Download Chapter 9: Energy and Chemistry

Chapter 9: Energy and Chemistry
•
•
•
•
•
•
Terminology
Conservation of Energy
Heat Capacity and Calorimetry
Enthalpy
Hess’s Law and Heats of Reaction
Energy and Stoichiometry
The rate of a reaction depends
on the pathway from reactants
to products; this is the domain
of kinetics. Thermodynamics
tells us whether a reaction is
spontaneous based only on the
properties of the reactants and
products. The predictions of
thermodynamics do not require
knowledge of the pathway
between reactants and
products.
1
Thermodynamics
– is concerned with energy transfer.
– predicts whether a reaction can occur or not.
– predicts how much heat is given out or taken in
during a reaction.
– it is not concerned with how fast a reaction will
occur – that is kinetics
Forms of Energy
• Two broad categories of energy: potential energy and kinetic energy.
– Potential energy - associated with the relative position of an object.
– Kinetic energy - associated with motion.
Kinetic Energy = ½mv²
• Internal energy - the combined kinetic and potential energies of atoms and
molecules that make up an object or system.
• Chemical energy - energy released or absorbed during a chemical reaction.
Chemical energy is a form of potential energy
• Other forms of energy include radiant, mechanical, thermal, electrical, and
nuclear.
• Thermochemistry - the study of the energetic consequences of chemistry
• The Joule is the SI unit of energy.
– 1 Joule = 1 kg m2/s2
2
Energy Transformation and Conservation of Energy
• During energy transformation, the total energy
must be conserved.
– The sum of all energy conversions and energy transfers
must equal the total energy present which must remain
constant.
• To account for energy transformations and
conversions, the system and surroundings must be
specified.
–
–
–
–
System - the part of the universe being considered.
Surroundings - the remainder of the universe.
System + Surroundings = Universe
System and surroundings are separated by a boundary.
Transfer of Energy
In the initial position, ball A has a higher potential energy than ball B.
After A has rolled down the hill, the potential energy lost by A has
been converted to random motions of the components of the hill
(frictional heading) and to the increase in the potential energy of B.
3
Heat
Energy is transferred between a system and its
surroundings as a result of a temperature
difference.
• Heat always flows from hotter to colder.
– Temperature may change.
– Phase may change (an isothermal process).
• This directionality corresponds to the spreading
out of energy over the greatest number of atoms,
molecules or ions.
Exothemeric and Endothermic
• Heat exchange accompanies chemical reactions.
– Exothermic: Heat flows out of the system (to the surroundings).
– Endothermic: Heat flows into the system (from the surroundings).
Exothermic
Endothermic
System is losing energy
System is gaining energy
4
The combustion of methane releases the quantity of energy change ∆(PE)
to the surroundings via heat flow. This is an exothermic process.
The energy diagram for the reaction of nitrogen and oxygen to form nitric
oxide. This is an endothermic process: Heat {equal in magnitude to
∆(PE)} flows into the system from the surroundings.
5
Work
• Work is the transfer of energy accomplished by a
force moving a mass some distance against
resistance.
– Pressure-volume work (PV-work) is the most common
work type in chemistry.
• Releasing an inflated balloon before it is tied off illustrates an
example of PV-work.
– Electrical work (IVt) is used in electrochemistry.
• Batteries, electrolysis
Pressure-Volume Work
• In addition to heat effects chemical reactions
may also do work.
• Gas formed pushes against
the atmosphere.
• Volume changes.
• Pressure-volume work.
6
Pressure-Volume Work
w=Fxd
= (P x A) x h
= PV
w = -PextV
Pressure-Volume Work Examples
Calculating Pressure-Volume Work.

Suppose the gas in the previous figure is 0.100 mol He at 298 K.
How much work, in Joules, is associated with its expansion at
constant pressure.

What is the maximum flow of energy in the form of work done
during an isothermal expansion?
7
Internal Energy
– Total energy (potential and kinetic) in a system.
•Translational kinetic energy.
•Molecular rotation.
•Bond vibration.
•Intermolecular attractions.
•Chemical bonds.
•Electrons.
• The total internal energy of a
sample depends on:
– Sample’s temperature
– The type of material in the sample
– The amount of material in the
sample
Energy Transformation
• A system contains only internal energy.
– A system does not contain heat or work.
– These only occur during a change in the system.
E = q + w
E = change in system’s internal energy
q = heat flow into system (positive value means
energy is flowing into the system)
w = work done on system (positive value means
energy is flowing into the system)
8
Work and Heat
State Functions
• Any property that has a unique value for a specified state of a
system is said to be a State Function (T, V, H, S, P, E, G).
• Depends only on the present state of the system - not how it
arrived there.
• It is independent of pathway
–
–
–
–
Water at 293.15 K and 1.00 atm is in a specified state.
d = 0.99820 g/mL
This density is a unique function of the state.
It does not matter how the state was established.
9
Functions of State
• Energy is a function of state.
– Not easily measured.
• The difference in energy has a unique value between
two states.
– Is easily measured.
Path Dependent Functions
• Changes in heat and work are not functions of
state.
– Remember pressure-volume work example,
w = -113 J in a one step expansion of gas:
– When expansion is done to get maximum work
w = -151 J in a reversibly expansion of gas:
10
Energy Transformation and Conservation of Energy
• First law of thermodynamics states that energy can
be transformed from one form to another but
cannot be created or destroyed.
Eunivierse = Esurroundings + Esystem = 0
– The energy of an isolated system is constant
Heat Capacity and Calorimetry
• Calorimetry is a laboratory method for observing
and measuring the flow of heat into and out of a
system.
• Different systems will absorb different amounts of
energy based on three main factors.
– The amount of material, m or n.
• m is mass and n is number of moles
– The type of material, as measured by c or Cp.
• c is the specific heat capacity, or specific heat, and Cp is the
molar heat capacity.
– The temperature change, T.
11
Heat Capacity and Specific Heat
• The specific heat capacity, or specific heat, is a
physical property of a substance that describes the
amount of heat required to raise the temperature of
one gram of a substance by 1ºC.
– Represented by c.
– Specific heat is compound and phase specific.
• The molar heat capacity is a physical property of
a substance that describes the amount of heat
required to raise the temperature of one mole of a
substance by 1ºC.
– Represented by Cp.
– Molar heat capacity is compound and phase specific.
Heat Capacity and Specific Heat
• The amount of heat energy absorbed can be
quantified.
– Specific heat capacity, c.
q = mcT
• System is one gram of substance.
– Molar heat capacity, Cp.
q = nCpT
• System is one mole of substance.
– Heat capacity
q = CT
• Mass dependent specific heat.
12
Heat Capacity and Specific Heat
• Specific heat and molar heat capacities for some common substances.
Calorimetry
• Heat flow is measured using a calorimeter.
• A calorimeter measures the heat evolved or
absorbed by the system of interest by
measuring the temperature change in the
surroundings.
qsystem - qsurroundings = 0
qsystem = -qsurroundings
qgained = -qlost
13
Determination of Specific Heat
100°C
Calculate the specific heat of lead.
Coffee Cup Calorimeter
• A simple calorimeter.
– Well insulated and therefore isolated.
– Measure temperature change.
• Reaction done at constant pressure
qrxn = -qcal
14
Bomb Calorimeter
Reaction done at constant volume
qrxn = -qcal
qcal = qbomb + qwater + qwires +…
Define the heat capacity of the
calorimeter:
qcal = miciT = CT
all i
Calorimetry
• There are two steps in a calorimetric
measurement.
– Calibration - the calorimeter constant, Ccalorimeter, is
determined by dividing the known amount of heat
released in the calorimeter by the temperature change of
the calorimeter.
Ccalorimeter = q/T
– Actual Measurement - heat released or absorbed in a
reaction of known quantity of material is measured.
qcal = CcalorimeterT
qrxn = -qcal
15
Heat of Reaction Example
•
Using Bomb Calorimetry Data to Determine a Heat of
Reaction.
•
The combustion of 1.010 g sucrose, in a bomb
calorimeter, causes the temperature to rise from 24.92 to
28.33°C. The heat capacity of the calorimeter assembly is
4.90 kJ/°C.
(a) What is the heat of combustion of sucrose, expressed in
kJ/mol C12H22O11
(b) Verify the claim of sugar producers that one teaspoon of
sugar (about 4.8 g) contains only 19 calories.
Reaction Conditions
• The conditions under which heat flow, q, occurs
will have an impact on the measurement that is
made.
– Combustion of octane releases 5.45 x 103 kJ under
constant volume conditions, represented as qv.
– Combustion of octane releases 5.48 x 103 kJ under
constant pressure conditions, represented as qp.
16
Heat Flow at Constant Volume
• The internal energy change for a reaction equals the sum of
the heat flow and the work.
E = q + w
– During an expansion, w = –PV.
E = q - PV
– Under constant volume conditions, V = 0, and E = qv.
• The change in energy is the heat flow under conditions of
constant volume.
Heat Flow at Constant Pressure
• Consider a process done under the condition of constant
pressure and with only PV-work.
E = qp + w = qp – PV
• Define Enthalpy, H, as H= E + PV
• The enthalpy change can be expressed as
H = E +(PV)
H = qp - PV + PV +VP
H = qp
• The change in enthalpy is the heat flow under conditions of
constant pressure.
17
Exothermic and Endothermic
• When a system releases heat, the process is said to
be exothermic.
– The value of H is less than zero; the sign on H is
negative.
– Enthalpy of products is less than enthalpy of reactants
• When a system absorbs heat, the process is said to
be endothermic.
– The value of H is greater than zero; the sign on H is
positive.
– Enthalpy of products is greater than enthalpy of
reactants
H of Phase Changes
• Phase changes occur under constant pressure conditions.
– The heat flow during a phase change is an enthalpy change.
– During a phase change, temperature does not change with heat flow due to
formation or breaking of intermolecular attractive forces.
18
∆H of Phase Changes
• The heat required to convert a liquid to a gas is the heat of
vaporization, Hvap.
– Hvap is endothermic with a positive value.
• The heat released to convert a gas to a liquid is the heat of
condensation, Hcond.
– Hcond is exothermic with a negative value.
• Hcond = –Hvap
– The values of enthalpy changes in opposite directions have equal
numeric values and differ only in their signs.
– The magnitude of enthalpy change depends on the substance
involved.
∆H of Phase Changes of Water
19
∆H of Phase Changes
• The value of H for a phase change is compound
specific and has units of kJ/mol.
– The heat flow can be calculated using the number of
moles of substance, n, and the value of the enthalpy
change.
H = nHphase change
Heating Curve
• The enthalpy change for the conversion of ice
to liquid and then to steam can be calculated.
– A heat curve breaks the calculation down into
specific heat calculations (sections of the heat
curve where temperature changes) and phase
change enthalpy calculations (sections of the heat
curve where temperature does not change).
20
Heating Curve for Water
• Heat curve for the heating of 500-g of ice at -50oC to 200oC.
Changes of State of Matter Example
Enthalpy Changes Accompanying Changes in States of Matter.
Calculate H for the process in which 50.0 g of water is
converted from liquid at 10.0°C to vapor at 25.0°C.
Break the problem into two steps: Raise the temperature of
the liquid first then completely vaporize it. The total enthalpy
change is the sum of the changes in each step.
21
Changes of State of Matter Example
A 25.0 g ice cube at 0.0°C is added to 100.0 g of liquid water at
22.0°C in a thermally isolated container. The heat of fusion of ice
is 6.01 kJ/mol and the specific heat of liquid water is 4.18 J g-1 °C-1.
(a) What is the final condition reached – liquid only, solid only, or
a mixture of solid or liquid?
(b) What is the final temperature?
Vaporization and Electricity Production
• Schematic
diagram of
the important
elements of a
standard
electric power
plant.
22
Heat of Reaction
• Enthalpy changes can be calculated for
chemical reactions, in addition to
temperature changes and phase transitions.
– The enthalpy change is commonly referred to
as the heat of reaction.
Standard States and Standard Enthalpy Changes
• Define a particular state as a standard state.
• Standard enthalpy of reaction, H°
– The enthalpy change of a reaction in which all reactants and
products are in their standard states.
• Standard State
– Element
• The form [N2(g), K(s)] in which it exists at 1 atm and 25°C.
– Compound
• For a gas, pressure is exactly 1 atmosphere.
• For a solution, concentration is exactly 1 molar.
• Pure substance (liquid or solid), it is the pure liquid or solid.
23
Bonds and Energy
• The enthalpy change for a reaction can be estimated using
bond energies.
• During a chemical reaction, reactant bonds are broken and
product bonds are made.
– Breaking bonds requires energy.
– Making bonds releases energy.
H 
D
Bonds Broken

D
Bonds Formed
• If the amount of energy released making product bonds is
greater than the amount of energy required to break
reactant bonds, the reaction is exothermic. If the energy
released is less than the energy required, the reaction is
endothermic.
Bonds and Energy
CH4 + 2O2  CO2 + 2H2O
• The combustion of methane breaks 4 C-H bonds and 2 O=O bonds.
2 C=O bonds and 4 O-H bonds are made.
24
Bonds and Energy
• The accuracy of enthalpy changes calculated from tabulated
bond energies is not very good.
– The bond energies used are averages.
– Bond energy method used to estimate enthalpy changes for reactions
involving compounds with no available thermochemical data.
• A thermochemical equation summarizes the overall energetics
for a chemical reaction.
– The sign on the H indicates whether the reaction is endothermic or
exothermic
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890.4 kJ
• The combustion of methane is an exothermic reaction and
releases 890.4 kJ of heat energy when 1 mole of methane reacts
with 2 moles of oxygen.
Heats of Reaction for Some Specific Reactions
• Some classes of chemical reactions are given their
own labels for heats of reactions.
– Heat of combustion, Hcomb
– Heat of neutralization, Hneut
– Heat of formation, Hf, is the heat of reaction for
formation of substances.
• Fractional coefficients are allowed for formation reactions
because only one mole of product can be formed.
C(s) + ½O2(g)  CO(g)
25