Name: ___________________________________
CHAPTER 17: Thermodynamics
Period:___________________
SPECIFIC HEAT CAPACITY AND CALORIMETRY
1. Units of heat (with abbreviations):
Joule (j):
Kilojoule (kj) – 1000joules are in 1 kilojoule
calorie (cal):
kilocalorie- also known as a food calorie (kcal or Cal): 1000 cal are in 1 kilocal or 1 Cal
We can convert between the different units of heat. The numerical relationships or
conversion factors are:
4.18J = 1 cal
1000 cal = 1 Cal or 1 kcal (also known as a food calorie)
1000 J = 1 kJ
1 kcal = 1 Cal= 4.18 kJ
Convert the following.
2500J = ? cal
4.25 Cal = ? J
11.2 kJ = ? Cal
2. Specific Heat Capacity
STOP AND THINK: Why is it better to stir hot food with a wooden
spoon, rather than a metal spoon?
Metal
Handle
Wooden
Handle
Specific heat capacity:
High Specific Heat Capacity v. Low Specific Heat Capacity
1
STOP AND THINK: Look at the table of specific heat capacities below and answer the questions that follow.
NOTE: Water has the highest specific heat in the table and silver has the smallest specific heat in the table.
Which substance heats up the quickest?_______________________________________
If equal masses of each of the substances above absorb 1000J of heat, which substance would become the
hottest and which would be the coolest ?
___________________________________________________________________________________
Why?_____________________________________________________________________________________
__________________________________________________________________________________________
__________________________________________________________________________________________
2
3. Calculations Involving Specific Heat - https://youtu.be/Ak7PN8tn4cU or
https://youtu.be/c2NaAHEusm4
mass
q is heat
(joules)
q = m x c x ΔT or
Heat = m x (sh) x ΔT
where ΔT (change in temperature) = (Tfinal – T initial)
Examples and Practice:
a. The specific heat capacity of aluminum is 0.897 J/g°C. Determine the amount of heat released when
10.5g of aluminum cools from 240°C to 25°C.
b. The specific heat of lead is 0.129 J/g°C. How much heat is necessary to raise the temperature of
45.0g of lead from 25°C to 38°C?
c.
The specific heat capacity of aluminum is 0.897 J/g°C. What mass of aluminum can be heated from
33°C to 99°C using 450J of heat?
d. A 16-g piece of iron absorbs 1090 joules of heat energy, and its temperature changes from 25ºC to
175ºC. Calculate the heat capacity of iron.
3
https://youtu.be/llsBOI7BORI (but I do heat lost + heat gained = 0)
4. Calorimetry: is the study of heat flow. Heat flows from hot to cold.
In a calorimeter, the heat gained by the water equals the heat lost by an object (i.e. metal) or produced by a
reaction.
So, heat gained by the water + heat lost by the metal (or produced in the reaction) = 0
They are equal and opposite values: gained is positive and lost is negative
1. 55.0g of a “mystery metal” at 93°C is placed in a calorimeter containing 100.0g of water at 25°C. The
metal sits in the water until the temperature levels off at 29°C. At this point, both the metal and the
water are at 29°C. The specific heat capacity of water is 4.18 J/g°C.
METAL
specific heat capacity
Tf
Ti
ΔT
WATER
4.18 J/g°C
Heat lost by the metal (negative) + Heat gained by the water (positive) = 0 (zero)
(mass metal)(cmetal)(Tfinal – Tinitial) + (mass water)(cwater) )(Tfinal – Tinitial)=0
(
)(
)(
)+(
)(4.18)(
)=0
4
2. 55.0g of a metal are heated to 112ºC, and then placed in a coffee cup calorimeter containing 60.0g of
water at 32ºC. The final temperature in the calorimeter is 42ºC. What is the specific heat of the metal?
METAL
mass
specific heat capacity
Tf
Ti
ΔT
WATER
4.18 J/g°C
__________________
3. The specific heat of aluminum is 0.902 J/gºC. 15.0g of aluminum are heated to 115ºC, and added to a
calorimeter containing water at 25ºC. The final temperature in the calorimeter is 45ºC. What mass of
water was in the calorimeter?
METAL
mass
specific heat capacity
Tf
Ti
ΔT
WATER
4.18 J/g°C
5
Endothermic vs Exothermic
Endothermic
Exothermic
 Energy is __________
 Energy is __________
 Energy of the __________ is greater
than the energy of the ___________
 Energy of the __________ is
greater than the energy of the
___________
 H = positive or negative
 Beaker feels - hotter or colder
 Heat is written on the __________ of
the equation
Example:
S2Cl2 + CCl4 + 150kJ  CS2 + 3Cl2
 H = positive or negative
 Beaker feels - hotter or colder
 Heat is written on the __________
of the equation
Example:
2Na + 2H2O  2NaOH + H2 +150kJ
ENERGY AND PHASE CHANGES
Directions: Do the words describe: an endothermic reaction or
an exothermic reaction.
1. Energy of reactants > energy of products.____________________
2. Energy of products > energy of reactants.____________________
3. ∆H is a negative (ex) ∆H= -100kj.__________________________
4. Temperature inside the reaction vessel increases._____________
5. N2 (g) + 3H2 (g)  2NH3(g) + 92.4 kj _____________________
6. ∆H is a positive (ex) ∆H= +100kj.__________________________
7. 2NH3(g) + 92.4 kj  N2 (g) + 3H2 (g) _____________________
8. Energy of reactants < energy of products.____________________
9. Energy of products < energy of reactants.____________________
10.
Energy is written on the side of the reactants.___________
11.
Energy is written on the side of the products.___________
6
HEATING CURVE FOR WATER - www.youtube.com/watch?v=qlizkLCSXmQ and
https://youtu.be/nZVMqYU6Yfg
T
E
M
P
E
R
A
T
U
R
E
Temperature remains constant during a CHANGE OF STATE
Boiling 
condensing 
100 -
Vapor warming
or cooling
(ºC)
Liquid
warming or
cooling
Melting 
freezing 
0
-solid warming or
cooling
Time
Time
(min)
0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
Temperatur
e (°C)
-6
-2
0
0
0
0
4
13
8.5
9.0
9.5
10.0
10.5
11.0
11.5
12.0
12.5
13.0
97
100
100
100
100
100
100
100
105
109
What is the temperature
during the entire melting
or freezing
process?_______
What type of energy is
involved during this
process? PE or KE
What is the temperature
during the entire
vaporizing or condensing
process?_______
What type of energy is
involved during this
process? PE or KE
7
http://youtu.be/sqkS9_MxRVU and https://youtu.be/0cUK4jcAEaU
Match the “PHASE CHANGE” with the correct description.
Up and Right – endo down and left - exo
A. absorb heat (endothermic – heat is +)
B. release heat(exothermic – heat is -)
____________ Freezing
____________ Melting
____________ Boiling / Vaporize
____________ Condensing
____________ Subliming (sg)
Answer the following questions using the graph above.
1. What is the freezing point of the substance? _____
2. What is the boiling point of the substance? _____
3. What is the melting point of the substance? _____
4. What letter represents the range where the vapor is being warmed? _____
5. What letter represents the range where the liquid is being warmed? _____
6. What letter represents the range where the solid is being warmed? _____
7. What letter represents the vaporization of the liquid? _____
8. What letter represents the melting of the solid? _____
9. What letter represents condensation? _____
8
Draw a heating curve for water going from -20 ºC to 125 ºC on the axis below. Write in all formulas used to
calculate heat. Also write KE for a kinetic energy change and PE for a potential energy change.
125
PE change
100
KE change
KE change
PE change
0
KE change
For diagonals – the formula will be q=mcΔT
NOTE: “c” is often different for the solid form, liquid form and gaseous form for the same
substance.
For plateaus (flat areas) – the formula will be …..
a. lower plateau q=mΔHfusion
b. upper plateau q=mΔHvaporization
9
The data below is for water (H2O)
use highlighters and copy this page from the document labeled “highlighted section” from my teacher page
Melting
point
Boiling
Point
Heat of
Fusion
Heat of
SH Capacity SH Capacity SH Capacity
Vaporization (solid)
(liquid)
(vapor)
0.0 C
100.0 C
334 J/g
2260 J/g
2.05 J/gºC
4.18 J/gºC
1.90/gºC
e. Draw a heating curve for water going from -20 ºC to 125 ºC on the axis below. Write in all formulas used to
calculate heat.
f. Determine the amount of heat required to convert 15.0 g of water at 100.0 ºC to steam at 100.0 ºC.
ΔH = positive or negative
ΔH (q) = _____
g. Determine the amount of heat required to convert 15.0 g of ice at 0.0 ºC to liquid at 0.0 º C.
ΔH = positive or negative
ΔH (q) = _____
h. Determine the amount of heat given off when 25 g of water cools from 56.2 ºC to 33.5 ºC.
i. Determine the amount of heat needed to raise the temperature of 19.6g of X from -15 ºC to -2.0 ºC.
10
MORE HEAT PRACTICE
The data below refers to an unknown substance, X.
Melting
Boiling
Heat of
Heat of
SH Capacity SH
point
Point
Fusion
Vaporization (solid)
(liquid)
32.0C
112.0C
425 kJ/g
695 kJ/g
2.3 J/gºC
5.9 J/gºC
SH (vapor)
1.1 J/gºC
a. Draw a heating curve for substance X going from 15 ºC to 125 ºC on the axis below. Write in all formulas
used to calculate heat.
b. Determine the amount of heat released when 15.0 g of gaseous X at 112.0 ºC changes to liquid at 112.0 ºC.
ΔH = positive or negative
ΔH (q) = _____
c. Determine the amount of heat required to convert 15.0 g of solid X at 32.0 ºC to liquid at 32.0 º C.
ΔH = positive or negative
ΔH (q) = _____
d. Determine the mass of X that can be heated from 38 °C to 102 °C using 4500 J of heat.
e. Find the final temperature if 150 g of X at 19°C receives 650 J of heat.
11
MULTI-STEP HEAT PROBLEMS
The data below refers to an unknown substance, X.
Melting
Boiling
Heat of
Heat of
SH Capacity SH
point
Point
Fusion
Vaporization (solid)
(liquid)
14.0C
86.0C
150 J/g
550 J/g
4.3 J/gºC
5.2 J/gºC
SH (vapor)
1.1 J/gºC
a. Draw a heating curve for substance X going from 2 ºC to 120 ºC on the axis below. Write in all formulas
used to calculate heat.
b. Determine the amount of heat necessary to change 30 g of X from solid at 4 °C to liquid at 14 °C.
STEP 1:
STEP 2:
TOTAL:
c. How much heat will be released when 25 g of X cools from 110 °C to 50 °C?
STEP 1:
STEP 2:
STEP 3:
TOTAL:
12
THERMODYNAMICS: ENERGY & ENTROPY IN CHEMICAL REACTIONS
Enthalpy of a Reaction (ΔH):
Entropy (ΔS): https://youtu.be/MALZTPsHSoo -
STOP and THINK: Which process(es) below would have a positive change (increase) in entropy (ΔS + )?
a. cleaning your room
c. a bottle breaking
b. leaves falling off of a tree
d. water freezing
Predicting Changes in Entropy:
1.
Temperature: When a cold glass of water heats up, the particles move faster and the water particles
become more disorganized. So,
 Entropy increases (ΔS is +) with increasing temperature (increasing average KE).
 Entropy decreases (ΔS is - ) with decreasing temperature (decreasing average KE).
2.
Change in State of Matter: When an ice cube melts, it changes state (solidliquidgas). The
particles of solid ice are in fixed positions whereas the particles of liquid water slide over each other,
and the particles of water vapor move in all directions randomly. So,
Entropy increases (ΔS is +) when a substance changes state from solid  liquid  gas
3.
Solutions: For example, when a solid ionic compound dissociates into ions in an aqueous solution, the
system becomes more disorganized. For example:
NaCl (s)  Na+(aq) + Cl- (aq).
So,
Entropy increases (ΔS is +) when an ionic solid  aqueous solution
However, entropy decreases when a gas dissolves in a solvent. The gas molecules do not move around
as easily in the solution.
CO2 (g)  CO2 (aq)
Entropy decreases (ΔS is - ) when gas  solution
4.
Assuming no change in physical state, the entropy of a system increases when the number of gaseous
product particles is greater than the number of gaseous reactant particles. Gases move in a random
chaotic manner and the more gas particles there are, the more random.
2 SO3(g)  2 SO2(g) + O2(g)
Entropy increases (ΔS is +) when: the number of gas particles in reactants is less than the
number of gas particles in the products.
13
STOP and THINK:
Which reactions below would have a positive change in entropy (ΔS + )?
a.
Al2(CO3)3(s)  Al2O3(s) + 3CO2(g)
b.
KCl(aq) + Br2(l)  KBr(aq) + Cl2(g)
c.
FeCl2(aq) + H2SO4(aq)  FeSO4(s) + HCl(aq)
How can you tell if a reaction will be spontaneous? (able to occur?)
https://youtu.be/huKBuShAa1w
ΔS
change in entropy
+ more disorder
ΔH
-
+ more disorder
+
-
less disorder/more order
-
- less disorder/more order
+
change in heat
exothermic
releases heat
endothermic
requires heat
exothermic
releases heat
endothermic
requires heat
Spontaneous???
Yes – it can occur at any temp.
Spontaneous at high temperature
Spontaneous at low temperature
No – it cannot occur
So, in order to determine if a reaction will be spontaneous, there are 2 things to consider:
1. Sign of ΔS.
2. Sign of ΔH.
Determine which reaction(s) below will naturally occur (be spontaneous)? Which will require added energy?
Which will only occur naturally at low temperatures?
a.
Na2CO3 (s)

ΔS = _____
Na2O (s) + CO2 (g)
ΔH = 321kJ
Spontaneous? ____________
b.
Mg (s) + 2 HCl (aq)  MgCl2 (aq) + H2 (g)
ΔS = _____ ΔH = - 457.6kJ
Spontaneous? _____________
c.
N2 (g) + 3 H2 (g)

2 NH3
(g)
ΔS = _____
ΔH= -91.8kJ
Spontaneous? _____________
d.
2 KCl (s) + 3 O2 (g)  2 KClO3
(s)
ΔS =_____
ΔH = 90.0kJ
Spontaneous? _____________
14
Thermochemistry – Use number 1 as an example for number 2.
1. Consider the following reaction:
2S + 3 O2  2SO3 + 791.4kJ
a. Is the reaction endothermic or exothermic?_exothrmic since kJ (heat) is in the products____
b. Is heat absorbed or released?___released (exo)______
c. The ΔH = ___-791.4kJ______
d. If 73.4g of S react, find the kJ value.
73.4gS x 1 mole S x 791.4kJ = 1810kJ of heat released.
32.07gS 2 mole S
e. If 1063kJ is measured, find the mass of oxygen consumed.
2S + 3 O2  2SO3 + 791.4kJ
1063kJ x 3 mole O2 x 32.00g O2 = 128.9g O2
791.4kJ 1 mole O2
f. Draw a sketch of the enthalpy diagram
Exothermic diagram (enthalpy of products is lower than that of the reactants)
2. Consider the following reaction:
N2 + O2  2NO
ΔH = 180 kJ
a. Is the reaction endothermic or exothermic?___________________________
b. Is heat absorbed or released?___________________
c. Write the equation showing the heat term in the proper location:___________________
d. If 106.7g of NO are produced, find the kJ value.
e. If 454kJ is measured, find the mass of oxygen consumed.
f. Draw a sketch of the enthalpy diagram
15
HOMEWORK #1 Specific Heat Worksheet
Show all work and proper units.
1. A 15.75-g piece of iron absorbs 1086.75 joules of heat energy, and its temperature changes from 25 °C to
175 °C. Calculate the specific heat capacity of iron.
1. How many joules of heat are needed to raise the temperature of 10.0 g of aluminum from 22 °C to 55 °C, if
the specific heat of aluminum is 0.90 J/g°C?
2. To what temperature will a 50.0 g piece of glass raise if it absorbs 5275 joules of heat and its specific heat
capacity is 0.50 J/g°C? The initial temperature of the glass is 20.0 °C.
3. Calculate the heat capacity of a piece of wood if 1500.0 g of the wood absorbs 6.75×104 joules of heat, and
its temperature changes from 32 °C to 57 °C.
4. 100.0 mL of 4.0 °C water is heated until its temperature is 37 °C. If the specific heat of water is 4.18 J/g°C,
and the density of water is 1.0 g/mL. Calculate the amount of heat energy needed to cause this rise in
temperature.
5. 25.0 g of mercury is heated from 25 °C to 155 °C, and absorbs 455 joules of heat in the process. Calculate
the specific heat capacity of mercury.
16
HOMEWORK #2 CALORIMETRY PRACTICE
METAL
mass
specific heat capacity
Tf
Ti
WATER
4.18 J/g°C
1. A piece of metal with a mass of 38.0 g and a temperature of 115 °C is placed into 75.0 mL of water at 15
°C. The final temperature of the system is 20 °C. The density of water is 1.0 g/mL. What is the specific
heat of the metal?
2. A piece of gold (c= 0.129 J/g°C) is heated to 100 °C and placed into 60 mL of water at 20.5 °C. The final
temperature of the system is 22 °C. What was the mass of the metal?
3. 25.0 g of copper (c=.389 J/g°C) at 95 °C is placed into water at 25 °C. The final temperature of the system
is 28.5 °C. What was the mass of the water?
4. A piece of aluminum (c=0.900 J/g°C) with a mass of 55.0 g is heated and placed into 45.0 mL of water at 18
°C. The final temperature of the system is 22 °C. What was the starting temperature of the metal?
17
HOMEWORK #3: HEAT PRACTICE
The data below refers to an unknown substance, X.
Melting
Boiling
Heat of
Heat of
SH Capacity SH Capacity SH Capacity
point
Point
Fusion
Vaporization (solid)
(liquid)
(vapor)
-12.C
76.0C
65 kJ/g
88 kJ/g
5.2 J/gºC
2.1 J/gºC
1.7 J/gºC
a. Draw a heating curve for substance X going from -20 ºC to 125 ºC on the axis below. Write in all formulas
used to calculate heat.
b. Determine the amount of heat released when 24.0 g of liquid X at -12.0 ºC changes to solid at -12.0 ºC.
ΔH = positive or negative
ΔH (q) = _____
ΔS =
c. What mass of X can be converted from liquid to vapor at 76.0 ºC using 425 kJ of heat?
mass =
ΔH = _____
ΔS = _____
d. Determine the amount of heat given off when 25 g of X cools from 96.2 ºC to 83.1 ºC.
e. Find the final temperature if 17 g of X at 115 °C absorbs 725 J of heat.
18
HW #4: ENTROPY AND SPONTANEITY OF CHEMICAL REACTIONS
1. Determine whether ΔS will be positive (disorder increases) or negative (disorder decreases) for
each process below:
a. burning paper
_____________________
b. organizing your binder
_____________________
c. water freezing
_____________________
d. dry ice sublimating
_____________________
2. Determine the sign of ΔS for each reaction below. Then determine under which conditions (if at all)
the reactions will be spontaneous.
a. CO2 (g) →
C (graphite) + O2
ΔH = +393.5 kJ/mol
ΔS = _________
b. 2 HCl (g)
→ H2 (g) + Cl2 (g)
spontaneous? _____________
ΔH = +184.6 kJ/mol
ΔS = _________
c. CH4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (l)
spontaneous?_____________
ΔH = -890.4 kJ/mol
ΔS = _________
spontaneous? _____________
d. 2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g) ΔH = -446 kJ/mol
ΔS = _________
spontaneous?____________
3. A student tries to dissolve a salt (ionic compound) in water. He is having a hard time, so he decides to
use the bunsen burner to heat up the water as the salt dissolves. It works! Determine the signs for
ΔH and ΔS for this process.
ΔH ______________
ΔS _______________
19
HW #5: ENTROPY AND SPONTANEITY OF CHEMICAL REACTIONS
1. ΔHo for the reaction is -167 kJ and ΔSo for the reaction is 88 J/K.
This reaction will be:
a) spontaneous at all temperatures
b) spontaneous only at high temperature
c) spontaneous only at low temperature
d) nonspontaneous at all temperatures
2. ΔHo for the reaction is 488 kJ and ΔSo for the reaction is 88 J/K.
This reaction will be:
a) spontaneous at all temperatures
b) spontaneous only at high temperature
c) spontaneous only at low temperature
d) nonspontaneous at all temperatures
3. ΔHo for the reaction is -167 kJ and ΔSo for the reaction is -123 J/K.
This reaction will be:
a) spontaneous at all temperatures
b) spontaneous only at high temperature
c) spontaneous only at low temperature
d) nonspontaneous at all temperatures
4. ΔHo for the reaction is 90 kJ and ΔSo for the reaction is 330 J/K.
This reaction will be:
a) spontaneous at all temperatures
b) spontaneous only at high temperature
c) spontaneous only at low temperature
d) nonspontaneous at all temperatures
5. In exothermic reactions, heat is ________________________.The temperature of the surroundings
__________________________ so the temperature of the system ________________________________.
6. In endothermic reactions, heat is ________________________.The temperature of the surroundings
__________________________ so the temperature of the system ________________________________.
7. All combustion reactions are __________________________________.
(endo or exo) (pos or neg ΔH)
8. Describe how entropy(ΔS) changes (positive or negative) when:
a. temperature increases________________
b. temperature decreases________________
c. A gas dissolves in water_______________
d. A solid dissolves in water______________
e. 2 gas particles turn into 1 gas particle_______
f. A bunch of different gases mix together_______
9. If you know the sign of the ΔH AND you know the sign of the ΔS, then you can determine the sign of the
_____________if you know either the temperature or the always/never situation. (think of the table).
20
HW #6: GENERAL REVIEW OF ENDOTHERMIC AND EXOTHERMIC REACTIONS
1.
Exothermic reactions _____________________ heat.
(absorb or release)
2. Endothermic reactions _____________________ heat.
(absorb or release)
3. What is the sign of the ∆H for an endothermic reaction?_____________________
(positive or negative)
4. What is the sign of the ∆H for an exothermic reaction?_____________________
(positive or negative)
5. Draw an energy diagram for an exothermic reaction
6. Draw an energy diagram for an endothermic reaction
7. Heat always flows from the ___________________object to the ______________object.
8. If an exothermic reaction takes place in a beaker, does the beaker get hot or cold after the reaction
occurs?__________________________________
9. If an endothermic reaction takes place in a beaker, does the beaker get hot or cold after the reaction
occurs?__________________________________
21
10. Water has a high specific heat capacity (c). What does this mean in terms of ease of heating it up or cooling
it down?
11. Label the Phase Change Diagram below
12. Going up the phase diagram chart (and right) is _________________________________.
(endothermic or exothermic)
13. Going down the phase diagram chart (and left) is _________________________________.
(endothermic or exothermic)
14. Which would hurt more (cause more damage): putting your hand in 100°C water or in 100°C steam?
Why?
22