Thermodynamics
New section in table of contents
Energy
 Energy Defined: Matter is stuff (atoms), Energy
is the stuff that makes matter do stuff.
 What are some different types of energy you
know of?
First Law of Thermodynamics
 KC 1: 1st Law of Thermodynamics - energy can
be neither created nor destroyed (conservation
of energy).
Mechanical Energy
 Energy due to a object’s motion (kinetic) or
position (potential).
Thermal Energy
Heat Energy
The heat energy of an object determines
how active its atoms are.
A hot object is one whose atoms and
molecules are excited and show rapid
movement.
 A cooler object's molecules and atoms
will show less movement.
Electromagnetic Energy
Light energy
 Includes energy
from gamma
rays, x-rays,
ultraviolet rays,
visible light,
infrared rays,
microwave and
radio bands
Electrical Energy
 Energy caused by the movement of electrons.
 Easily transported through power lines and
converted into other forms of energy.
Nuclear Energy
 The energy released during nuclear fission or fusion.
Fusion: Stars
Fission: Nuclear
Power Plants; Earth Core
Chemical Energy
 Energy that is available for release from
chemical reactions.
Energy
Matter
Energy
 Anything that has mass
and takes up space
 Energy is of central
importance in physics.
 Anything made of
atoms or subatomic
particles.
 It is impossible to give
a comprehensive
definition of energy
because of the many
forms it may take.
 KC 2: Energy is the
ability to do work or
produce heat.
Energy
 KC 3: Chemical potential energy is energy
stored in a substance because of its
composition
 Chemical potential energy is important in
chemical reactions
Chemical Energy
Temperature
 KC 4: Temperature is a quantity proportional to
molecular kinetic energy.
 Temperature is the measure of the movement
of atoms.
 Average kinetic energy of molecules is
measured in terms of temperature.
Rank the following pictures according to
their average kinetic energy (low to high)
A
B
C
Heat
KC 5: Heat is a
form of energy;
that is in the
process of
flowing from a
warmer object
to a cooler
object.
Chemical
Energy
Demos
2 more things to do
 Ticket out the door- Post-it Note with name
 Write one complete sentence that shows how
energy, temperature, and heat are related.
 Closure- Exit Pass- The system with the most
heat will burn you the worst (200mL boiling
water). Write the question below and think!
Why can dry ice, solid CO2 burn you, but CO2
in the gas phase cannot? Gases have more
kinetic energy then liquids or solids.
Cold and Hot Packs
 How do you think instant hot and cold packs
work? Keep in mind chemical reactions and
energy.
Cold or Hot Pack?
 Experiment A:
A spoon full of sodium bicarbonate (NaHCO3)
with 2 squirts of water in a sandwich bag.
 Experiment B:
A spoon full of calcium chloride (CaCl2) with 2
squirts of water in a sandwich bag.
KC 6: Exothermic Reaction- is a
chemical reaction that releases energy
in the form of light or heat.
Activation
Energy
Energy of reactants
•Energy
Energy of products
Reactants
-DH
Products
•Reaction Progress
KC 7: Endothermic Reaction-
chemical reaction accompanied by the
absorption of heat.
Energy of reactants
Energy of products
Energy
Activation
Products
Energy
Reactants
+DH
Reaction progress
Complete the sentences below for both
experiment using the correct word.
1. Experiment ____ is similar to ____(hot /cold
pack)
2. The salt which dissolves in water is_____.
3. The heat energy is _____(given off/ taken in).
4. The temperature _____ (falls/rises).
5. The reactions is _____
(endothermic/exothermic).
Chemical Energy and the Universe
 Chemists are interested in changes in energy
during reactions.
 Enthalpy is a measure of the total energy in a
chemicals bonds.
 KC 8: Enthalpy= heat of reaction= ΔH is the
change in enthalpy during a reaction
symbolized as ΔHrxn.
ΔHrxn = Hfinal – Hinitial
ΔHrxn = Hproducts – Hreactants
KC 9: Endothermic Reaction = positive ∆H
KC 10: Exothermic Reaction = negative ∆H
Enthalpy of a Reaction
KC 11: ΔHrxn = ∑n∆Hproducts – ∑n∆Hreactants
Example
 KC 12: Use the table to find the
∆Hrxn for the following reaction:
SO2(g) + NO2(g) → SO3(g) + NO(g)
Substance
∆Hf (kJ/mol)
SO2(g)
-296.8
NO2(g)
33.1
SO3(g)
-395.8
NO(g)
90.3
Enthalpy of Reaction
 There will NOT be values given for substances
that are in their “natural” state
 Example – oxygen: It does not take energy to
make O2(g)
carbon: It does not take energy to
make C(s)
 These heats of formation (∆Hf) will be zero
 You MUST know this!!
Specific Heat
 KC 13: The specific heat of
any substance is the amount
of heat required to raise one
gram of that substance one
degree Celsius.
 Some objects require more
heat than others to raise their
temperature.
Specific Heat
 Intrinsic Property
 It is independent of how much of the material is present.
Measuring Heat
Measuring Heat
 KC 14: A calorie is defined as the amount of
energy required to raise the temperature of
one gram of water one degree Celsius.
 Food is measured in Calories, or 1000 calories
(kilocalorie).
 KC 15: A joule is the SI unit of heat and energy,
equivalent to 0.2390 calories.
Specific Heat of Water
 Life on earth evolved in water, and all life still
depends on water.
 KC 16: Water has a specific heat capacity of 4.2
J/ goC, which means that it takes 4.2 joules of
energy to heat 1g of water by 1oC.
 KC 17: This is unusually high and it means that
water does not change temperature easily.
KC 18: Calculating Heat
q = c × m × ΔT
 q = heat absorbed or released
 c = specific heat of substance
 m = mass of substance in grams
 ΔT = change in temperature in Celsius (Final –
Initial)
Calculating Heat
 KC 19: A 12.5g sample of aluminum was
originally at 20°C. It was heated to a
temperature of 34.3°C. How much heat was
absorbed by the sample? (specific heat of
aluminum = 0.9J/g°C)
Calculating Heat
 KC 20: A sample of zinc was heated from an
initial temperature of 234°C to 439°C. There
was 198.5J of heat absorbed and the specific
heat of zinc is 0.388J/g°C. What was the mass
of the sample?
Calculating Heat
 KC 21: Calculate the amount of heat needed
in calories to raise 23g of water from 49°C to
236°C.
Closure – Exit Pass
 Why did one ice cube melt faster than the
other?
 You must use the correct vocabulary in your
response.
Calorimetry
 How do scientists measure the heat of
reactions?
 Scientists use a calorimeter to measure the
temperature change of a given reaction
 They can then use q=mc∆T to calculate the
heat of that reaction
Calorimetry
 KC 22: Calorimetry – using temperature change
to calculate the heat of a reaction as it moves
from reactants to products
Calorimetry
 This process of calorimetry is how scientists
measure the amount of calories in food
 You will be doing a similar process to rank
different snack foods from smallest kcal/g to
largest amount of kcal/g
 Video
Calorimetry Example
 Calculate the number of calories absorbed by the
water, based on the information below.
Mass of empty beaker: 300g
Temperature of water before reaction: 27.0 ˚C
Mass of beaker with H2O: 465g
Temperature of water after reaction: 65.0 ˚C
Lab – Burning Food
 Measure approximately 15mL of water into the can. Record
the mass of the water in data table (NOTE: 1mL of water = 1g)
 Measure and record the initial temperature of the water in the
can then remove the thermometer.
 Place the food sample on the paper clip and record its mass in
the data table.
 Quickly ignite the food sample with a Bunsen burner and place
the platform under the test tube. Place the thermometer in the
can. Allow the food to burn completely.
 Measure and record the highest temperature achieved by the
water in the can during the burning of the food.
 Record the mass of the food sample after burning and record
in data table.
Data Table
Food Sample:
Mass of water
Initial temp of water
Initial mass of food sample
(this also includes the paper clip and
cork)
Final temp of water
Final mass of food sample
(this also includes the paper clip and
cork)
Change in temp of water
Change in mass of food
sample
q of water (from
calculation)
Kcal/g of food sample
(from calculation)
Peanuts
Doritos
Cheetos
Pretzels
Hess’s Law
 KC 23: Hess’s law – the amount of heat
released or absorbed in a chemical reaction
does not depend on the number of steps in the
reaction
 The ∆Hrxn can be determined from using other
reactions
 This is like systems of equations:
2x + 3y = 23
6y + x = -45
Hess’s Law
RULES:
Flip an equation, change the sign on the ∆H
Multiply by a factor, multiply the ∆H by the
same factor (can be a whole number or a
fraction, ½, ¼ )
Add equations down to get the goal
reaction, add ∆H’s down to get the ∆H for
the goal reaction
KC 24: Example Hess’s Law
Reaction
∆H
2 C(s) + O2(g) → 2 CO(g)
-221kJ
C(s) + O2(g) → CO2(g)
-393kJ
C(s) + CO2(g) → 2CO(g)
?
Reaction
∆H
BRN – find ∆H
C2H5OH(l) + 2 O2(g) --> 2 CO2(g) +
2 H2O(l)
-875.
kJ
C (s) + O2 (g) --> CO2 (g)
-394.51
kJ
H2 (g) + ½ O2 (g) --> H2O (l)
-285.8
kJ
2H2(g) + 2C(s) + O2(g) --> C2H5OH(l)
?
Lab Reminders
 Keep the water boiling while you are doing the experiment
 Be careful when placing the metal in the test tube (I’m
going to give you 2 samples)
 Leave test tube in boiling water for 5-10min
 Use test tube holder to pour metal into calorimeter
 Don’t forget to record initial temperature of water
 Use about 15-20mL of water in the calorimeter
 Record highest temperature of water reached for final
temp
 Assume initial temp of metal is 100°C
Enthalpy
 KC 25: enthalpy is the measure of the energy
that is released or absorbed by the substance
when bonds are broken and formed during a
reaction
 When bonds are formed, energy is released
 When bonds are broken, energy is absorbed
 Exothermic reaction: -∆H
 Endothermic reaction: +∆H
Entropy
 KC 26: entropy (∆S) is a measure of the randomness
or disorder of the system; the greater the disorder of
a system, the greater its entropy
 As molecules or ions become more dispersed, their
disorder increases and their entropy increases
Reaction
Entropy change ∆S
NaCl(s) → Na+(aq) + Cl-(aq)
+43
2Na(s) + Cl2(g) → 2NaCl(s)
-181
Gibbs Free Energy
 KC 27: Gibbs free energy (∆G) is a measure of the
spontaneity of the process
 KC 28: A spontaneous reaction is one that does
occur or is likely to occur without continuous
outside assistance, such as input of energy
 A reaction is spontaneous if the Gibbs energy
change is negative
 If a reaction has a ∆G greater than 0, the reaction
is nonspontaneous
Gibbs Free Energy
 Reactions that have large negative ∆G values
often release energy (-∆H) and increase disorder
(+∆S)
 ∆G = ∆H - T∆S
 Example:
2 K(s) + 2 H2O(l) → 2K+(aq) + 2 OH-(aq) + H2(g)
∆H = -392kJ, ∆S = 0.047kJ/K, T = 298K
KC 29: Spontaneity
(thermodynamically favored)
∆H
-
+
_
+
∆S
+
+
-
T
∆G
Low
+
+
+
+
High
Low
High
Low
High
Low
High
Spontaneous?
Yes
No
No
Yes
Yes
No
Gibbs Free Energy
 KC 30: Given that the changes in enthalpy and
entropy are -139kJ and 277J/K respectively for the
reaction given below, calculate the change in
Gibbs energy. Then, state whether the reaction is
spontaneous at 25°C.
C6H12O6(aq) → 2 C2H5OH(aq) + 2 CO2(g)
Practice
 Which of the following will be true when a pure
substance in liquid phase freezes spontaneously?
A.∆G, ∆H, and ∆S are all positive
B.∆G, ∆H, and ∆S are all negative
C.∆G and ∆H are negative, but ∆S is positive
D.∆G and ∆S are negative, but, ∆H is
positive
Practice
 Which of the following pairs of conditions will
favor a spontaneous reaction?
A. A decrease in entropy and a decrease in enthalpy
B. A decrease in entropy and an increase in enthalpy
C. An increase in entropy and a decrease in enthalpy
D. An increase in entropy and an increase in enthalpy
Review Video - closure
 Entropy
 3, 2, 1
 3 things you learned
 2 things you found interesting
 1 question you have
Bellringer
 What three phases can H2O exist?
 Can liquid water that is boiling ever get hotter
than 100 C?
 What are the bubbles made of when water
boils?
Heating Curve
•temperature
Heating Curves
•gas
•liquid
•solid
•added energy
boiling/
cond. pt
melting/
freezing pt
•temperature
Heating Curves
•gas
•liquid
•solid
•added energy
Heating Curves
•boiling/condensing
•melting/freezing
melting/
freezing pt
•occurring here
•temperature
boiling/
cond. pt
•occurring here
•gas
•liquid
•solid
•added energy
 Cice = 2.09J/g°C
∆Hfus = 334J/g
 Cwater = 4.184J/g°C
∆Hvap = 2260J/g
 Csteam = 2.01J/g°C
q=mc∆T
q = m∆H
A 12g sample of ice initially at -5°C is heated to a
temperature of 95°C. Calculate the amount of heat
that is absorbed during this process.
Cice = 2.09J/g°C
∆Hfus = 334J/g
Cwater = 4.184J/g°C
∆Hvap = 2260J/g
Csteam = 2.01J/g°C
A 3.5g sample of water at 13.6°C is heated to a
temperature of 112°C. Calculate the amount of heat
absorbed during this process.
Cice = 2.09J/g°C
∆Hfus = 334J/g
Cwater = 4.184J/g°C
∆Hvap = 2260J/g
Csteam = 2.01J/g°C
•temperature
Cooling Curves
Removal of energy
Heating and Cooling Curve
Assignment
 You will be given 2 element cards to create a
heating and cooling curve. The melting points
and boiling points are on the card for each
element.
 You will need to draw a heating curve for one
and a cooling curve for the other.
Closure: Jeopardy
d
b B c
D
e
C
a
A
 Answer- This letter Indicated when the matter is a solid.
 Answer- 95°C
 Answer- 70°C