First Law of
Thermodynamics
Thermochemistry
Enthalpy
Enthalpies of
Reaction
Hess’s Law
Foods and Fuels
Enthalpies of
Formation
History of thermodynamics



Study of energy and its
transformations
Study began during industrial
revolution to develop relationships
among heat, work, and fuel in
steam engines.
Examines relationships between
chemical reactions and energy
changes that involve heat.
Nature of Energy (definitions galore!)

Define energy.

Define work.

Define heat.
Kinetic vs. Potential Energy



Objects possess kinetic energy. How
can we represent kinetic energy in
an equation?
How does potential energy differ
from kinetic energy?
How is potential energy converted
to kinetic energy?
Potential Energy under the
microscope


Gravitational forces play little to no role in
interactions between atoms and
molecules. Predict what plays a much
more important role.
**electrostatic potential energy**


This energy is proportional to electrical charges
on two interacting objects, Q1 and Q2, and
inversely proportional to the distance, d.
Predict an equation.

Electrostatic Potential Energy (Eel)
Eel = κQ1Q2 / d
κ
is a constant of
proportionality: 8.99 x 109 J-m/C2


How can we make electrostatic potential
energy absolute zero?

At finite
separation
distances for
two charged
particles, Eel is
positive for like
charges and
negative for
opposite
charges. 
Think about it…

If a cyclist is at the top of the hill
(not moving) and stops at the
bottom of the hill, are the potential
energies and kinetic energies the
same at both locations?

Why?
Units of Energy

SI unit for energy is the joule (J) (more
often than not, use kJ)


Using the kinetic energy equation, determine
what 1 J of units would be.
Energy changes associated with
chemical reactions often expressed as
calories (cal)



Originally defined as: amount of energy require
to raise the temperature of 1 g of water from
14.5 C to 15.5 C.
1 cal = 4.184 J (exactly)
Cal does not = cal.
Which is the system, which is the
surroundings?
Open system:
matter and energy
can be exchanged
with surroundings
(ie: boiling water)
Closed System: can
exchange energy but not
matter with surroundings
(ie          
Isolated System: no
exchange of matter or
energy. (ie: coffee
thermos)
Transferring Energy: Work and
Heat


Define Force.
How can we represent work in an
equation?
Example Problem

A bowler lifts a 5.4 kg (12 lb)
bowling ball from ground level to a
height of 1.6 m (5.2 ft) and then
drops it.



What happens to the potential energy
of the ball as it is raised?
What is the quantity of work, in J, used
to raise the ball?
After the ball is dropped, it gains
kinetic energy. If all the work done in
part b has been conver
November 29th, 2012
DO NOW:
 How are potential energy and
kinetic energy related?
 What is the kinetic energy, in J, of
the following:


An Ar Atom moving at a speed of 650
m/s
A mole of Ar atoms moving at 650
m/s?
THE FIRST LAW OF
THERMODYNAMICS

Energy can neither be created nor
destroyed


Any energy lost by a system must be
gained by the surroundings (vice
versa)
Energy is conserved
Internal energy



How do we define internal energy?
Chemistry is mainly concerned with
the change in energy in a system
How do we represent the change in
energy in a system?
Change in Energy



Change in internal energy is denoted:
∆E = Efinal – Einital
Generally we can not determine Efinal
– Einital for a practical system, but we
only need the ∆E to apply the law.
Compare and contrast positive and
negative values of ∆E. What does the
sign change mean?
Example




For the reaction:
2 H2(g) + O2(g)  2 H2O(l)
Which is the initial state and which is
final? Label them on the equation.
As water is created, is energy lost or
gained to the surroundings?
Does this indicate a positive ∆E or a
negative ∆E?
Lets tie it together a little bit.

How may a
system
exchange
energy with
its
surroundings?
Using the previous slide…

How can we algebraically represent
a change in internal energy?
∆E = q + w
As work is done on a system from its
surroundings, w > 0.
 As heat is gained in the system, q > 0.

Example Problem
Calculate the change in internal
energy for a process in which a
system absorbs 140 J of heat from
the surroundings and does 85 J of
work on the surroundings.
Endothermic Vs. Exothermic

Define Endothermic.

Define Exothermic.

Is the formation of H2 gas and O2
gas from H2O considered
endothermic or exothermic? Why?
State Functions

A state function is a property of a
system that is determined only by
the present state of the system, not
the path that system took.

IE:
 Change in sea level
Vs.
Path Driven

How does this relate to ∆E?
ENTHALPY



A system of a gas confined to a
container can be characterized by
several important properties. Which
properties are important?
How are these properties similar to
∆E?
Define Enthalpy.
Why is this equation useful?


Since ∆E involves work, we must
consider the most commonly work
produced by chemical or physical
changes open to the atmosphere.
What is this type of work?
Why is the equation or work we
gave yesterday no longer useful?
Pressure – Volume Work



Work involved in expansion or
compression of gas
When pressure is constant:
 w = -P∆V
If a system does not change its
volume during the course of a
process, does it do pressure-volume
work?
Enthalpy Cont…



When a change occurs at constant
pressure, how can we rewrite our
equation?
When ∆H is positive, that means
the system has gained heat. That
means this is endothermic
When ∆H is negative, that means
the system has lost heat. This
means it is exothermic.
Practice, Practice, Practice!

Complete the worksheet.
November 30th, 2012
Do Now:

Indicate the sign of the enthalpy
change, ∆H, in these processes carried
out under atmospheric pressure and
indicate whether each process is
endothermic or exothermic:
An ice cube melts
 1 g of butane is combusted in sufficient
oxygen to give complete combustion to
CO2 and H2O

Blah, Blah, Blah… so what did it all
mean!

Enthalpy is ∆H= ∆E + P∆V


So what does Enthalpy really tell
us?


HEAT!
So how is this different than ∆E?



MEANING: ∆H = (q + -P∆V) + P∆V
∆E = change in heat at constant
volume
∆H = change in heat at constant P.
The difference between the two is
Enthalpy of Reactions
∆H = Hfinal – Hinitial

How could we re-write this
equation to represent ∆H for a
reaction?

∆H = Hproducts – Hreactants
Reactions


When a numerical value is given for
∆Hrxn , a numerical value along with
the equation must be given.
Suppose you are considering the
reaction seen at the beginning of
class, and that reaction has an
enthalpy of -848. How would we
write this correctly?
Thermochemical Reactions


A thermochemical equation is a
balanced chemical equation that
shows the associated enthalpy
change but does not specify amount
of chemical involved.
The coefficients in the balanced
equation represent the number of
moles of reactants and products
**IMPORTANT GUIDELINES**

Enthalpy is EXTENSIVE.


Enthalpy change for a rxn is equal in
magnitude, but opposite in sign, to ∆H
for the reverse reaction.


If 1 mol of CH4 is reacted with 2 moles of O2
and has an enthalpy of -890 J, how many J
would a reaction of 2 mol CH4 and 4 mol O2
give off?
The reverse rxn of the previously mentioned
rxn would have what value for ∆H?
∆H for a rxn depends on states of
reactants and products
Practice
How much heat is released when 4.50
g of methane gas is burned in a
constant-pressure system? ( 1 mol
CH4 = -890 kJ)
Was this an endothermic or
exothermic process?
Calorimetry
•The value of ∆H can be determined
experimentally by measuring heat
flow accompanying a reaction at
constant pressure.
•Measure of heat flow = calorimetry
•Determined by measure of
temperature change heat flow
produces
Temperature change


Why does the temperature change
felt by certain objects/materials
differ?
How does molar heat capacity differ
from specific heat capacity?
Specific Heat Capacity



How might we solve for specific
heat capacity?
Specific heat capacity values are
slightly different at different
temperatures. Temperature is
generally given.
How can we rearrange our equation
to solve for q?
Practice!

How much heat is needed to warm
250 g of water from 22 degrees
celcius to 98 degrees celcius? What
is the molar heat capacity of water?
December 5th, 2012
Do Now:

How much heat is needed to warm
300 g of water from 20 degrees
celcius to 104 degrees celcius?
What is the molar heat capacity of
water?
Coffee Cup Calorimetry
•Constant Pressure Calorimetry
•Why does the coffee cup act as a
constant pressure calorimeter?
•How do we measure the heat
change in a coffee cup
calorimeter?
•How does the q of the solution
relate to the q of the water
surrounding it?
Example

When a student mixes 50 mL of 1.0 M
HCl and 50 mL of 1.0M NaOH in a
coffee-cup calorimeter, the temperature
of the resultant solution increases from
21.0 C to 27.5 C. Calculate the enthalpy
change for the reaction in kJ/mol HCl,
assuming the calorimeter loses only a
negligible quantity of heat, that the total
volume of the solution is 100 mL, that
its density 1.0 g/mL, and specific heat is
4.18 J/gK
BOMB Calorimetry
Bomb Calorimetry





Constant-Volume Calorimetry
Studies Combustion
May only use after a standardization
measurement
Once you have specific heat
capacity of calorimeter, you may
use:
qrxn = -Ccal X ∆T
Difference between ∆E and ∆H is
very small
Example

The combustion of methylhydrazine
(CH6N2), a liquid rocket fuel, produces
N2(g), CO2 (g), and H2O (l).
When 4.00 g of methylhydrazine is
combusted in a bomb calorimeter, the
temperature of the calorimeter increases
from 25.00 C to 39.50 C. In a separate
experiment the heat capacity of the
calorimeter is measured to be 7.794 kJ/C.
Calculate the heat of reaction for the
combustion of a mole of CH6N2.
Enthalpy of Formation

How can we define Enthalpy of
Formation?

An enthalpy of formation, Hf, is defined
as the enthalpy change for the reaction in
which a compound is made from its
constituent elements in their elemental
forms]
**BE AWARE: enthalpy of
formation values are generally
for 1 atom.
Hess’s Law


If a reaction is carried out in a
series of steps, ∆H for the overall
reaction equals the sum of the
enthalpy changes for the individual
steps.
Example: Calculate the heat of
reaction at 1 atm and 298 K for the
following reaction:

2HI (g) + F2  2HF (g) + I2 (2)
PRACTICE

Complete worksheet