Thermochemistry
3.3 Temperature
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


3 temperature scales
temperature conversions
change of state temperatures
absolute zero and energy
3.4 Energy
9.5 Energy in Chemical Reactions



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3.5 Specific Heat
 energy and Work
 heat (thermal energy)
 Joule (SI energy unit)
system & surroundings
enthalpy
heat of reaction(Δ enthalpy) & calculations
endothermic and exothermic processes


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
heat capacity & specific heat
uniqueness of water
heat equation & calculations
calorimetry
specific heat of water
10.5 Changes of State
 changes of state
 heat of fusion/ vaporization & calculations
 heating / cooling curves & calculations
I Can Statements:
(3.3) Temperature
1. State the name and explain the scalar differences between the following 3 Temperature Scales
 Fahrenheit
 Celsius
 Kelvin
2. Convert Kelvin temperatures to Celsius temperatures and vice-versa
3. Cite the temperatures at which phase changes occur for water on the Kelvin, Celsius, and Fahrenheit
temperature scales.
4. Define the term absolute zero and explain its significance in terms of kinetic energy of particles
(3.4) Energy
1.
2.
3.
4.
5.
6.
Explain the concept of the Internal Energy of a substance
Explain the relationship between the temperature , particle motion, and kinetic energy of a substance
Explain the relationship and dynamics of EK and Heat (thermal)Energy transfers between substances
Explain the concept of Potential and Kinetic Energy in terms of chemical bonds and particle position
Explain how chemical reactions change internal energy as heat energy is transferred between substances
Identify the Joule (J) as the SI unit of energy and the Calorie (Cal = kcal) as the food unit of energy
(3.5) Specific Heat
1. Explain the concepts of the Heat Capacity and Specific Heat of a substance and compare the heat capacity
of substances based on their specific heats
2. Explain how a substance’s specific heat can be used to calculate thermal energy changes for that substance
3. Explain each of the terms in the Heat Equation and use it to calculate thermal energy changes of and
between substances
4. Define the terms Calorimeter and Calorimetry and explain the use of a calorimeter
5. Use a calorimeter to calculate the thermal energy changes for a chemical reaction
6. Explain the significance of Water’s Specific Heat to global temperature regulation
(9.5) Energy in Chemical Reactions
1.
2.
3.
4.
Define and distinguish between the System and the Surroundings for a chemical reaction
Define the term enthalpy
Define the term heat of reaction and relate it to the change in the enthalpy of a system
Define the terms exothermic and endothermic and correctly apply the terms to heats of reaction for
various chemical processes
1
(10.5) Changes of State
1. State the name of all Changes of State and explain them in terms of energy entering or leaving the system
2. Define the terms Heats (Enthalpy)of Fusion and Vaporization and correctly apply them to enthalpy changes
occurring during changes of state
3. Calculate the heat of fusion and/or vaporization for water
4. Interpret the changes of state and energy information contained in a Heating/ Cooling Curve
5. Calculate the total energy change for water from below freezing to above boiling
Vocabulary 3.3-3.5, 9.5, 10.5
Absolute zero
Boiling
Calorimeter
Calorimetry
Celsius
Change of state
Condensation
Cooling/ Heating Curve
Deposition
Endothermic
Energy
Enthalpy(H)
Exothermic
Evaporation
Fahrenheit
Freezing
Heat (q)
Heat Capacity
Heat equation
Heat of fusion(ΔHfus)
Heat of reaction (ΔHrxn)
Heat of vaporization (ΔHvap)
Joule (J)
Kelvin(K)
Kinetic energy(EK)
Melting
Potential energy(EP)
Specific heat(c)
Sublimation
Surroundings
System
Temperature (T)
Thermal Energy
Work
Equations and Constants 3.3-3.5, 9.5, 10.5
Constants for Water
TK = TC + 273
q = m × c × ΔT
ΔT = Tfinal ─ Tinitial
q = m × Δ𝑯𝒗𝒂𝒑𝑯
q = m × Δ𝑯𝒇𝒖𝒔𝑯
𝐉
cice= 2.03
𝐠−𝐊
cwater = 4.184
𝟐𝑶
csteam = 1.841
𝟐𝑶
𝚫𝑯𝒗𝒂𝒑𝑯
Δ𝑯𝒇𝒖𝒔𝑯
𝟐𝑶
𝟐𝑶
=
=
𝐉
𝐠−𝐊
𝐉
𝐠−𝐊
𝟐𝟐𝟔𝟎 𝐉
𝐠
𝟑𝟑𝟒 𝐉
𝐠
Achievement Scale 3.3-3.5, 9.5, 10.5
Goal
3.3
Temperature
C Level
B Level
 Can name the 3 temperature scales
 Can convert oC to Kelvin and vice-versa
 Can cite the phase change temperatures of
water for the Kelvin and Celsius
temperature scales
 Can define absolute zero and determine its
value on the Kelvin and Celsius scale
 Can explain the scalar
differences between the 3
temperature scales
 Can explain the significance of
absolute zero in terms of
energy
A Level
2
Achievement Scale 3.3-3.5, 9.5, 10.5
Goal
3.4
Energy
3.5
Specific
Heat
9.5
Energy in
Chemical
Reactions
10.5
Changes
of
State
C Level
 Can explain Internal Energy and Heat (thermal) Energy
of a substance
 Can explain the relationship between temperature and
kinetic energy
 Can explain the relationship between EK and Heat
(thermal) energy transfers between substances
 Can explain how temperature changes are used to
measure heat energy
 Can explain kinetic and potential energy in terms of
chemical bonds or particle position
 Can explain how chemical reactions cause a
change in internal energy through a transfer of
heat (thermal) energy
 Can identify the Joule and Calorie as energy units
 Can define the concepts and compare the heat
capacity and specific heat of various substances
 Can explain how the specific heat of a substance can
be used to calculate changes in the thermal energy of
that substance
 can explain each of the terms in the heat equation
 Can define calorimeter and calorimetry and explain
how a calorimeter is used to measure thermal energy
changes
 Can use a calorimeter to measure thermal energy
changes
 Explain how the specific heat of water helps to
regulate Earth's weather
 Can define and distinguish between system and
surroundings for a chemical reaction
 Can define and relate the terms enthalpy and heat of
reaction
 Can define and correctly apply the terms exothermic
and endothermic to the heats of reaction for various
chemical processes
 Can state the name of all changes of state and explain
them in terms of energy entering or leaving the system
 Can define the terms heats (enthalpy) of fusion and
vaporization and correctly apply them to enthalpy
changes occurring during changes of state
 Can interpret the change in state and energy
information contained in a heating/cooling curve
B Level
A Level
 Can use the heat
equation to calculate
thermal energy changes
of a substance
 Can use the
heat equation
to calculate
energy changes
between two
substances
 Can calculate the
enthalpy change for a
chemical reaction and
state if that change for
was exothermic or
endothermic based on
the mathematical sign
 Can calculate the heat of  can calculate the
fusion or vaporization
energy change
for a substance
for water from
solid through gas
phase when
given a heating/
cooling curve
3
Sample Questions 3.3-3.5, 9.5, 10.5
C Level:
1. What is the name of the SI unit of temperature? Kelvin
2. What is the freezing temperature of water on the Celsius scale? 0oC
3. Convert the boiling point of water on the Kelvin scale to the Celsius temperature. 373 - 273 = 100oC
4. What is 'absolute zero'? the lowest possible temperature on the Kelvin scale = 0 Kelvin
5. Which has more internal energy - an iceberg or a boiling cup of coffee? Explain
an iceberg has more internal energy - b/c it has more mass than the cup of coffee
6. Which substance has particles with more kinetic energy - an iceberg or a boiling cup of coffee? Explain
the particles of a boiling cup of coffee have more EK - the greater the temperature the greater the EK
of the particles
7. In which direction will heat energy be transferred when an iceberg and a boiling cup of coffee are put into
physical contact with one another? - when will the transfer stop?
heat energy will be transferred from the boiling cup of coffee(substance with the higher
temperature) to the iceberg (substance with the lower temperature) - energy will always travel from
the substance with greater temperature to the substance with lower temperature - the transfer will
stop when both substances are at the same temperature
8. When the following reaction is performed the temperature increases dramatically - in other words, heat is
one of the products of this reaction.
CH4 (g) + 2 O2(g) → CO2 (g) + 2 H2O (l) + heat(energy)
ΔERXN = − 890 kJ/ mol
Explain the meaning ΔERXN = − 890 kJ/ mol in terms of bonding and energy changes for the reaction.
bonds broken = energy IN ........ must break bonds of the reactants
bonds made = energy OUT......must make the new bonds of the products
ΔERXN = − 890 kJ/ mol = the amount of energy change for the reaction;
−ΔERXN = energy was released by the reaction (energy IN < energy OUT)
+ ΔERXN = energy was absorbed by the reaction (energy IN > energy OUT)
9. When a chemical reaction is carried out in a coffee cup calorimeter, explain what is the system and what
are the surroundings.
system = the thing(s) which will be the cause of the energy change
surroundings = the thing(s) which will register the energy change
10. A student dissolved some solid ammonium nitrate in water in a coffee cup calorimeter to form an
aqueous solution of ammonium nitrate. After calculating the average energy change for three trials the
student determined the following:
energy + NH4NO3(s) → NH4NO3(aq)
ΔESolution = + 330 J/ g
State whether the temperature increased or decreased during this process. Explain
DECREASED = b/c the ΔESolution = + 330 J/ g
+ΔESol'n = energy IN = temperature goes down
4
11. Explain how the process of calorimetry allows for the measurement of energy changes.
for 2 substances in thermal contact and in a closed system.......
...........the energy lost by one substance must be equal to the energy gained by the other substance
in calorimetry a reaction is carried out in water (or a solution) of a known temperature and mass. A
thermometer is used to measure the change of temperature of the water (or solution) caused by the
reaction. This temperature change is then used in the following mathematical relationship:
mathematical representation:
heat lost/gained by the water = heat lost/ gained by the chemical reaction
qlost/gained H2O = qlost/gained chemical reaction
heat lost/gained by the water is determined by the following:
qH2O = mH2OSHH2OΔTH2O
mH2O = mass of water
SHH2O = specific heat of water
ΔTH2O = change in temperature of water
12. What are heat of reaction and enthalpy and how are they related.
heat of reaction = the change in energy that occurs during a chemical reaction
enthalpy = the scientific term for heat of reaction occurring at a constant pressure
13. Is the enthalpy of the following reaction exothermic or endothermic? Explain
energy + NH4NO3(s) → NH4NO3(aq)
ΔH = + 330 J/ g
ENDOTHERMIC = b/c the ΔESolution = + 330 J/ g
+ΔESol'n = energy IN = temperature goes down during endothermic processes
14. Snowing is an example of what phase change? What is the NAME of this phase change?
phase change = gas - to - solid
name = deposition
15. To calculate the enthalpy change when liquid water changes to ice, would you need ΔHfus or ΔHvap?
Explain
ΔHfus = b/c it measures the change in energy when a liquid changes to a solid
16. Use the following heating curve for water to answer the questions at the right.
Identify the line segment for each of the following:
d.
d.
b.
e.
a.
b.
c.
1.
2.
3.
4.
5.
6.
7.
the boiling of water?
ΔHvap
a water-ice combination
all gas
all solid
ΔHfus
the heating of liquid water
5
B Level:
1. What the significance of 'absolute zero'?
lowest temperature on the Kelvin scale - where there is no longer any kinetic energy - meaning no
particle movement
2.. Which of the three temperature scales measures the smallest change in temperature? Explain.
the Fahrenheit Scale measure the smallest change in temperature : 1.8oF = 1oC and 1 Kelvin
1oC = 1 Kelvin
3.
(a) 25.0 g of liquid water are cooled from 55.0oC to 0oC. Calculate the energy released in kJ, as this
water was cooled.
q = mH2O × SHH2O × ΔTH2O and don't forget that
ΔT = Tfinal ─ Tinitial
𝐉
q = 25.0g x 4.184 𝐠−𝐊 x (0oC − 55.0oC) → 105 J x (− 55.0) → − 5575 = − 5.58 kJ
(b) Was this process exothermic or endothermic? Explain
exothermic = b/c the sign on the answer was negative (−) which indicates energy was released
(d) Calculate the energy chang, in kJ, which would occur if this water were then turned completely to
ice.
q = m × Δ𝑯𝒇𝒖𝒔𝑯
q = 25.0 g x
𝟐𝑶
𝟑𝟑𝟒 𝐉
𝐠
and
Δ𝑯𝒇𝒖𝒔𝑯
𝟐𝑶
=
𝟑𝟑𝟒 𝐉
𝐠
= 8350 → − 8.35 kJ
A Level:
Use this heating curve for water for the following.
1. Calculate the amount of heat energy
needed to warm and completely melt
10.0 g of ice from the lowest temperature on this graph
step 1: heat ice from −50.0 oC to 0.0oC
𝐉
10.0 g x 2.03
[0.0oC − ( −50.0 oC )] = + 1020 J
𝐠−𝐊
step 2: melt the ice
10.0 g x
𝟑𝟑𝟒 𝐉
𝐠
= + 3340 J
step 3: sum the answers
= 4360 J
6