CP Chapter 17 Thermochemistry 2014-2015
Thermochemistry
• Thermochemistry is the study of energy __________________ that occur during chemical and
physical changes (changes of state)
The Nature of Energy
• Energy is the ability to do work or produce __________________
• Heat, q or Q, is _____________; flows due to temperature differences (always _________ to
__________)
Law of Conservation of Energy
• Energy CANNOT be created nor ______________________; only converted into different
types
Kinetic vs. Potential Energy
• Two main types of energy – kinetic and potential
o Kinetic – energy of _________________
o Potential – energy due to position or energy ____________________ in chemical
bonds
§ Chemical potential energy - the energy stored in a substance because of its
composition
§ Example: gasoline
Temperature vs. Heat
• Temperature is a measure of the _____________________ in a sample.
• Temperature is a _______________________________ of heat.
• Heat is the total energy of molecular motion, dependent upon amount, size, and type of
particles. Heat is _______________________.
Units of Heat
• calorie - the amount of heat required to raise the temperature of one gram of pure
_________________by one degree Celsius
• Calorie – nutritional calorie;
o 1 Calorie = 1000 calories = 1 kilocalories (kcal)
• Joule – SI unit of heat
o 1 calorie = ________________________J
Converting Energy Units
Calorie/calorie/ kilocalorie
calorie/Joule
Ex 1) A cereal has 155 nutritional Calories per serving. How many calories, kilocalories and Joules is
this?
Ex 2) A person on a diet consumed 1350 Calories in one day. How many calories, kilocalories and
Joules is this?
1 System and Surroundings
• Universe = system + surroundings
• System – The part of the universe you wish to _____________. In Chemistry this is your
chemical ______________/physical process.
• Surroundings – Everything else in the ___________________.
• When heat is transferred it can flow in or out of the system
Endothermic vs. Exothermic
• An Exothermic process is one that _____________________ (evolved) heat to its
surroundings (feels __________)
o Energyproducts < Energyreactants
• An Endothermic process is one that _____________________ heat from the surroundings
(feels _________)
o Energyproducts > Energyreactants
Exothermic Process
Endothermic Process
E________ > E________ E________ < E________ Q and heat flow
• Exothermic process, heat is _________________, q is _________________.
• Endothermic process, heat is _____________________, q is positive (+)
Specific Heat = Cp
• Specific heat of a substance is the amount of heat required to raise the temperature of one
gram of that ________________________ by one degree Celsius.
• Unit for specific heat is J/goC
• Each substance has a _________________________ specific heat
• Water = 4.184 J/goC
• Gold = 0.129 J/goC
• Copper = 0.386 J/goC
• The lower the specific heat the lower the amount energy is required to raise its
temperature.
2 Calculating Heat Released and Absorbed
• Q = mCpΔT
o Q = Heat (J)
o m is _________(gram)
o ΔT (______) is temperature change Tfinal -Tinitial
o Cp is specific heat at a constant pressure
Calculating Heat
Ex1:If the temperature of 56.6 g of ethanol increases from 45.0oC to 80.0oC, how much heat has
been absorbed by the ethanol? Specific heat of ethanol = 2.44 J/goC
Ex2: A 4.00 g sample of a substance was heated from 274 K to 314 K and absorbed 32 J of heat,
what is the specific heat of the substance?
Ex3: If 98,000 J of energy are added to 6200 g of water, what will the change in temperature of the
water be? Specific heat of water = 4.184 J/goC
Calorimetry and Enthalpy
Calorimetry
• Calorimetry is a laboratory ____________________used to measure heat flow
o Based on
o the law of conservation of ________________
o Idea that the heat released by the system _________________heat absorbed by the
surroundings or vice versa; -q = q
Calorimetry Example
• An ice cube is added to a warm cup of water.
o The amount of heat used to melt the ice cube is the same amount of heat lost by the
warm water.
Calorimeter
• Insulated device used to measure ________________ flow
• Either measures the heat absorbed ______ released during chemical/physical change
3 Enthalpy and Enthalpy Changes
• Enthalpy (H) is the _____________ content of a system at constant ____________
• Thermochemistry uses the ______________ in enthalpy(ΔH) to study heat changes
• At constant pressure, Q = ΔH
• Heat of reaction (ΔHrxn) – The change of enthalpy in a _____________________ reaction
Endo, Exo, q, and ∆H
• All chemical and physical changes release or absorb heat
• Exothermic: _____________ heat; Q is negative; ∆H is ________________
• Endothermic: absorbs _________; Q is positive; ∆H is positive
Thermochemical Equation
• equation that includes the physical states of all _____________ and products (1 atm, 25oC)
and the energy ____________, ΔH
• Endo) NH4NO3(s) à NH4+(aq) + NO3-(aq) ΔH = 27 kJ
• Exo) CaO(s) + H2O(l) à Ca(OH)2(s) ΔH = -65.2 kJ
Thermochemical Equations for Endothermic Reactions
• Endothermic : ΔH is positive; heat is on the _______________________ side of the equation
(Heat absorbed)
• NH4NO3(s) à NH4+(aq) + NO3-(aq) ΔH = 27 kJ
• NH4NO3(s) + Heat à NH4+(aq) + NO3-(aq)
• NH4NO3(s) + 27kJ à NH4+(aq) + NO3-(aq)
Thermochemical Equations for Exothermic Reactions
• Exothermic : ΔH is negative; heat is on the _________________________ side of the
equation (Heat released)
• CaO(s) + H2O(l) à Ca(OH)2(s) ΔH = -65.2 kJ
• CaO(s) + H2O(l) à Ca(OH)2(s) + Heat
• CaO(s) + H2O(l) à Ca(OH)2(s) + 65.2 kJ
Heat of Combustion
• Heat of combustion is the enthalpy change for the complete burning of a substance
• CH4(g) + 2O2(g) à CO2(g) + 2H2O(l) ΔH = -891kJ
• CH4(g) + 2O2(g) à CO2(g) + 2H2O(l) + Heat
• CH4(g) + 2O2(g) à CO2(g) + 2H2O(l) + 891kJ
Ex 1) A hot piece of metal (at 155 °C) with a mass of 4.68 g is placed into 53.9 grams of water at 22
°C. The water (Cp =4.184 J/g °C) heats up to 37.1 °C. What is the specific heat of metal?
Ex 2) How much heat is released by the combustion of 250.0 g of octane, C8H18?
Octane ΔHcomb = -5471 kJ/mol
4 Changes of State (Phase Changes)
Phase Change
• A phase change is going from one _____________________ of matter to another (Physical
change)
o Gas to liquid
o Liquid to solid
• Changes of state either ____________________ or ________________________
(evolve/liberate) energy
• Endothermic phase changes (Absorb energy)
o Melting (solid to liquid)
o Evaporation/___________________________ (liquid to gas)
o Sublimation (________________ to gas)
• Exothermic phase changes (Release energy)
o ____________________________ (gas to liquid)
o Freezing (liquid to solid)
o Deposition (____________ to solid)
Temperature and Phase Change
• During a phase change, as energy is added or removed, there is _________ temperature
change.
Latent Heat
• Since there is no __________________________ change, we cannot use Q = mCPΔT to
calculate enthalpy change.
• Instead we use latent (hidden) heat, which is the quantity of heat released or absorbed during
a phase change.
• Units are usually Joule/gram or kilojoule/mole
Changes of State
• Latent Heat of Vaporization is the heat required to ________ one mole of a liquid
o Going from liquid to gas
o For water, ΔHvap = 40.7 kJ/mol
• Latent Heat of Condensation is the heat required to ________ one mole of a gas
o Going from gas to liquid
o For water, ΔHcond = -40.7 kJ/mol
Thermochemical Equations for Vaporization and Condensation
• The reverse process of vaporization is condensation
o The _________________________ process has an opposite sign
o ΔHvap = - ΔHcond
o H2O(l) à H2O(g) ΔHvap = 40.7 kJ/mol
o H2O(g) à H2O(l) ΔHcond = -40.7 kJ/mol
Changes of State
• Latent Heat of Fusion is the heat required to _______________________ one mole of a solid
substance.
o Going from solid to liquid
o For water, ΔHfus = 6.01 kJ/mol (Endothermic)
5 •
Latent Heat of Solidification is the heat required to _____________________ one mole of a
liquid substance.
o Going from liquid to solid
o For water, ΔHsolid = -6.01 kJ/mol (Exothermic)
Thermochemical Equations for Fusion and Solidification
• The reverse process of fusion is solidification.
o The reverse process has an ________________________ sign
o ΔHfus = - ΔHsolid
o H2O(s) à H2O(l) ΔHfus = 6.01 kJ/mol
o H2O(l) à H2O(s) ΔHsolid = -6.01 kJ/mol
Heating Curve
Explanation of Heating Curve
• When there is a temperature change use Q=m CpΔT to calculate the amount of energy used.
• When the heating curve is flat (no temperature change), there is a phase change. Use the
latent heat to calculate the amount of energy used.
Calculate the heat required to melt 45.6 g of water at its melting point.
Latent heat of fusion (∆Hfus) = 6.01 kJ/mol
Calculate the heat evolved to condense 75.4 g of water at its boiling point
Latent heat of vaporization (∆Hvap) = 40.7 kJ/mol
6 What mass of ammonia (NH3) must be vaporizes that absorbs 345 kJ of heat?
Latent heat of vaporization (∆Hvap) = 23.3 kJ/mol
Hess’s Law - Calculating Enthalpy Change (Heat of Reaction)
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•
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Hess’s Law states that you can ___________ two or more ____________________ equations
to produce a __________________ equation,
the reaction enthalpy of unknown is the ______________ of the related reactions
Ex: H2(g) + ½ O2(g) à H2O(g)
DH = -242 KJ
H2O(g) à H2O(l)
DH = -44 KJ
H2(g) + ½ O2(g) à H2O(l)
DH = -286 KJ
General principles for combining thermochemical equations.
1. If a reaction is reversed, the sign of ΔH is reversed
2. If the coefficients are multiplied, then the ΔH is multiplied by the same factor.
Calculate the enthalpy change for:
Ex1: 2S(s) + 3O2(g) à 2SO3(g)
∆H = ?
S(s) + O2(g) à SO2(g) ∆H = -297 kJ
2SO3(g) à 2SO2(g) + O2(g)
∆H = 198 kJ
Ex2: 2NO(g) + O2(g) à 2NO2(g)
∆H = ?
N2(g) + O2(g) -à 2NO(g) ∆H = 180.6 kJ
N2(g) + 2O2(g) à 2NO2(g) ∆H = 66.4 kJ
Ex3: 2C (s) + O2(g) à 2CO(g) ∆H = ?
C(s) + O2(g) à CO2(g)
∆H = 393.5 kJ
2CO + O2 à 2CO2
∆H = -566kJ
Ex4: C(s) + 2H2(g) à CH4(g)
∆H = ?
C(s) + O2(g) à CO2(g)
∆H =-393.5 kJ
2H2(g) + O2(g) à 2H2O(l)
∆H =-571.6 kJ
CH4(g) + 2O2(g) à CO2(g) + 2H2O(l) ∆H = -890.8 kJ
7 Standard Heat of Formation
•
•
•
Standard heat of formation is defined as the _______________ in ______________ that
accompanies the ___________________ of one mole of the compound at its
________________ state from its __________________ in their standard state.
Standard state is _________________________________________
Every free element in its standard state is assigned a ∆Hf of _____________. Free elements
include ___________________ elements and any ___________________ molecule
(H2O2N2Cl2Br2I2F2)
Heat of Formation Equation
• Ex: Ca(s) + C(s) + 3/2 O2(g) à CaCO3(s)
• Practice: Write a formation equation for 1 mole of Na2CO3(s)
ΔHrxn =
Use the standard enthalpies of formation to calculate ∆Hrxn
Ex1: CH4(g) +2O2(g) à CO2(g) + 2H2O(l)
∆Hf(CO2(g)) = -394kJ
∆Hf(H2O(l)) = -286 kJ
∆Hf(CH4(g)) = -75kJ
∆Hf(O2(g)) = 0 kJ
Use Table C-13 to calculate ∆Hrxn for the following reactions
Ex2: 2HF(g) à H2(g) + F2(g)
Ex3: 2H2S(g)
+ 3O2(g) à 2H2O(l)
+ 2SO2(g)
Ex4: 4Fe(s) + 3O2(g) à 2Fe2O3(s)
Ex5: 2NO(g)
+ O2(g)
à 2NO2(g)
Ex6: 2CO(g)
+ O2(g)
à 2CO2(g)
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