# Topic 5: Energetics

Document Sample

```					Topic 5: Energetics

5.1 Exothermic and endothermic reactions
5.2 Calculation of enthalpy changes
Thermochemistry (Energetics)
   The study of energy involved during chemical
reactions
   Heat:
   the energy of motion of molecules
   All matter has moving particles at stp
   Temperature:
   transfer of heat to a substance because of faster
molecular movement (as long as there is no phase
change)
   A temperature change is explained as a change in
kinetic energy (movement)
   Temperature depends on the quantity of heat (q)
flowing out or in of the substance.
   Energy flowing in the system = Endothermic
   Has a positive value
   Energy entering (feels cool)
   Energy flowing out of the system = Exothermic
   Has a negative value
   Energy exiting (feels warm)
Heat (q)
q=mc ∆t
   q=heat
   m=mass
   ∆t=change in temperature (tf-ti)
   c=specific heat capacity (J (g oC)-1)

Specific heat capacity is the quantity of heat required to
raise the temperature of a unit mass of a substance
by one degree Celsius.
Law of conservation of energy
   ∆E universe = O
   The total energy of the universe is constant, it
is not created or destroyed, however it can be
transferred from one substance to another.
   ∆E universe = ∆E system + ∆E surroundings
First Law of thermodynamics
   Any change in energy of a system is
equivalent by an opposite change in energy of
the surroundings.
   ∆E system = - ∆E surroundings
   According to this law, any energy released or
absorbed by a system will have a transfer of
heat, q.
   So, q system = - q surroundings
Sample Problem
   15 g of ice was added to 60.0 g of water. The
Ti of water was 26.5 oC, the final temperature
of the mixture was 9.7 oC. How much heat
was lost by the water?
   q=mc ∆t
q=(60.0 g) (4.18 J/g oC) (9.7-26.5 oC)
q= - 4213.44 J [-4.2 kJ]
Watch this flash video
http://www.mhhe.com/physsci/chemistry/animat
ions/chang_7e_esp/enm1s3_4.swf
Enthalpy (∆H)
   Total kinetic and potential energy of a system under constant
pressure.
   The internal energy of a reactant or product cannot be
measured, but their change in enthalpy (heat of reaction) can.

∆ H = Hproducts – Hreactants

   A change in enthalpy occurs during phase changes, chemical
reactions and nuclear reactions.

∆ H system = q surroundings
Endothermic Reactions
Method 1: enthalpy level diagram
Endothermic Reactions
Method 2: enthalpy term outside of the
equation
2 HgO (s)  2 Hg (l) + O2(g) ΔH=181.67 kJ

Method 3: enthalpy term within the equation
2 HgO (s) + 181.67 kJ  2 Hg (l) + O2(g)
Exothermic Reactions
Method 1: Enthalpy level diagram
Exothermic Reactions
Method 2:
4Al(s)+ 3O2(g) 2 Al2O3(g) ΔH=-1675.7 kJ

Method 3:
4 Al(s) + 3O2(g) 2 Al2O3(g) +1675.7 kJ

   Neutralization and combustion reactions
Calorimeters (qwater = -qsystem)
   Used to measure the
amount of energy involved
in a chemical reaction.
   To be treated like
   Isolate/closed system
   Specific mass of water used.
   Energy flows to or from the
water in the cups
   Measure the temperature
change related to the water.
Problem
1.   An 25.6 g of an unknown metal with an
initial temperature of 300 oC, is placed in
150.0 g of water with an initial temperature
of 35.0 oC. If the water’s temperature
stabilizes at 55.0 oC, calculate the specific
heat capacity of this metal.
Bomb Calorimeter
Heat Capacity
   Related to bomb calorimeters
   Unit is (J/oC) because its always with a set
mass, so it is redundant to repeat the term
over.
15.1 Standard enthalpy
changes of reaction
Higher level
15.1.1 Define and apply the terms standard state, standard enthalpy
change of formation, and stand enthalpy change of combustion
15.1.2 Determine the enthalpy change of a reaction using standard
enthalpy changes of formation and combustion.
Standard Molar Enthalpy of Formation
   Standard implies the states of the particle at 1 atm
and at 0oC.
   Quantity of energy released (-) or absorbed (+) when
one mole of a compound is formed directly from its
elements at standard temperature and pressure.
   We use a table to find them.
   Unit for ΔHof: kJ/mol
Practice:
   What is the standard molar enthalpy of
formation for the following reaction?
   2 Na (s) + Cl2(g) 2 NaCl (s) + 814 kJ

By definition, the standard molar enthalpy
of formation is for ONE mole of product formed.

∆Hfo = -407 kJ/mol for NaCl
Standard Molar Enthalpy of
Combustion (∆H o     comb

   Energy changes involved with combustion
reactions of one mole of a substance.
   Remember that these reactions are only
measured once cooled to 25oC
   Combustion is a reaction with oxygen as a
reactant (burning)
   Will need a table of values to use.
Combustion reaction with alkanes:
   Always form water and carbon dioxide.
   Ex: CH4 + 2O2  CO2 + 2H2O
   Meth: C=1, Eth: C=2, Prop: C=3, But: C=4,
Pent: C=5, Hex: C=6, Hept: C=7, Oct: C=8,
Non: C=9 and Dec: C=10.
Standard heats of reactions (∆H o)                 rxn

standard states. (if a solution, concentration = 1M)
   All elements at standard state ΔHof = 0
   Most compounds have a negative ΔHof
   Use balanced equation, where n= number of moles

H rxn   nH products   nH reactants

f

f
Calculating enthalpy changes
   Amount of a substance reacting matters, so
can use q= nΔH.
   Remember n=amount of moles.
   If you are given a mass (g) and molar mass
(g/mol), then you can solve for n by dividing
mass by molar mass. (review from Topic 1
stoichiometry section)
Practice:
    Calculate the ΔHorxn for:
1.   4NH3 (g) + 5O2 (g)  4NO (g) + 6H2O (g)
2.   CO (g) + H2O (g)  CO2 (g) + H2 (g)
    Calculate the ΔHocomb for:
1.    2CH3OH(l) + 3O2(g)  2 CO2(g) + 4H2O(l)
2.   2C2H6 (g) + 7 O2 (g)  4CO2 (g) + 6 H2O (l)
Practice:
1/8 S8 (s) + H2 (g)  H2S (g) ∆H o= -20.2 kJ
rxn

a) Is this an endo or exothermic reaction?

b) What is the ∆H o for the reverse reaction?
rxn

c) What is the ∆H when 2.6 mol of S8 reacts?

d) What is the ∆H when 25.0 g of S8 reacts?
Don’t Forget to…