6H2O (l)
Given the standard enthalpy of reaction , use the standard enthalpies of
formation to calculate the standard enthalpy of formation of CuO
CuO(s) + H2(g)  Cu(s) + H2O(l) ΔH°reaction = -129.7 kJ
Hess’s Law
 Is a state function- independent of the path
 A reaction can be carried in a single step or in a series of steps.
 If carried out in a series of steps, we can add equations to come up
with the desired final product and thus add ΔH for the overall reaction

One step:

Two step:
N2(g) + 2 O2(g)  2NO2(g)
ΔH = 68 kJ
N2(g) + O2(g)  2NO(g)
ΔH =180 kJ
2NO(g) + O2(g)  2NO2(g)
ΔH =-112 kJ
_________________________________________________
N2(g) + 2 O2(g)  2NO2(g)
ΔH = 68 kJ
Two Rules to Remember:
 If the reaction is reversed, the sign of ΔH is reversed
N2(g) + 2 O2(g)  2NO2(g)
ΔH = 68 kJ
2NO2(g)  N2(g) + 2 O2(g)
ΔH = -68 kJ
 The magnitude of ΔH is directly proportional to the moles of
reactants and products. If the coefficient of a reaction are multiplied
by and integer, the ΔH is also multiplied by the integer
******when using Hess’s law, work by adding the equations to make it look
like the answer, the other parts will cancel out******
Example 8
C(s) + 2 H2(g) CH4(g)
C(s) + O2(g)  CO2(g)
ΔH1 = -393.5 kJ
H2(g) + ½ O2(g)  H2O(l)
ΔH2 = -285.8 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
ΔH3 = -890.3 kJ
ΔHrxn = ?
Example 9
Given O2(g) + H2(g)  2 OH(g)
ΔºH = +77.9 kJ
O2(g) 2O(g)
ΔºH = +495 kJ
H2(g) 2H(g)
ΔºH = +435.9 kJ
Calculate the ΔºH for this reaction
O(g) + H(g)  OH (g)
Calorimetry
 process that measures the transfer of heat
 use a calorimeter( an insulated container): usually the change in
temperature of water is measured but often the heat capacity of the
calorimeter is known since the calorimeter also absorbs some heat as
well as the water in the device.
Heat Capacity (C)
 Amount of heat energy needed to raise temperature by 1º Celsius
C =__q
∆T
Specific Heat Capacity (c) or (s)
 amount of energy needed to raise 1 gram of a substance by 1º Celsius
c=_ q
rearranges to q= mc∆T
m · ∆T
 Units: J/g·°C or J/g·K or cal/g·ºC or cal/g· K
 Each substance has its own specific heat
Example 10
The specific heat of graphite is .71 J/gºC. Calculate the energy needed to
raise the temperature of 75 Kg of graphite from 294 K to 348 K.
Example 11
How many joules of heat will raise the temperature of exactly 50.0 g of
water at 25 ºC to 75 ºC
Molar Heat Capacity
 amount of heat required to raise 1 mole of substance by 1°C
C = ___q
or
Cp=∆H/∆T
n ·∆T
rearranges to q = nc ∆T
Example 12
What is the molar heat capacity for water?
Coffee Cup Calorimetry (constant pressure)
 used to determine ΔH for reactions in solution
 uses an insulated cup full of water as the calorimeter
 use q= mc∆T to solve
 Water has a specific heat of 4.184 J/g·C
 Assume:
Heat released in a reaction = heat absorbed in water
Negative heat
positive heat
 qrxn + qsoln = 0
 qrxn = - qsoln if endothermic
 qsoln = - qrxn if exothermic
Example 13
A 46.2 g sample of copper is heated to 95.4 ºC and then placed in a
calorimeter containing 75.0 grams of water at 19.6 ºC. The final
temperature of both the water and the copper is 21.8 C. What is the
specific heat of copper?
Example 14
Suppose you place .500 g of magnesium chips into a coffee-cup calorimeter
and then add 100.0 ml of 1.00 M HCl. The reaction occurs is
Mg(s) + 2HCl (aq) --> H2(g) + MgCl2(aq)
The temperature of the solution increases from 22.2 to 44.8º C. What is the
enthalpy change per mole of magnesium. Assume specific heat capacity of
the solution is 4.20 J/g·K and the density of the HCl solution is 1.00 g/ml.
Example 15
Assume you mix 200. ml of .400 M HCl with 200. ml of .400M NaOH in a
coffee-cup calorimeter. The temperature of the solutions before mixing was
25.10 C; after mixing and allowing the reaction to occur, the temperature is
26.60C. What is the molar enthalpy of neutralization of the acid? Assume
the densities of the solution are 1.00g/ml and their specific heat capacities
are 4.20J/gK.
Bond Energies and Enthalpy
Bond energy is required to break a bond.
 Since the breaking of a bond is an endothermic process, the bond
energy is always a positive number.
 When a bond is formed, energy equal to the bond energy is released.
Σ Bond energy of Reactants – Σ Bond energy of Products = Δ°Hrxn
Example 16:
Bond
H–H
O=O
O–H
Bond Energy(kj/mol)
436
499
463
Find the Δ°Hrxn for the following reaction
2 H2(g) + O2(g)  2H2O(g)