Energy is commonly defined as the capacity to do work or
transfer heat. Work is defined as the energy used to cause and
object to move. Heat is defined as the energy used to cause the
temperature of an object to increase.
Slide 3
Kinetic Energy: Energy of motion KE = mv 2 Example: A car
moving at 55 mph has a greater kinetic energy than a car moving at
40 mph. All atoms and molecules are in motion and have kinetic
energy. Potential Energy: Associated with the position of an
object. Chemical potential energy is the energy stored in chemical
bonds.
Slide 4
The SI unit of energy is the joule (J) This is derived from the
formula for kinetic energy. Another unit of energy is the calorie.
1 cal = 4.184 J
Slide 5
Things in the universe we single out to study and observe are
referred to as systems. Everything else is the surroundings.
Systems can be: Open: Where matter and energy can be exchanged with
the surroundings. Closed: Where energy but not matter can be
exchanged with the surroundings. Isolated: Where neither matter nor
energy can be exchanged with the surroundings.
Slide 6
The two ways that we experience energy changes are work and
heat. Work is defined as the energy used to move an object. = F x d
The other way to transfer energy is heat. Heat is the transfer of
energy from a warmer object to a cooler one.
Slide 7
The first law of thermodynamics states that energy cannot be
created or destroyed, only transferred. Any energy lost by the
system must have been transferred to the surroundings.
Slide 8
The internal energy of a system is the sum of ALL kinetic and
potential energy of all of its components. The internal energy of a
system is defined as E. We can usually never calculate the actual
numerical value of E. We can calculated the change in the internal
energy of a system ( E).
Slide 9
A closed system may exchange energy with the surroundings as
heat or work. This would obviously change the amount of internal
energy resulting in a positive or negative E. The first law of
thermodynamics can be written algebraically as: E = q +
Slide 10
When heat is added to a system or work is done on the system,
its internal energy would increase ( E = +) If heat is released
from the system or the system does work on the surroundings, its
internal energy would decrease ( E = -)
Slide 11
For q+ Means system gains heat - means system loses heat For w+
means work is done on the system - means work is done by the system
For E+ means net gain of energy by system - means net loss of
energy by system
Slide 12
Two gases, A(g) and B(g), are confined in a cylinder and piston
arrangement. Substances A and B react to form a solid product: A(g)
+ B(g) C(s). As the reaction occurs, the system loses 1150 J of
heat to the surroundings. As the volume of the gas decreases under
constant atmospheric pressure, the surroundings do 480 J of work on
the system what is the change in the internal energy of the
system?
Slide 13
When a process occurs in which the system absorbs heat, the
process is called endothermic. A process in which the system loses
heat is called exothermic.
Slide 14
Internal energy is an extensive property. Meaning a 25 g sample
of water at 25 o C would contain less total internal energy than a
50 g sample of water at 25 o C. Suppose we define our system as 50
g of water at 25 o C. This system could have arrived at this state
by cooling 50 g of water from 100 o C, or by melting ice from 0 o
C. The internal energy after each process would be the same. A
state function is a property of a system that depends only on its
current state, not how it got there.
Slide 15
Suppose you drive from Chicago to Denver. Chicago is 596 ft
above sea level; Denver is 5280 ft above sea level. No matter what
rout you take your altitude change will be 4684 ft. Altitude is a
state function. The rout you take however can have a big difference
on the distance your travel.
Slide 16
Consider a battery
Slide 17
Consider the reaction between zinc and an acid Zn(s) + 2 H +
(aq) Zn 2+ (aq) + H 2 (g) If we carry this reaction out in an open
beaker, we can see the hydrogen gas being evolved. It may not be
obvious but the hydrogen gas is doing work on the surroundings.
This is more evident if we add a cylinder and piston set up to the
experiment.
Slide 18
This is the most common form of work discussed in chemistry.
This is called pressure-volume (PV) work. When pressure is constant
PV work is described as: w = -P V The thermodynamic function called
enthalpy accounts for both the follow of heat into or out of a
system as well as PV work. Positive enthalpy indicates a total
increase in energy. Negative enthalpy indicates a total decrease in
energy.
Slide 19
Indicate the sign of the enthalpy change, H, in each of the
following processes carried out under atmospheric pressure, and
indicate whether the process is endothermic or exothermic. (a) An
ice cube melts. (b) 1 g of butane (C 4 H 10 ) is completely
combusted in excess oxygen.
Slide 20
Just like internal energy: H = H final H inital In a chemical
reaction: H = H products H reactants When the combustion of
hydrogen gas is controlled so that 2 mol of H 2 (g) burn to form 2
mol of H 2 O(g) at constant pressure the system releases 483.6 kJ
of energy.
Slide 21
Enthalpy is an extensive property. The magnitude of H,
therefore, is directly proportional to the amount of reactant
consumed. CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O(l) H = -890 kJ 2
CH 4 (g) + 4 O 2 (g) 2 CO 2 (g) + 4 H 2 O(l) H = -1780 kJ
Slide 22
The enthalpy of a reaction is equal in magnitude but opposite
in sign to the enthalpy of its reverse reaction. CO 2 (g) + 2 H 2
O(l) CH 4 (g) + 2 O 2 (g) H = 890 kJ The enthalpy of a reaction
depends on the state of the reactants and products. CH 4 (g) + 2 O
2 (g) CO 2 (g) + H 2 O(g) H = - 802 kJ
Slide 23
How much heat is released when 4.50 g of methane gas is burned
in a constant pressure system?
Slide 24
Calorimetry is the experimental determination of the H for any
physical or chemical process. In order to do a calorimetry
experiment we need to know some background. Heat Capacity (C): Heat
capacity, C, is the amount of heat energy required to raise the
temperature of a substance by 1 K (1 o C).
Slide 25
A substance with a high heat capacity requires more heat energy
to raise its temperature than a substance with a low heat capacity.
Heat capacity is an extensive property. 10 mL of water heat up much
faster than 1 L of water. The heat capacity of one mole of a
substance is called its molar heat capacity C m. The heat capacity
of one gram of a substance is called its specific heat capacity or
just specific heat C s.
Slide 26
C s = Example: It requires 209 J of heat to increase the
temperature of 50.0 g of water by 1.00 K. What is the specific heat
of water?
Slide 27
(a) How much heat is needed to warm 250 g of water from 22 o C
to 98 o C? The specific heat of water is 4.18 J/g-K. (b) What is
the molar heat capacity of water?
Slide 28
Many enthalpies of reactions have been tabulated by scientists.
Having this information available and being able to use it allows
us to calculate the heat of reaction for almost every chemical
reaction. Because enthalpy change ( H) is a state function,
reactions can be thought of as happening in one step or in a series
of steps.
Slide 29
The combustion of methane (CH 4 ) gas to produce carbon dioxide
and liquid water happens in two steps.
Slide 30
The enthalpy of reaction for the combustion of C to CO 2 is:
C(s) + O 2 (g) CO 2 (g) H = -393.5 kJ The enthalpy of reaction for
the combustion of CO to CO 2 is: CO(g) + O 2 (g) CO 2 (g) H =
-283.0 kJ Use this information to calculate the enthalpy of
reaction for the combustion of C(s) to CO(g)
Slide 31
Calculate H for the reaction: 2 C(s) + H 2 (g) C 2 H 2 (g)
Given the following chemical equations and their respective
enthalpy changes C 2 H 5 (g) + 5/2 O 2 (g) 2 CO 2 (g) + H 2 O(l) H
= -1299.6 kJ C(s) + O 2 (g) CO 2 H = -393.5 kJ H 2 (g) + O 2 (g) H
2 O(l)
Slide 32
Many different enthalpies of many different processes have been
tabulated. We have already used the enthalpy of fusion and the
enthalpy of vaporization. One other important type of enthalpy
change that has been tabulated for many compounds in the enthalpy
of formation. This is the heat change associate with the formation
of a compound from its constituent elements.
Slide 33
Enthalpies of formation change depending on things like,
pressure, temperature, and the physical states of the reactants and
products. In order to compare different enthalpies of formations we
need to study then under the same conditions. Most enthalpies of
formation (an of other reactions as well) are tabulated under
standard conditions standard conditions are defined as 1 atm and
298 K. Standard enthalpy change is that of a reaction where all of
the components are in the physical state they would be in at
standard conditions.
Slide 34
The standard enthalpy of formation of a compound ( H o f ) is
the enthalpy change for the reaction that forms one mole of the
compound at standard state. Elements (in standard state) 1 mole of
Compound (in standard state) The value of H o f is always reported
as kJ/mol The standard enthalpy change for any element in its
standard state is 0 kJ/mol.
Slide 35
The enthalpy of any reaction can be calculated from the
addition of the tabulated H o f for all the components. Example:
Calculate the standard heat of reaction for the combustion of
propane (C 3 H 8 ) with oxygen to from CO 2 and water.
Slide 36
A spontaneous process is one that occurs without any outside
assistance. The melting of ice at temperatures higher than 0 o C is
a spontaneous process. Any process that is spontaneous is one
direction is not spontaneous in the reverse direction. Think about
trying to freeze water at 25 o C. Not impossible but not easy. The
majority of spontaneous processes are exothermic. All spontaneous
processes are considered to be irreversible. A reversible process
is one where the system and the surroundings can be converted back
to their original states without any net change.
Slide 37
Now that we know what a spontaneous process is we can start to
predict if an unfamiliar process will be spontaneous or not. We
will use the thermodynamic quantity of entropy (S) to do this. In
general we will define entropy as randomness. The more free the
molecules are to move in random order, the higher the entropy of
the system.
Slide 38
Just like E, and H, entropy (S) is a state function. We can
also not calculate the specific value of S for any system. We can
only calculate the change in S. S = S final S initial In the
special case of an isothermal process S is equal to the heat that
would be transferred in the reverse process divided by the
temperature.
Slide 39
The melting of a substance at its melting point and the
vaporization of a substance at its boiling point are isothermal
processes. We achieve these changes by adding a certain amount of
heat from the surroundings. ( H fus or H vap ) The S for these
processes can be calculated easily because these are isothermal
processes.
Slide 40
The normal freezing point of mercury is - 38.9 o C, and its
molar enthalpy of fusion is H fusion = 2.29 kJ/mol. What is the
entropy change of the system when 50.0 g of Hg(l) freezes at the
normal freezing point?
Slide 41
We saw that in the first law of thermodynamics energy is always
conserved. Entropy is different. The entropy change in any
spontaneous process will always be positive. If we calculate the
entropy change of one mole of ice melting to liquid water we
get:
Slide 42
In the last example we saw that the total entropy change of the
universe was positive. This corresponds to a spontaneous and
irreversible process. If the sum of the S of both the system and
surroundings equal zero the process is reversible.
Slide 43
There are three properties of matter that we can use to predict
the sign (+ or -) of S: Temperature Volume The number of
independently moving particles. Think of these three factors and
how they affect the movement and disorder of molecules.
Slide 44
When water vaporizes the molecules spread out into a larger
volume. This results in S = + Consider the reaction: 2 NO(g) + O 2
(g) 2 NO 2 (g) In this process we are combining two molecules into
one. The result of this is less total movement of molecules in the
system and a negative entropy.
Slide 45
In general we expect entropy to increase when: Gases are formed
from either solids or liquids. Liquids or solutions are formed from
solids. The number of gas molecules increases during a chemical
reaction.
Slide 46
Predict whether S would be negative or positive for each of the
following processes. H 2 O(l) H 2 O(g) Ag + (aq) + Cl - (aq)
AgCl(s) 4 Fe(s) + 3 O 2 (g) 2 Fe 2 O 3 (g) N 2 (g) + O(g) 2
NO(g)
Slide 47
As we cool a substance down the molecules begin to move slower.
If we continue to cool a substance down to absolute zero (0 K) the
molecules will have no kinetic energy. The third law of
thermodynamics states that the entropy of a pure crystalline solid
substance at absolute zero is 0.
Slide 48
We talked about how calorimetry is used to determine the
enthalpy change for a reaction. No such type of experiment exists
to determine the entropy change of a reaction. The absolute value
of S can however be determined through complex
experimentation.
Slide 49
The entropies of substances in their standard states are
usually tabulated as molar quantities: As you can see, unlike
enthalpies of formation, the molar entropy of a pure element is NOT
zero. The standard molar entropies of gases are, in general,
greater than those of liquids and solids. Standard molar entropies
generally increase with increasing molar mass. Standard molar
entropies generally increase with an increasing number of atoms in
the formula of a substance.
Slide 50
Just like using the standard heats of formation to calculate
the enthalpy change in a reaction, we can use the standard molar
entropies in the same way. S o =
Slide 51
Calculate the S o for the synthesis for ammonia from N 2 (g)
and H 2 (g) at 298 K.
Slide 52
The change in entropy of the surroundings during any process
will depend on how much heat is given off or absorbed by the
system. for a reaction at constant pressure
Slide 53
So far we have studied two thermochemical concepts, entropy and
enthalpy. These two quantities must be used together to tell
definitively if a reaction is spontaneous or not. Gibbs free energy
tells us if a reaction is truly spontaneous.
Slide 54
If G is negative, the reaction is spontaneous in the forward
direction. If G is zero, the reaction is at equilibrium If G is
positive the reaction is nonspontaneous in the forward direction
but the reverse reaction is spontaneous.
Slide 55
Calculate the standard free energy change for the formation of
NO(g) from N 2 (g) and O 2 (g) at 298 K. Given that H o = 180.7 kJ
and S o = 24.7 J/K. Is the reaction spontaneous under these
circumstances?
Slide 56
Just like the standard enthalpies of formation there is a
tabulated list of standard free energies of formation. We use this
tabulated data in the same way we use the standard enthalpies of
formation.
Slide 57
Use the tabulated standard free energy of formation to
calculate. The standard free- energy change for the following
reaction at 298 K: P 4 (g) + 6 Cl 2 (g) 4 PCl 3
Slide 58
We saw that G o f values are tabulated. These all correspond to
the formation reactions at 298 K (25 o C). Many reactions do not
occur at 298 K.
Slide 59
HH SS-T S GGReaction Character
Slide 60
For a certain reaction, H o = -35.4 kJ and S o = -88.5 kJ/K. Is
this reaction exothermic or endothermic? Does the reaction lead to
an increase or increase in randomness or disorder of the system?
Calculate G o for this reaction at 298K. Is the reaction
spontaneous at 298K? Would the spontaneity of this reaction change
if we changed the temperature? How?
