Download Thermochemistry Note

Thermochemistry/Thermodynamics – Unit one
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Students will determine and interpret energy changes in chemical reactions.
Students will explain and communicate energy changes in chemical reactions.
Definitions:
Thermochemisty: The study of heat changes that occur during chemical reactions.
Thermochemistry is interested in the flow of heat from the system to the surroundings or
vice versa.
System: The area, usually contained, that we focus our attention on.
Eg. Chemicals in a flask undergoing a reaction
Surroundings: The region in the vicinity of the system.
Eg. A water bath surrounding the flask
Energy: The capacity for doing work or supplying heat
Kinetic Energy – A molecules’ random, chaotic motion.
Potential Energy – Stored energy in chemicals due to their composition.
-The bonding forces between and within molecules.
-It depends on the kinds of atoms and their arrangements within the
substance.
Eg. Fuels (coal, oil, gas) have high chemical potential energy whereas
their combustion products (carbon dioxide and water) have relatively low
chemical potential energy. Where did the energy go?
Heat: (a form of energy) -Energy that is transferred from one body to another because of
a temperature difference. Heat will always flow from a warmer body to a cooler body
until both temperatures are the same. Heat  Temperature
Temperature: The average kinetic energy of the molecules within the substance being
measured.
A system’s internal energy is the total energy of the system. This cannot be measured,
but the heat changes that accompany a chemical or physical change can be.
***All chemical reactions involve the release or absorption of heat.
***No energy is created or destroyed, only transferred. What law is this?
Laws of Thermodynamics (James Joule)
1. Energy can be converted from one form to another but is conserved in an isolated
system.(no energy or matter enters or leaves)
In=out
Gain=loss
Kinetic  Potential
Chemical  Thermal
2. Heat flows spontaneously from a hot object to a cool one.
ENDOTHERMIC – Energy in put INTO a body (positive in sign)
(add heat)
Reactants +heat Products
EXOTHERMIC – Energy is TAKEN OUT OF (EXITS) a body (negative in sign)
(heat lost)
Reactants  Products +heat
TYPES OF CHANGES IN A SYSTEM:
1. Temperature Change
2. Phase Change
3. Chemical Change
4. Nuclear Change
Comparision:
TEMPERATURE CHANGES
 You will recall the application of Q = mc∆t to the analysis of heat transfer
***A change in temperature is a change in the movement (Kinetic Energy) of the
molecules but involves no significant change in chemical bonding (Potential
Energy).
Mathematics:
Q=mct
Where:
Q (∆H)= heat energy gained or lost (depending on the sign) kJ or J
m =mass of the substance undergoing the change (*1mL of water = 1g) kg or g
c =the specific heat capacity of the substance J/goC
t = the change in temperature (final minus initial) oC
CHEMICAL CHANGE
 You will explain, in a general way, how stored energy in the chemical
bonds of hydrocarbons originated from the sun
 You will define enthalpy and molar enthalpy for chemical reactions
 You will use and interpret ∆H notation to communicate and calculate
energy changes in chemical reactions
 You will explain the energy changes that occur during chemical reactions,
referring to bonds breaking and forming and changes in potential and
kinetic energy
A change in which one substance is converted into another substance. The energy
involved can exhibit itself as heat, light, mechanical or electrical energy.
What type of bonds would have to be broken/formed for this to occur?
Is this a change of Potential or Kinetic energy to the system (substance)?
Therefore called ENTHALPY!
They can be either exothermic or endothermic just as the other types of changes are.
Eg. Burning of fossil fuels – ________________
Note that this requires some energy input to occur regardless of which type it is.
***The rearrangement of chemical bonds requires energy. Energy is added to separate
the atoms within a molecule that are held together by a chemical bond. This energy is
stored as potential energy in the separate atoms. The energy is then released when the
particles are allowed to come together to form new molecules (bonds).
Bond breaking (reactants) is generally endothermic (positive sign)
Bond making (products) is generally exothermic (negative sign)
Eg. 2 H2O(l) 2H2(g) + O2(g)
Strong bonds in the water make it very endothermic to break them.
The new bonds in the gases release energy when formed but not as much as was
required to break the water molecules’ bonds.
The overall reaction is therefore endothermic. ∆H = +285.8kJ/mol of water
**Whether the entire reaction is endo or exothermic depends on which side is higher in
energy. If more energy is released when forming products than required from breaking
reactants, the reaction is exothermic and vice versa.
Does potential energy increase or decrease with an exothermic reaction? Why?
What about an endothermic reaction?
Bond Energy – the energy required to break a bond (the same amount of energy is
required to form that bond)
-‘r’ represents the chemical reaction –can be replaced with’f’, ‘dr’, ‘sr’,etc
-‘m’ represents that the value is ‘molar’
where:
rH = the total amount of heat energy gained or lost in the system. Also termed the
Enthalpy of Reaction, or Heat of Reaction (kJ)
n = the number of moles of one of the substances involved in the reaction (mol)
∆rHm = the amount of heat gained or lost PER MOLE of the same substance above. Also
termed the Molar Enthalpy of Reaction, or the Molar Heat of Reaction. (kJ/mol)
Eg. Find the enthalpy if 3.294g of pentane (∆cHm = -3272.8kJ/mol) burn in excess O2 .
If this energy is used to heat 7.0L of water at 29.0 C, what is the final temperature of the
water?
Potential Energy Diagrams (Energy Level Diagram)
 You will classify chemical reactions as endothermic or exothermic, including
those for the processes of photosynthesis, cellular respiration and hydrocarbon
combustion.
 You will compare energy changes associated with a variety of chemical reactions
through the analysis of data and energy diagrams.
 You will analyze and label energy diagrams of a chemical reaction, including
reactants, products, enthalpy change and activation energy
 You will draw and interpret enthalpy diagrams, indicating changes in energy for
chemical reactions
The change in energy must be Potential in nature to show any change on a potential
energy diagram. Therefore, what type(s) of changes could NOT be demonstrated this
way?
-An increase in Potential energy is endothermic (kinetic energy was absorbed and used to
break bonds creating more freedom of movement of molecules-stored energy). Eg.
Melting
-A decrease in potential energy is exothermic (bonds are formed causing the release of
energy usually converted to heat in the surroundings). eg. Condensation
Activation Energy- read pg 524-530
 You will define activation energy as the energy barrier that must be overcome for
a chemical reaction to occur
 You will explain that catalysts increase reaction rates by providing alternate
pathways for changes, without affecting the net amount of energy involved; e.g.,
enzymes in living systems.
The extra energy required to initiate the reaction (the spark for the fire) - to get
chemicals to transition state.
-A catalyst lowers the activation energy for a given reaction. The key is that it
does NOT change the overall enthalpy of reaction.
Ways of communicating Enthalpy/ Energy Changes for a Chemical reaction:
 You will write balanced equations for chemical reactions that include energy
changes
eg. a. In a reaction of sulfur dioxide with oxygen to form sulfur trioxide (as shown
below), if 98.9 kJ of energy is given off for every mole of SO2(g) then how much energy
will be given off for the reaction as written?
2SO2(g)+ O2(g)2SO3(g)
b. Write the energy change as a term in the reaction.
c. Draw a potential energy diagram for this reaction.
d.
What is the molar enthalpy for sulfur trioxide?
oxygen?
CALORIMETRY
 use calorimetry data to determine the enthalpy changes in chemical reactions
 You will identify that liquid water and carbon dioxide gas are reactants in
photosynthesis and products of cellular respiration and that gaseous water and
carbon dioxide gas are the products of hydrocarbon combustion in an open system
 You will design a method to compare the molar enthalpy change when burning
two or more fuels (e.g., octane, propane, ethanol and historic fuels such as seal or
whale oil), identifying and controlling major variables.
 You will perform calorimetry experiments to determine the molar enthalpy
change of chemical reactions.
 You will explain the discrepancy between the theoretical and actual efficiency of
a thermal energy conversion system
-A method of finding the enthalpy of a reaction.
Energy is conserved. In an isolated system, the energy given off from one substance (or
reaction) has to be gained by some other substance.
A calorimeter is designed to be an isolated system, in other words contain the energy
(heat) of the system.
Therefore: heat lost H = heat gained H
The key is that heat lost has a negative sign whereas heat gained is positive in sign. So
they are equal in amount, but not sign. This has to be accounted for mathematically.
Hreaction = -Hcalorimeter
nHm= -mc∆t
*there are many assumptions that involved in calorimeters – know them-pg 486
*bomb calorimeters involve combustion reactions in a closed system – water as a product
is under pressure and therefore in liquid form!
*calorimeters that are used repeatedly have a constant mass and constant specific heat
values to all their contents. These are therefore precalculated as a ‘calorimeter heat
capacity-C ’ ‘mc’ in the equation is replaced by ‘C’
Eg.1: A student mixed 100.0mL of 1.50 mol/L sulfuric acid with 200.0mL of 1.50 mol/L
sodium hydroxide. Both solutions were at 19.67 oC initially and the highest temperature
reached by the reaction mixture was 34.06 oC. Calculate the molar enthalpy of
neutralization for sodium hydroxide.
Eg. 2: A bomb calorimeter has a heat capacity of 3.78 kJ/oC. The complete combustion
of 4.56 g of octane caused the temperature to change from 21.6 oC to 64.8 oC. Calculate
the molar enthalpy for the combustion of octane.
HESS’S LAW
 You will explain and use Hess’ law to calculate energy changes for a net reaction
from a series of reactions
This law is used to allow for determination of enthalpy change of a reaction without
having to perform a calorimetry lab.
If the sum of two or more thermochemical equations (of known enthalpies) result in the
final equation (unknown enthalpy), then the enthalpies can also be added to give the
enthalpy for the final equation.
Rules:
1. If a chemical equation is reversed, then the sign of the enthalpy must be reversed
2. If the coefficients of a chemical equation are altered by multiplying or dividing by
a factor, then the enthalpy is altered the same way.
Eg.
H from MOLAR HEATS OF FORMATION
 You will predict the enthalpy change for chemical equations using standard
enthalpies of formation
 You will compare energy changes associated with a variety of chemical reactions
through the analysis of data and energy diagrams
This is another method to determine enthalpy for a reaction without having to perform
calorimetry and a shortcut to Hess’ law depending on the info available.
Molar heats of formation can be found in the data booklet for many substances. The
molar heat of formation is defined as the amount of heat released or absorbed when one
mole of a compound is formed from its elements at standard conditions. For free
elements, this value is zero.
When looking at a chemical reaction, it can be interpreted that in order for the chemicals
to rearrange, the reactants must be broken down (decomposed) into elements as an initial
step, and then the products are formed from the elements. So the heats of formation can
be used in this same method.
The heats of formation for the reactants must be reversed (sign changed) if they were
decomposed.
The heats of formation for the products can remain as they are since the products were
formed!
By adding all these values together, we are performing Hess’s Law to find the overall
heat of reaction. In short form:
H = nfH (products) - nfH (reactants)
nfH is the sum of (numbers of moles of substance x the heat of formation for the
substance).
Steps:
1. Write a balanced equation for the reaction (use a coefficient of 1 for the substance
for which you are determining the molar heat of. That way rH = ∆rHm)
2. Obtain molar heats of formation of reactants and products from the data sheet and
write them under the compound.
3. Multiply each ∆rHm by the moles (balancing number in the reaction) for each
reactant and product.
4. Add the values together for products. Do the same for reactants.
5. Subtract the reactants value from the products as per the equation. This value is
the rH for the reaction as written. (If using this value further, convert it to the
molar heat of reaction rHm for the substance specified.)
Eg. Calculate the molar enthalpy for the combustion of octane forming gaseous products.
Find the percent error with the calorimetry example done previously in the notes.