Transcript enthalpy change

Thermochemical Changes
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Thermochemistry
• The study of the energy
• Evolved or absorbed in chemical reactions
• Changes of physical transformations (Example: melting, boiling)
Thermochemistry
• Deals with the energy exchange
of transformations (chemical
reactions phase transitions,
nuclear reactions, dissolving
of ionic compounds)
• Includes calculations of
quantities such as the heat
capacity, heat of combustion,
heat of formation, enthalpy and
free energy.
• The atoms and molecules in any sample of matter are in constant
random motion.
• Depending on the phase or state of a sample of matter, the
atoms and molecules are capable of vibrating, rotating and
translating to varying extents.
• Heat is the name given to the
energy transferred between two
substances that have different
temperatures.
Temperature Measurement
• Celsius and Kelvin temperatures scales have the same size units
(Example:1 oC change is equal to 1 K change)
Temperature Conventions
• ∆t is used for Celsius measurements
• ∆T is used for Kelvin measurements
• Increase in temperature - positive
• Decrease in temperature - negative
Primary Sources of Energy
• Chemical - fossil fuels, plants
• Solar - direct radiant, wind,
water
• Nuclear - fission, fusion
• Geothermal - geysers, hot springs
Useful forms of Energy
• Heat
• Electrical
• Mechanical
• Light
• Sound
Heat as Energy
• Almost all forms of energy are eventually converted to thermal
energy
• If heat is so spread out that it
cannot be used it is called
low level heat
An Introduction to Energetics
 Kinetic Energy (Ek) is related to the motion of an entity
 Molecular motion can by translational (straight-line),
rotational and vibrational
 Chemical Potential Energy (Ep) is energy stored in the
bonds of a substance and relative intermolecular forces
 Thermal Energy is the total kinetic energy of all of the
particles of a system. Increases with temperature.
 Symbol (Q), Units (J), Formula used (Q=mcΔT)
 Temperature is a measure of the average kinetic energy of
the particles in a system
 Heat is a transfer of thermal energy. Heat is not possessed by
a system. Heat is energy flowing between systems.
DO YOU REMEMBER??
THE LAW OF CONSERVATION OF ENERGY
• During physical and chemical
processes, energy may change form,
but it may never be created nor
destroyed.
• If a chemical system gains energy, the
surroundings lose energy
• If a chemical system loses energy, the
surroundings gain energy
Examples:
• When octane (C3H8, the main
component of gasoline) is burned in
your car engine, chemical bond
energy (potential energy) is
converted into mechanical energy
(pistons moving in the car engine;
kinetic energy) and heat.
• When we turn on a light switch,
electrical energy is converted into
light energy and, you guessed it,
heat energy.
DO YOU REMEMBER??
EXOTHERMIC
• A change in a chemical energy
where energy/heat EXITS the
chemical system
• Results in a decrease in chemical
potential energy
ENDOTHERMIC
• A change in chemical energy
where energy/heat ENTERS the
chemical system
• Results in an increase in
chemical potential energy
Heat as Energy
• Exothermic change - reaction releases E (surroundings increase in temp.)
• Endothermic change - reaction absorbs
E (surroundings decrease in temp.)
THE LAW OF CONSERVATION OF ENERGY
Thermal Energy Calculations
 There are three factors that affect thermal energy (Q = mcΔt):
 Mass (m)
 Type of substance (c)
 c is the specific heat capacity, the quantity of energy required to raise the temperature of
one gram of a substance by one degree Celsius
 Temperature change (Δt)
 Example: Consider a bathtub and a teacup of water! All water has the same
specific heat capacity which is 4.19J/g°C. However, the bathtub would take
considerably more energy to heat up!
Thermal Energy Calculations
 Example: Determine the change in thermal energy when 115 mL of
water is heated from 19.6oC to 98.8oC?
MASS = DENSITY X VOLUME
SHOW HOW L = kg AND mL = g
The density of a dilute aqueous solution is the same as that of water;
that is, 1.00g/mL or 1.00kg/L
c water = 4.19J/g °C
or
4.19 kJ/kg °C
or 4.19 kJ/L °C
Thermal Energy Calculations
 Example: A sample of ethanol absorbs 23.4 kJ of energy when
its temperature increases by 14.25°C . The specific heat
capacity of ethanol is 2.44 J/g°C . What is the mass of the
ethanol sample?
Q = 23.4 kJ
Q = mcΔt
Δt = +14.25°C
m = Q/cΔt
c = 2.44 J/g°C
m=
23.4 kJ
.
(2.44J/g°C)(+14.25°C)
m=?
m = 0.6729 kg = 0.673 kg
Example
• How much heat energy is absorbed when 10.0kg of
aluminum initially at a temperature of 20 degrees
Celsius, warms up to its melting point , 660 degrees
Celsius, but does not melt?
Example
• A 2.50kg sample of water has a temperature of 85
degrees Celsius. What is the final temperature of
the water if it loses 220 kJ of heat energy?
Example
• What mass of water, at 25 degrees Celsius, requires 2.25 MJ of
heat energy to warm up 100 degrees Celsius? (Assume the water
being heated is below sea level in order to allow for heating to 100
degrees Celsius without changing states.)
Example: Day 2
• A 55.0g block of aluminum cools from 53.2°C to
25.0°C. What amount of energy is lost by the
aluminum block?
Endothermic/Exothermic Reactions
• Recall that energy is required to break bonds & is released when
bonds are formed.
Endothermic/Exothermic Reactions
• A net absorption of energy results in an endothermic reactions while a net
release of energy gives an exothermic reaction.
Energy of Chemical Systems
• Kinetic Energy
• Moving electron w/in atoms
• Vibration of atoms connected by chemical bonds
• Rotation & translation of molecules
• Potential Energy
• Covalent and /or ionic bonds between atoms
• Intermolecular forces
How do we measure Q?
• With a simple laboratory calorimeter, which
consists of an insulated container made of three
nested polystyrene cups, a measured quantity of
water, and a thermometer.
• The chemical is placed in or dissolved in the water
of the calorimeter.
• Energy transfers between the chemical system and
the surrounding water is monitored by measuring
changes in the water temperature.
• “Calorimetry is the technological process of
measuring energy changes of an isolated system
called a calorimeter”
Includes: Thermometer, stirring rod,
stopper or inverted cup, two Styrofoam
cups nested together containing
reactants in solution
Comparing Q’s
Negative Q value
• An exothermic change
• Heat is lost by the system
• The temperature of the
surroundings increases and the
temperature of the system
decreases
Positive Q value
• An endothermic change
• Heat is gained by the system
• The temperature of the system
increases and the temperature of
the surroundings decreases
• Example: Hot Pack
• Example: Cold Pack
• Question Tips: “How much
energy is released?”
• Question Tips: “What heat is
required?”
ENTHALPY
 The total of the kinetic and potential energy within a
chemical system is called its enthalpy. (Energy
possessed by the system)
 Enthalpy is communicated as a difference in
enthalpy between reactants and products, an
enthalpy change, ΔrH
.
 Units (usually kJ)
Enthalpy Change (ΔH)
• It is not possible to determine the absolute enthalpy of a system (H) but
changes in enthalpy can be measured (∆H)
• ΔH is a difference in enthalpy between reactants and products of a particular
chemical system when the pressure or volume of a system is held constant
•ΔH = HP – HR
(IUPAC defn.)
• Enthalpy of reaction is measured in kJ
Enthalpy Change
• Enthalpy change can be communicated in the following four
ways
• The molar enthalpy of a process, such as formation or
combustion, (units kJ/mol)
• The enthalpy change of a reaction (units kJ)
• The enthalpy of a reaction written as a reactant or product term
in the equation (units kJ)
• An energy diagram representing a reaction (units kJ)
Enthalpy Change
• First it is important to be familiar with some common definitions
• The standard molar enthalpy change of formation of a compound,
is
the heat energy absorbed or released when 1 mol of compound is formed
from its elements in their standard states.
• Values of
can be found on the table of standard molar enthalpies
of formation. The variable
can also be represented as
• The standard molar enthalpy of formation of elements in their standard
states is defined as 0 kJ/mol.
+
Signs
-
• Negative enthalpy values indicate exothermic reactions
(energy is released to the surroundings thus is lost by the
system)
• Positive enthalpy values indicate endothermic reactions
(energy is gained by the system from its surroundings)
Types of Enthalpy
• As noted earlier enthalpy changes occur during other
processes (eg. phase changes, nuclear changes,
dissolving of ionic compunds) so your text uses the
symbol ∆rH to indicate the enthalpy of reaction
IUPAC Symbols
• If the reaction is:
• combustion the symbol used can be ∆c H
• formation the symbol used can be ∆f H
• decomposition the symbol used can be ∆d H
• dissolving to form a solution the symbol used can be ∆sol H
• dilution of a solution the symbol used can be ∆dil H
Standard Enthalpy of Reaction (∆ r H o)
• Since enthalpy changes differ depending upon pressure and
temperature standard enthalpy of reaction values are for
SATP (1 mol/L, 25 oC & 100 kPa)
• This means that the initial and the final conditions of the
reaction system must be in a standard state.
Communicating Enthalpy Changes of Chemical
Reactions
- Enthalpy changes can be communicated by four methods:
1.
2.
3.
4.
Stating the molar enthalpy of a reaction
Thermochemical equations
ΔH notation
Potential energy diagrams
ENTHALPY CHANGES
 In a simple calorimetry experiment involving a burning candle and a can of water, the temperature
of 100 mL of water increases from 16.4°C to 25.2°C when the candle is burned for several
minutes. What is the enthalpy change of this combustion reaction?
Assuming: ΔcH = Q (The energy lost by the chemical system, (burning candle), is equal to the
energy gained by the surroundings (calorimeter water)
Assuming: Q = mcΔt then ΔcH = mcΔt
We will use
ΔcH = - mcΔt
• Is the value of ΔcH going to be positive or negative??
• If the surroundings (water) gained energy, then the system (burning candle) lost it.
So based on the evidence, the enthalpy change of combustion for this reaction is 3.69J
ENTHALPY CHANGES
 When 50 mL of 1.0 mol/L hydrochloric acid is neutralized completely by 75 mL of 1.0
mol/L sodium hydroxide in a polystyrene cup calorimeter, the temperature of the
total solution changes from 20.2°C to 25.6°C. Determine the enthalpy change that
occurs in the chemical system.
Is this an Endothermic or
Exothermic reaction??
 Based upon the evidence available, the enthalpy change for the neutralization of
hydrochloric acid in this context is recorded as -2.83 kJ.
DO YOU REMEMBER??
EXOTHERMIC
• A change in a chemical energy where
energy/heat EXITS the chemical system
• Results in a decrease in chemical potential
energy
• ΔH is negative
ENDOTHERMIC
• A change in chemical energy where
energy/heat ENTERS the chemical system
• Results in an increase in chemical potential
energy
• ΔH is positive
MOLAR ENTHALPY
• Molar enthalpy: ΔrHm the change in enthalpy expressed per mole of a substance
undergoing a specified reaction (kJ/mol)
• Have we had other quantities expressed per mole? YES!
• How will we calculate this?
• Molar Enthalpy: change in enthalpy per mol of a specified
chemical undergoing a change.
• Important because it lets us do calculations using chemical
reactions
• Enthalpy change and Molar enthalpy use the same symbol
ΔH, so pay attention to the context and units of the
question.
• Enthalpy = kJ
• Molar Enthalpy= kJ/mol
• IUPAC Symbol = ΔrHm
ΔH Notation
• Describing the enthalpy change (ΔH) for a balanced
equation
• To calculate an enthalpy change a molar enthalpy value and
a
balanced equation is required.
Example
• Standard molar enthalpy of formation of CuO (s) is
Hof (CuO) = - 157.3 kJ/mol
(Data Booklet)
Example: 157.3 kJ is released to the surroundings during the
formation of CuO (s) (exothermic) under SATP conditions
Molar enthalpy of a reaction
• Note: if the balanced chemical reaction indicates that the specified
compound has a coefficient other than 1 the molar enthalpy must
be multiplied by the number of moles to obtain the total enthalpy
of reaction.
• When determining a molar enthalpy change, it is critical to specify
which substance a molar enthalpy is for.
Try This
• Formation of water
MOLAR ENTHALPY
1. Predict the change in enthalpy due to the combustion of 10.0 g of propane used in a
camp stove. The molar enthalpy of combustion of propane is -2043.9 kJ/mol.
2. Predict the enthalpy change due to the combustion of 10.0 g of butane in a camp
heater. The molar enthalpy of combustion of butane is -2657.3 kJ/mol.
Reaction for the combustion of propanone:
CH3COCH3 (l) + 4O2(g)  3CO2(g) + 3H2O(g)
Molar Enthalpy of propanone combustion:
HCH3COCH3(l) = -1 659.1 kJ/mol
What is the molar enthalpy of the combustion of
propanone per mole of carbon dioxide formed?
MOLAR ENTHALPY AND CALORIMETRY
• Can we measure the molar enthalpy of reaction using calorimetry?
• Yes, but indirectly. We can measure a change in temperature, we can then
calculate the change in thermal energy (Q=mct). Then, using the law of
conservation of energy we can infer the molar enthalpy.
• In doing so, we must assume that the change in enthalpy of the chemicals involved
in a reaction is equal to the change in thermal energy of the surroundings.
From this equation, any one of
the five variables can be
determined as an unknown.
MOLAR ENTHALPY
1. In a research laboratory, the combustion of 3.50 g of ethanol in a sophisticated
calorimeter causes the temperature of 3.63 L of water to increase from 19.88°C to
26.18°C. Use this evidence to determine the molar enthalpy of combustion of
ethanol.
** You don’t have to equate the two formulas to solve this. Instead, you can calculate Q, then use that value as
ΔrH, and solve for either the chemical amount or the molar enthalpy of reaction.
Q = 95.8 kJ = ΔH
ΔcHm = 1.26 x 103 kJ/mol
= 1.26 MJ/mol
Heat Capacity
Q = CΔt
Thermochemical Equations Example
Combustion of propane
Thermochemical Equations Example
• The formation of dinitrogen tetraoxide
Decomposition Reactions
• The values in the data booklet apply for decomposition
reactions if you
reverse the enthalpy sign.
Example: decomposition of water
H2O(l) –> H2(g) + ½ O2(g)
ΔH = + 285.8 kJ
Thermochemical Equations Example
• If the equation 2NaOH + H2SO4  Na2SO4 + 2H2O +57.0 kJ was rewritten such
that the energy term had a value of 171.0 kJ the respective balancing
coefficients for the equation would be what?
Thermochemical Equations Example
• If the equation 2Ag + I2  2AgI + 123.6 kJ was rewritten to produce 1 mol of
AgI, the magnitude of the energy term would be
Try This
• If the combustion of 1.28 mol of benzene
(C6H6 (l)), produces an enthalpy change of - 8027 kJ, what
is the enthalpy change for this combustion as shown for a
balanced chemical equation?
2 C6H6(l) + 15 O2(g) –> 12 CO2(g) + 6 H2O(g)
Step 1 Find the molar enthalpy of combustion
of benzene
Step 2 Find the enthalpy change for 2 moles of
benzene
Example
• How much energy would be released by the combustion of propane if 30.0g of
carbon dioxide was formed?
Example
• How much energy is evolved from the combustion of 25.0g of ethane if
ΔcH° = -1428.4 kJ/mol?
Example
• Commercial drain cleaner typically contain sodium hydroxide and aluminum.
When the solid cleaner is poured down the drain water is added, a reaction
occurs. The given equation represents this reaction
• If the given reaction produces 6.75 moles of NaAlO2, then the heat released is
what?
Thermochemical equations
• Chemical equations which show an energy change as
part of the balanced chemical equation
Example: Combustion of benzene
2 C6H6(l) + 15 O2(g) –> 12 CO2(g) + 6 H2O(g)
As a thermochemical equation becomes:
2 C6H6(l) + 15 O2(g) –> 12 CO2(g) + 6 H2O(g) + 12.5 MJ
Note:
Endothermic energy goes on reactant side (ΔH = + )
Exothermic energy goes on product side (ΔH = - )
ΔcH = - 12.5 MJ
COMMUNICATING ENTHALPY
• We will be learning how to communicate enthalpy changes in four ways:
1. By stating the molar enthalpy of a specific reactant in a reaction
2. By stating the enthalpy change for a balanced reaction equation
3. By including an energy value as a term in a balanced reaction equation
4. By drawing a chemical potential energy diagram
COMMUNICATING ENTHALPY #1
1. By stating the molar enthalpy of a specific reactant in a reaction
• Why do we use standard conditions in chemistry (i.e. SATP)?
We use a standard set of conditions so that scientists can
create tables of precise, standard values and can compare
other values easily
• Do we have standard conditions for enthalpy??
Yes, we will be using SATP (but liquid and solid compounds
must only have the same initial and final temperature – most
often 25°C)
• How do we communicate that standard conditions are used for
reactants and products?
•
With a ° superscript, such as ΔfHm° or ΔcHm° (See data
booklet pg. 4 and 5)
•
*For well-known reactions such as formation and
combustion, no chemical equation is necessary, since they
refer to specific reactions with the Δf or Δm
•
** Would the sign for ΔfHm° be the opposite of the sign
for ΔdHm° (decomposition)? YES!
•
*For equations that are not well known or obvious, then
the chemical equation must be stated along with the molar
enthalpy.
COMMUNICATING ENTHALPY #1
1. By stating the molar enthalpy of a specific reactant in a reaction
Example #1:
This means that the complete combustion of 1 mol of
methanol releases 725.9 kJ of energy according to the
following balanced equation
•
Example #2:
This does not specify a reaction, so a chemical equation must be stated along with
the molar enthalpy.
• This is not a formation reaction, since not all of the reactants are elements, so this
could not have been communicated with Δf
•
COMMUNICATING ENTHALPY #2
2. By stating the enthalpy change beside a balanced reaction
equation
Do we know how to calculate enthalpy change??
• The enthalpy change for a reaction can be determined by multiplying
the chemical amount (from the coefficient in the equation) by the
molar enthalpy of reaction (for a specific chemical)
• Example: Sulfur dioxide and oxygen react to form sulfur trioxide. The
standard molar enthalpy of combustion of sulfur dioxide, in this
reaction, is -98.9 kJ/mol. What is the enthalpy change for this
reaction?
1)Start with a balanced chemical equation.
2)Then determine the chemical amount of SO2 from the equation =
2 mol
(this is an exact #, don’t use for sig digs)
3) Then use
whole reaction.
to determine the enthalpy change for the
4) Then report the enthalpy change by writing it next to the balanced
equation.
COMMUNICATING ENTHALPY #2
2. By stating the enthalpy change beside a balanced reaction
equation
3. THE ENTHALPY CHANGE DEPENDS ON THE ACTUAL CHEMICAL AMOUNT OF REACTANTS
AND PRODUCTS IN THE CHEMICAL REACTION. THEREFORE, IF THE BALANCED EQUATION
IS WRITTEN DIFFERENTLY, THE ENTHALPY CHANGE SHOULD BE REPORTED DIFFERENTLY
Both chemical reactions agree with the
empirically determined molar enthalpy of
combustion for sulfur dioxide
COMMUNICATING ENTHALPY #2
2. By stating the enthalpy change beside a balanced reaction
equation
•
THE ENTHALPY CHANGE DEPENDS ON THE ACTUAL CHEMICAL AMOUNT OF
REACTANTS AND PRODUCTS IN THE CHEMICAL REACTION. THEREFORE, IF
THE BALANCED EQUATION IS WRITTEN DIFFERENTLY, THE ENTHALPY
CHANGE SHOULD BE REPORTED DIFFERENTLY
Example 2:
2Al(s) + 3Cl2(g)  2AlCl3(s)
ΔfH° = -1408.0 kJ
• What is the molar enthalpy of formation of aluminum chloride?
ΔfHm° = -1408.0kJ = -704.0 kJ/mol AlCl3
2 mol
COMMUNICATING ENTHALPY #2
2. By stating the enthalpy change beside a balanced reaction
equation
•
EXAMPLE: The
standard molar enthalpy of combustion of hydrogen
sulfide is -518.0 kJ/mol. Express this value as a standard enthalpy
change for the following reaction equation:
•
SOLUTION:
COMMUNICATING ENTHALPY #3
3. By including an energy value as a term in a balanced reaction
equation
•
If a reaction is endothermic, it requires additional energy to react, so is listed along with the
reactants
•
If a reaction is exothermic, energy is released as the reaction proceeds, and is listed along with
the products
•
In order to specify the initial and final conditions for measuring the enthalpy change of the
reaction, the temperature and pressure may be specified at the end of the equation
COMMUNICATING ENTHALPY #3
3. By including an energy value as a term in a balanced reaction equation
• EXAMPLE: Ethane is cracked into ethene in world-scale quantities
in Alberta. Communicate the enthalpy of reaction as a term in the
equation representing the cracking reaction.
DOES THE +136.4 kJ MEAN
EXOTHERMIC OR ENDOTHERMIC?
COMMUNICATING ENTHALPY #3
3. By including an energy value as a term in a balanced reaction
equation
•
EXAMPLE: Write the thermochemical equation for the formation of 2 moles of
methanol from its elements if the molar enthalpy of formation is -108.6kJ/mol
2 C(s) + 4 H2(g) + O2(g)  2 CH3OH(l) + ___?_____
ΔfH = 2 mol (-108.6 kJ/mol) = -217.2 kJ (Exothermic)
2 C(s) + 4 H2(g) + O2(g)  2 CH3OH(l) + 217.2 kJ
COMMUNICATING ENTHALPY #4
4. By drawing a chemical potential energy diagram
•
During a chemical reaction, observed energy changes are due to changes in
chemical potential energy that occur during a reaction. This energy is a stored form
of energy that is related to the relative positions of particles and the strengths of the
bonds between them.
•
As bonds break and re-form and the positions of atoms are altered, changes in
potential energy occur. Evidence of a change in enthalpy of a chemical system is
provided by a temperature change of the surroundings.
•
A chemical potential energy diagram shows the potential energy of both the
reactants and products of a chemical reaction. The difference is the enthalpy
change (obtained from calorimetry)
•
Guidelines: The vertical axis represents Ep. The reactants are written on the left,
products on the right, and the horizontal axis is called the reaction coordinate or
reaction progress.
Chemical Potential Energy Diagrams
• Show the Ep of the reactants before the reaction and
the products after the reaction
• As in chemical equations, the reactants are written on
the left and the products go on the right
Exothermic Reactions
• Chemical potential energy diagrams can illustrate enthalpy changes
Exothermic Reactions
• An exothermic reaction releases heat to its surroundings thus losing
energy
• Temperature of surroundings increases
Endothermic Reactions
• An endothermic reaction gains heat from its surroundings thus the
temp. of the surroundings decreases
COMMUNICATING ENTHALPY
During an exothermic reaction, the enthalpy of During an endothermic reaction, heat flows
the system decreases and heat flows into the
from the surroundings into the chemical
surroundings. We observe a temperature
system. We observe a temperature
increase in the surroundings.
decrease in the surroundings.
COMMUNICATING ENTHALPY #4
Summary
• In an exothermic reaction the products have less Ep than the
reactants & energy is released to the surroundings as the
products form
• In an endothermic reaction the products have more Ep than the
reactants & E is absorbed from the surroundings
COMMUNICATING ENTHALPY #4
•
EXAMPLE: Communicate the following enthalpies of reaction as a chemical
potential energy diagram.
• The burning of magnesium to produce a very bright emergency flare.
•
The decomposition of water by electrical energy from a solar cell.
Your Task
• Practice Questions #9-13 page 493
• Practice Questions #3,4,5 page 494
Due at the start of class tomorrow
Energy Diagrams
Let’s Review Quickly…
• Activation Energy (Ea) - The energy level that the reactant molecules
must overcome before a reaction can occur
Endothermic Reaction
More to come later
In an endothermic reaction in which
the change in enthalpy between the
products and the reactants is positive,
there must be an extra input of energy
above the energy level of the products in
order for a reaction to occur.
Exothermic Reaction
Even in an exothermic reaction in which
the change in enthalpy between the
products and the reactants is negative,
there must be an input of energy to start
the reaction.
• Enthalpy (H) - The sum of the internal energy of the system plus the product of
the pressure of the gas in the system and its volume:
Exothermic - Reaction in which a system RELEASES heat to its
surroundings. H is negative ( H < 0)
Ea is the activation energy
Endothermic - Reaction in which a system ABSORBS
heat from its surroundings.
H is positive ( H > 0)
Review Example
• Write the equation representing the reaction for the formation
for magnesium carbonate, and determine the standard molar
enthalpy of formation for magnesium carbonate
Review Example Continued…
• Determine the reaction enthalpy change for the equation
representing the formation of magnesium carbonate balanced
using whole number coefficients. Include the value of enthalpy
change as an energy term in the equation.
Review Example Continued…
• Draw the potential energy diagram that represents the
formation of magnesium carbonate
Review Example 2
• The standard molar enthalpy of neutralization for sodium
hydroxide with any monoprotic strong acid is -56.8 kJ/mol.
What quantity of heat energy would be released if 750 mL of
1.56 mol/L sodium hydroxide were neutralized with an
appropriate quantity of strong acid?
Review Example 3
• Heptane burns in a camping stove. What quantity of energy
is released for every 1.00 kg of carbon dioxide produced?
Review Example 4
• Heptane burns in a camping stove. What quantity of energy is
released for every 1.00 kg of heptane is produced?
Predicting Enthalpy Changes
• The enthalpy change for any reaction is the difference between the sum of the
formation enthalpies of the products and the sum of the formation enthalpies
of the reactants.
• The following equation is used for predicting enthalpy change of a reaction
Example
• Determine the enthalpy change for the complete combustion
of 1 mol of propane gas
Example
• Determine the enthalpy change for the reaction of ethane with water
Example
• Determine the enthalpy change when 1 mol of ammonium nitrate decomposes
to produce water vapour and dinitrogen monoxide during an explosion
Example
• Butane is added to gasoline to make it perform better under rigorous
conditions of the Canadian winter. Approximately how much energy is
produced when one mole of butane burns in an automobile engine to give
gaseous products?
For Your Viewing Pleasure…
• https://www.youtube.com/watch?v=Nj6euCKpa6U
• https://www.youtube.com/watch?v=sJob0_V9ers
• https://www.youtube.com/watch?v=4jQmc07vss0
• https://www.youtube.com/watch?v=_UcGFQpYeMc
• https://www.youtube.com/watch?v=uyIzU1fPUNI
HESS’ LAW
•
Do you think it is convenient or possible to use a calorimeter to test all chemical reactions?
•
NO! Sometimes two products are created simultaneously, sometimes a reaction is too small to be
able to measure accurately. So what do scientists do?
•
Theoretically, we assume that the enthalpy change of a physical
or chemical process depends only on the initial and final conditions.
It is independent of the pathway, process or number of
intermediate steps required.
•
Illustration: Bricks are being moved from the ground up
to the second floor. But there are two pathways to do this:
•
Move from the 1st to 2nd floor
•
Move to third floor and then carry down one flight
•
In both cases, the overall change in position is the same.
Germain Henri Hess
Is important primarily for
his thermochemical studies
●
(1802 - 1850)
Hess’ Law of Constant Heat Summation
aka Law of Additivity of Enthalpies of Reaction
G. W. Hess suggested in 1840 that:
- The change in enthalpy for any reaction depends only on the nature of the reactants
and products and is independent of the number of steps or the pathway taken
between them.
Hess’s Law can be written as an equation:
•
The uppercase Greek Letter, Σ (sigma) means
“the sum of”
Hess’ Law
• Hess’ discovery allows us to determine enthalpy change without
direct calorimetry, using two rules that you already know:
1) If a chemical equation is reversed, then the sign of ΔrH changes
2) If the coefficients of a chemical equation are altered by
multiplying or dividing by a constant factor, then the ΔrH is
altered by the same factor
Hess’ Law
Hess's Law
The enthalpy change for any reaction depends only
on the energy states of the initial reactants and final
products and is independent of the pathway or the
number of steps between the reactant and product.
Predicting Enthalpy Changes
Whether a product is formed from a one step
reaction or a series of reactions the enthalpy
change will be the same provided the initial
reactants and final products are the same and
the same initial and final conditions are present.
Using Hess’s Law
•
Write a balanced net chemical reaction
•
Manipulate given formation equations to yield the net equation (x,
÷, and/or reverse ∆rH)
•
Cancel and add equations to yield the net equations
•
Add component enthalpy changes to obtain the net enthalpy
change
•
Determine molar enthalpy if required
Hess’s Law Example - Combustion of Methane
ΔrH = - 802.5 kJ
CH4(g) + 2 O2(g)
CO2 (g) + 2 H2O (g)
Reactants
Products
Alternate Path
CH4(g) + 2 O2(g)
ΔHdecomp = + 74.6 kJ
ΔHr = - 802.5 kJ
Hess said sum
these two energies
CO2(g) + 2 H2O(g)
ΔHformation = - 393.5 kJ
+ (- 483.6 kJ)
Reactants
break into
elements
C(s) + 2 H2(g) + 2 O2(g)
Products form
from these
elements
Hess’ Law #1
•
Example: Use Hess’ Law to determine the enthalpy change for the formation of carbon monoxide.
•
This reaction can not be studied calorimetrically but we are given the following information to help solve
this equation
•
Our job now, is to manipulate the equations so they will add to yield the net equation
•
We need 1 mol of C(s) to start the equation, so leave (1) unaltered
•
However, we want 1 mol of CO as a product, so reverse equation (2) and divide all terms by 2
•
** Remember whatever you do to the equation, affects the ΔH the same way
ΔcH = -566.0kJ (original equation)
1) Reversed equation; ΔH = + 566.0kJ
2) Divide equation by 2; Divide ΔH by 2 = +283.0kJ
Hess’ Law #1
•
Example: Use Hess’ Law to determine the enthalpy change for the formation of
carbon monoxide.
Now cancel and add the remaining reactants and products to yield the net
equation.
• Add the component enthalpy changes to obtain the net enthalpy change.
•
The process of using Hess’ Law is a
combination of being systematic and using
trial and error. Do what needs to be done to
the given equations so they add to get the
equation you want.
Hess’ Law #1
•
Example: Use Hess’ Law to determine the enthalpy change for the formation of
carbon monoxide.
•
Sketching a potential energy diagram might help you ensure that you have made
the appropriate additions and subtractions
Example Calculations
1) Nitrogen and oxygen gas combine to form nitrogen dioxide according to
the following reaction:
Calculate the change in enthalpy for the above overall reaction, given:
From the following enthalpy changes:
Calculate the value of
H
for the reaction:
Example
• The formation of liquid hexane is represented by the following overall net
equation: 6C(s) + 7H2(g)  C6H14(l)
C6H14(l) + 19/2 O2(g)  6CO2(g) + 7H2O(l)
C(s) + O2(g)  CO2(g)
H2(g) + ½ O2(g)  H2O(l)
Using Hess’ Law, determine the molar enthalpy of the formation for hexane
Hess’ Law #2
•
Example: One of the methods the steel industry uses to obtain metallic iron is to
react iron(III) oxide with carbon monoxide
Fe2O3(s) + 3CO(g)  3CO2(g) + 2Fe(s)
1) CO(g) + ½ O2(g)  CO2(g)
ΔfH = -283.0 kJ
2) 2Fe(s) + 3/2O2(g)  Fe2O3(s)
ΔfH = -822.3 kJ
3( CO(g) + ½ O2(g)  CO2(g))
reverse
ΔrH = ??
ΔfH =3(-283.0 kJ) = -849.0 kJ
Fe2O3(s)  2Fe(s) + 3/2O2(g)
Fe2O3(s) + 3CO(g)  3CO2(g) + 2Fe(s)
ΔfH = -822.3 kJ
= +822.3 kJ
ΔrH = -26.7 kJ
Hess’ Law #3
•
Example: What is the standard enthalpy of formation of butane? ΔfHm° = ???
•
First, we need to be able to write this balanced formation equation.
4C(s) + 5H2(g)  C4H10(g)
•
The following values were determined by calorimetry:
•
What will we need to do to get our net equation?
-Reverse equation (1) and change the
ΔH sign
-Multiply equation (2) and its ΔH by 4
-Multiply equation (3) and its ΔH by
5/2
ΔfHm° = -125.6 kJ/1 mol = -125.6 kJ/mol
C4H10
Hess’ Law
• If you can add equations then you can add ΔH’s
ΔHnet = Σ ΔrH
Note: Σ - means “sum of”
Recall these rules of enthalpy change:
1. If a chemical equation is reversed then the sign for ΔHr changes
2. If the coefficients are changed by a constant factor then ΔHr changes in the
same way
Why does Hess’s Law work?
• There are only a limited number or formation reactions (one for each
compound)
• But there are a seemingly infinite number of pathways for each
reaction….
• Therefore, if we always go through the elements pathway, it will limit
the amount of information required to calculate the theoretical
energy of reaction
Enthalpy change for a reaction can be obtained by:
ΔrHo
↓
--------------------------------------------------
↓
↓
Calorimetry
Hess’s Law
(Experimental value)
nH = mcΔt
(Predicted value)
- ΔrHo = Σn∆fpHo - Σn∆frHo
- manipulating intermediate
steps
Hess’ Law Example
20kJ of heat is involved when 4.2 mol of Y is used in the exothermic
reaction above. Which of the following reactions is correct?
2X(s) + 3Y(g)  X2Y3(g)
H = +20 kJ
2X(s) + 3Y(g)  X2Y3(g)
H = -20 kJ
2X(s) + 3Y(g) + 14kJ  X2Y3(g)
2X(s) + 3Y(g)  X2Y3(g) + 14kJ
Example
• Determine the enthalpy change for the hydrogenation reaction of benzene to
make cyclohexane.
Molar Enthalpy of Formation
• Molar enthalpies of formation are defined as the enthalpy change
when one mole of a compound forms from its elements
• NOTE: The enthalpy of formation for an element is 0 kJ
• Examples: Using your data booklet, find the following:
• Δf Hm ° CH4(g) = -74.6 kJ/mol
• Δf Hm ° O2(g) = 0 kJ/mol (Δf H elements = 0)
• Δf Hm ° CO2(g) = - 393.5 kJ/mol
• Δf Hm ° H2O(g) = - 241.8kJ/mol
Reference energy state
• A convention which describes elements as the reference point at
which potential energy is zero
• Allows chemists to compare enthalpy changes to the reference
state of zero and tabulate these values
Thermal stability
• Tendency of a compound to resist decomposition when
heated.
Thermal Stability Example
• Example: List the following compounds in decreasing order of thermal
stability;
A.
B.
C.
D.
carbon dioxide,
ethyne,
magnesium chloride,
hydrogen bromide
• Answer (C, A, D, B )
Standard Molar Enthalpies of Formation
Hess’s Law can be stated mathematically as:
ΔrHo = Σ(n∆fHo products) - Σ(n∆fHo reactants)
[Enthalpy change = the sum of the molar enthalpy of formation of products
minus the sum of the molar enthalpy of formation of reactants]
[ Note: When using this equation do NOT change any signs of values, the equation
accounts for whether the enthalpy values are for products or reactants. ]
MOLAR ENTHALPY OF FORMATION
•
Why do we care about the standard molar enthalpies of formation, ΔfH° ???
•
Because we are going to use them to predict standard enthalpy changes for
chemical reactions. How? Using this crazy formula!!
•
What does it mean? The net enthalpy change for a chemical reaction, ΔrH°, is
equal to the sum of the chemical amounts times the molar enthalpies of
formation of the products, ΣnΔf pHm °, minus the chemical amounts times the
molar enthalpies of formation of the reactants, ΣnΔf RHm °
•
Clear as mud?? Basically, the equation says that the change in enthalpy is the
total chemical potential energy of the products minus the reactants. Epproducts –
Epreactants
•
We will need to use an example to figure this out.
Molar Enthalpy of Formation
• Calculate the molar enthalpy of formation for two moles of carbon
monoxide from its elements.
2C(s) + O2(g)  2CO(g)
ΔfHm = 2 mol(-110.5 kJ) - 2 mol(0 kJ) + 1 mol(0 kJ)
mol
mol
mol
= -221.o kJ
2 mol
= -110.5 kJ/mol
MOLAR ENTHALPY OF FORMATION
•
Methane is burned in furnaces and in some power plants. What is the standard
molar enthalpy of combustion of methane? Assume that water vapor is a
product.
•
Need a balanced chemical equation: CH4(g) + O2(g)  CO2(g) + 2H2O(g)
•
Use the formula and the data booklet to calculate the ΔcH°
We found all of the Δf Hm for the compounds
Are we finished with -802.5 kJ?? NO!
MOLAR ENTHALPY OF FORMATION
•
Methane is burned in furnaces and in some power plants. What is the standard
molar enthalpy of combustion of methane? Assume that water vapour is a
product.
•
This can also be communicated as an enthalpy change diagram. Note that the
labeling of the y-axis is different from that in a chemical potential energy diagram.
Epproducts – Epreactants
Molar Enthalpy of Formation
Practice
• How can molar enthalpies of formation be used to calculate enthalpies of a
reaction? Consider the slacking of lime, calcium oxide, represented by the
following chemical reaction equation
Your Task
• Questions #3abc, 4ab, 5ab,6ab,8abcd on page 514-15
• Questions #24,25 page 520
Enthalpy Changes
Three types of Change:
Phase - intermolecular bonds break/form
- 100 – 102 kJ/mol
Chemical - intramolecular bonds break/form
- 102 – 105 kJ/mol
Nuclear - nuclear binding energy break/form
- > 1010 kJ/mol
Enthalpy Changes Compared
Phase Changes
• When a substance changes phase there is no change in temperature.
• Thus there is no change in the average Ek of the substance
• Since heat (energy) has been added the Ep must have increased
• Example:
H2O(g ) 
 H2O(l )
Chemical Changes
• Chemical changes involve the rearranging of atoms which changes the total
bond energy of the system
• If the temperature is held constant only Ep will change
• Example:
CH4(g )  2O2(g ) 
 CO2(g )  2H2O(g )
Nuclear Change
• Nuclear reactions involve the rearrangement of nucleons resulting in a
change in Ep
• Example:
4
2

He  C 
 2   N
12
6
0
1
14
7
Energy Efficiency
•% Efficiency = E output x 100%
E input
Input – theoretical (Hess calculation)
Output – actual (calorimetry)
Energy Efficiency
http://www.iflscience.com/plants-and-animals/caution-dead-whale-contents-under-pressure
Reaction Rate
• Some reactions occur very quickly while others
occur more slowly.
Example: combustion of butane vs formation of rust
KINETICS = RATES OF REACTION
• Collision-Reaction Theory
•
A chemical sample consists of entities (ions, atoms, molecules) that are in
constant, random motion at various speeds, rebounding elastically from
collisions with each other (kinetic energy is conserved during elastic
collisions)
•
For a reaction to proceed, reactants must collide
•
An effective collision requires sufficient energy to react and the correct
orientation, so that bonds can be broken and new bonds formed
•
The more collisions there are, the greater the potential for effective collision.
KINETICS = RATES OF REACTION
• Collision-Reaction Theory
• Ineffective collisions involve entities that rebound and do not
rearrange and form new substances.
KINETICS = RATES OF REACTION
Factors affecting Reaction Rate:
• Concentration: more reactant particles in a given
volume increases the number of collisions per second
• Surface Area: more opportunity for collisions, the more
collisions there will be
• Temperature: the faster the particles are moving, the
more energy they have to create an effective collision
Bond Energy and Enthalpy Changes
• Bond energy is the energy required to break a chemical bond; it is also the energy released
when a bond is formed.
- bonded particles + energy  separated particles
- separated particles  bonded particles + energy
• The change in enthalpy represents the net effect from breaking and making bonds.
• ΔrH = energy released from bond making – energy required for bond breaking
• Exothermic reaction: making > breaking (ΔrH is negative)
• Endothermic reaction: breaking > making (ΔrH is positive)
• In general, the following rules regarding the energy of
reactions apply:
• If the total energy input is greater than the total energy output,
ΔrH° is positive, and the reaction is endothermic (the
reaction absorbs energy from the surrounding)
• If the total energy output is greater that the total energy
input, ΔrH° is negative, and the reaction is exothermic (the
reaction releases energy to the surroundings)
Potential Energy Diagram for the Formation of
Hydrogen Iodide
LET’S SEE IF YOU GET IT
Draw energy pathway diagrams for general endothermic and a general exothermic reaction.
Label the reactants, products, enthalpy change, activation energy, and activated complex.
ACTIVATION ENERGY OF A REACTION
Activation Energy – (EA)
•
The minimum collision energy
required for effective collision
•
Dependent on the kinetic energy of
the particles (depend on T)
•
Analogy: If the ball does not have
enough kinetic energy to make it over
the hill – the trip will not happen.
Same idea, if molecules collide
without enough energy to rearrange
their bonds, the reaction will not
occur. (ineffective collision)
https://www.youtube.com/watch?v=VbIaK6PLrRM
ACTIVATION ENERGY OF A REACTION
The activated complex
occurs at the at the
maximum potential
energy point in the
change along the energy
pathway.
Is this an exothermic or
endothermic change?
Exothermic. This means the initial
energy absorbed to break the
nitrogen-oxygen bond is less than
the energy released when a new
carbon-oxygen bond forms.
ACTIVATION ENERGY OF A REACTION
In general, the greater the EA, the slower
the reaction. It takes longer for more
particles to achieve kinetic energy
necessary for effective collision.
ACTIVATION ENERGY OF A REACTION
What does this diagram indicate?
At Temperature 2, a greater
number of particles will have the
activation energy required
Will this increase the rate of the
reaction?
Yes
ACTIVATION ENERGY OF A REACTION
Is this an exothermic or endothermic
change?
Endothermic. A continuous input of
energy, usually heat, would be needed to
keep the reaction going, and the enthalpy
change would be positive.
Energy Exchanges in Chemical Reactions
• Chemical reactions can produce energy, in exothermic
reactions, or absorb energy, in endothermic reactions
• The energy released or absorbed is linked with the
formation or breaking of bonds.
• In fact, all chemical reactions involve bond breaking
(endothermic) and bond making (exothermic) steps.
Bond Energy and Reactions
Bond Energy and Enthalpy Changes
• Bond energy is the energy required to break a chemical bond; it is also the
energy released when a bond is formed.
• The change is enthalpy represents the net effect from breaking and making
bonds
• Exothermic reactions:
• Endothermic reactions:
Your Task
• Question # 1-6, 8abc
CATALYSTS AND REACTION RATE
 A catalyst is a substance that increases the rate of a chemical reaction without
being consumed itself in the overall process.
 A catalyst reduces the quantity of energy required to start the reaction, and results
in a catalyzed reaction producing a greater yield in the same period of time than
an uncatalyzed reaction.
 It does not alter the net enthalpy change for a chemical reaction
Catalysts lower the activation energy,
so a larger portion of particles have the
necessary energy to react = greater
yield
CATALYSTS AND REACTION RATE
How do catalysts work??
 Scientists do not really understand the actual mechanism. Catalysts are also usually discovered
through trial and error.
 What they do know is that they provide an alternative, lower energy pathway from reactants to
products.
 Most of the catalysts (enzymes) for biological
reactions work by shape and orientation. They fit
substrate proteins into locations on the enzyme as a
key fits into a lock, enabling only specific molecules to
link or detach on the enzyme.
 Almost all enzymes catalyze only one specific reaction
CATALYSTS AND REACTION RATE
 Reaction Mechanisms
 Steps making up the overall reaction
 Each step = elementary reaction
 Reaction intermediates: substances formed in one elementary reaction and consumed in
another
The rate-determining step of a
reaction is the step with the highest
activation energy.
It is called the rate-determining step
because it is the slowest step.
Potential Energy Diagram for a Catalyzed and
Uncatalyzed Reaction
Practice
• Molecules A-A and B-B combine to form 2 A-B. The bond in A-B is stronger than
the bond in A-A or B-B. Draw an energy diagram to represent the given
reaction.
Uses of Catalysts
• The oil industry uses catalysts in the cracking and reforming of crude
oil and bitumen to produce consumer products (gasoline)
• Catalysts allow oil companies to increase the reaction rate while
decreasing the energy (often decreasing temperature) required.
Your Task
• These questions will be due at the start of class on Monday, NO EXCEPTIONS.
You will have today’s class and whatever is left after we go through your exam
on Friday to work through them.
• Practice Questions #1-3ab page 526
• Questions # 4,5abcd,7 page 531
• Questions #2, 3, 5ab, 7ab,8abc,9 page 534
• Read and take notes on page 540
• Questions #2abcd, 3abcde,4,5 page 542
• Question 16a-h page 546
• Question 10 page 547
• Question 13,14,16, 21abc, 22,23 page 548/49
• Question 31 page 550
The Technology of Energy Measurement
• “Calorimetry is the technological process of measuring energy changes
of an isolated system called a calorimeter”
Includes: Thermometer,
stirring rod, stopper or
inverted cup, two Styrofoam
cups nested together
containing reactants in
solution
Assumptions of Simple Calorimeter
• A polystyrene foam cup is assumed to be perfect insulation  small
amounts of energy transfer is ignored
• Only the water releases, or absorbs, heat energy to or from the reaction. It is
assumed that no heat is gained or lost to the surroundings unless otherwise
indicated.
• The specific heat capacity of solution= the specific heat capacity of liquid
water
• VSolution= VH2O= mH2O(l). When a solution’s concentration is less than 2mol/L,
water constitutes the majority of the solution. Liquid water has a density
very close to 1.00g/mL. This is assumed to be the density of aqueous
solutions if no extra information is given.
Equations
• The given assumptions allow you to calculate the kinetic energy change of
the calorimeter
Qcal= (𝑚𝑐∆𝑇)
• A combination of the first and second laws of Thermodynamics, where
energy is transferred rather than created or destroyed, allows for the
assumption that the kinetic energy change of the calorimeter is equal to the
enthalpy change of the calorimeter is equal to the enthalpy change of the
reaction ∆𝐻, given the following equation:
∆𝐻 = 𝑛∆𝐻°
• Combining ∆𝐻 = 𝑛∆𝐻° with the idea that energy gained by the water in a
calorimeter equals the energy lost by the reaction, the following equation
may be written as:
∆𝐻 = −𝑄 = 𝑚𝑐∆𝑇 ⇒ 𝑛∆𝐻° = −(𝑚𝑐∆𝑇)
Example
• When 50.0 mL of 1.00 mol/L NaOH, with an initial temperature of 17.4 degrees
Celsius is mixed with 50.0mL of 1.00 mol/L HCl, with an initial temperature of
17.4 degrees Celsius, the final reaction mixture reaches a temperature of 24.2
degrees Celsius. What is the molar enthalpy change for the neutralization of
the base?
Example
• When excess zinc is added to 50.0mL of 0.250 mol/L aqueous copper
(II) sulfate, the calorimeter warms by 12.5 degrees Celsius. What is
the molar enthalpy of reduction of aqueous copper(II) ions?
Energy - Recall Science 10:
A heating curve of water shows the types of heat energy that exist.
Ek – related to motion and directly related to temp.
Ep – related to bonds and bonding forces
Energy Measurement
•Kinetic heat energy is very easy to calculate
Ek = mc ∆ t
• Potential energy, can not be measured by may be deduced!
• There is no way that we can measure the energy stored between
two bonds.
• BUT - If a potential energy change occurs in an insulated container...
• AND – we consider the Laws of Thermodynamics…
• First Law of Thermodynamics:
Heat gained (q) = Heat lost (q)
• Second Law of Thermodynamics:
heat flows from a warmer body to a cooler body
• We can use a kinetic energy change to determine the energy stored
in bonds… indirectly!
Calorimetry
• Information about potential energy changes comes
from calorimetry experiments
• We will consider two types of calorimetry
measurements:
• Solution calorimetry using a simple calorimeter
• Bomb calorimetry
Limitation of Calorimetry
• Not every reaction of interest to chemists and
engineers can be studied by calorimetry so a method
of predicting enthalpies of reaction is an important
part of thermochemistry