General Overlay: Topic 8, 9, 10, 11, 12, 13
Topic 8: Heat and Work
Video A: Energy In Reactions
● Chemical reactions are based on changes in ENERGY as bonds are broken and formed.
● Energy cannot be created or destroyed and can be transferred from one body to
another.
● Exothermic: release energy during a chemical reaction to form bonds
○ Energy transferred out of the system to the surroundings.
○ Potential energy decreases
○ Surroundings gain energy
○ Heat is a product
● Endothermic: gain energy during a chemical reaction to break bonds
○ Energy is transferred into the system
○ Potential energy increases
○ Surroundings lose energy
○ Heat is a reactant
● Heat of a chemical reaction (q) is due to the change in potential energy or in other words
the making and breaking of bonds.
Heat = q = ΔH (change in enthalpy)
● Enthalpy: how much energy is released/absorbed within a reaction.
ΔH ﹤0 = Exothermic
ΔH ﹥0 = Endothermic
● The net ΔHrxn depends on the strength of the bonds that are broken and the bonds that
are formed.
ΔHrxn= ΔHformed + ΔHbroken
○ For most reactions we cannot predict ΔHrxn by simply looking at the reaction.
Exothermic: Endothermic:
|ΔHformed| ﹥ ΔHbroken |ΔHformed| ﹤ΔHbroken
Video B: First Law of Thermodynamics
● The only way to make a system experience change in its energy is through the process
of energy transfer to or from the surroundings as either heat or work.
ΔEsystem = KEAVE + PEAVE
PEAVE = interactions between molecules, intermolecular forces and bonding interactions
KEAVE = motion
● Internal energy of the system is the sum of the average kinetic energy and the average
potential energy of the molecules in the system.
● For the Universe: ΔE = 0
, ● There are two different ways to change the energy of the system
○ Heat: random motion of particles in multiple directions
○ Work: motion against an opposing force and causes matter to move
● Work is creating a force acting over a distance.
● Heat is when energy is lost/dissipated and there is no work.
● Work and heat represents the transfer of energy.
ΔE = q+w
● We track ENERGY transfer between a system and its surroundings
Work Heat
Work done by the surroundings Heat into the system
W﹥0 Q﹥0
Energy of the system increases Energy of the system increases
Work done by the system Heat out of the system
W﹤0 Q﹤0
Energy of the system decreases Energy of the system decreases
Video C: Pressure-Work Volume
Review: Ideal Gas Postulates
1) Gas molecules have no volume of their own.
2) Gas molecules are perfectly elastic collisions and are in constant linear motion.
3) Gas molecules experience no intermolecular forces acting between one another.
4) The Kinetic Energy of an ideal gas is directly proportional to its temperature measured in
degrees K.
Equations: PV = nRT
R = 0.082 L-atm/mol-K
Internal Energy to Temperature
KE = (3/2) nRT
KEMolar = (3/2) RT
Any collection of gases at same temperature have the same KE molar.
● Work due to the expansion or compression of gas (PV work) is equal to the negative
external pressure of gas (Pexternal) times the change in volume of gas (ΔV).
WPV = - (Pexternal) ((ΔV)
, ● For any gas compression: ΔV ﹤0
○ When a gas expands (more moles of gas on product side) its volume increases.
It pushes against an external pressure and its energy goes down.
○ Work is being done to the gas; work must be positive (w > 0).
● For any gas expansion: ΔV﹥0
○ When a gas is compressed (less moles of gas on the product side) its volume
decreases. An external force pushes on the gas and its energy goes up.
○ Work is being done by the gas; work must be negative (w < 0).
Video D: Calorimetry Calculation (1 system problem)
● Calorimetry: measures heat of a reaction at constant temperature in ℃.
○ Heat: process for the transfer of energy
○ Temperature: internal kinetic energy of the system.
● Specific Heat Capacity: energy required to raise the temperature of 1 gram of a
substance by 1 ℃ or in other words, how much energy is required to change the
temperature of a substance.
C = J/(g x deg) or Cal/(g x deg)
Equation for Calorimetry:
Q = mcΔT
T = temperature
m = mass of a substance
c = specific heat capacity
● Substances with small heat capacity values heat up very quickly (e.g metals)
● Substances with high heat capacity values will need a lot of heat to raise the
temperature of that substance.
● The specific heat capacity of water is 4.184 J/(g・deg)
● When energy is absorbed by the system as heat, q is positive.
● When energy is released by the system as heat, q is negative.
HOW TO APPROACH 1-SYSTEM CALORIMETRY CALCULATION
● List out the variables given, m (mass), c (specific heat), ΔT ( Tf - Ti), q (heat in Joules)
● Isolate the variable that is being asked for and solve.
Sample Problem:
1. If I have 32 grams of unknown metal ( c= 3.68 J/(℃-g) ) at 25 ℃ and I transfer 120 of
energy to the metal, what is the new temperature?
m= 32 grams
c= 3.68 J/(℃-g)
ΔT= ?
Topic 8: Heat and Work
Video A: Energy In Reactions
● Chemical reactions are based on changes in ENERGY as bonds are broken and formed.
● Energy cannot be created or destroyed and can be transferred from one body to
another.
● Exothermic: release energy during a chemical reaction to form bonds
○ Energy transferred out of the system to the surroundings.
○ Potential energy decreases
○ Surroundings gain energy
○ Heat is a product
● Endothermic: gain energy during a chemical reaction to break bonds
○ Energy is transferred into the system
○ Potential energy increases
○ Surroundings lose energy
○ Heat is a reactant
● Heat of a chemical reaction (q) is due to the change in potential energy or in other words
the making and breaking of bonds.
Heat = q = ΔH (change in enthalpy)
● Enthalpy: how much energy is released/absorbed within a reaction.
ΔH ﹤0 = Exothermic
ΔH ﹥0 = Endothermic
● The net ΔHrxn depends on the strength of the bonds that are broken and the bonds that
are formed.
ΔHrxn= ΔHformed + ΔHbroken
○ For most reactions we cannot predict ΔHrxn by simply looking at the reaction.
Exothermic: Endothermic:
|ΔHformed| ﹥ ΔHbroken |ΔHformed| ﹤ΔHbroken
Video B: First Law of Thermodynamics
● The only way to make a system experience change in its energy is through the process
of energy transfer to or from the surroundings as either heat or work.
ΔEsystem = KEAVE + PEAVE
PEAVE = interactions between molecules, intermolecular forces and bonding interactions
KEAVE = motion
● Internal energy of the system is the sum of the average kinetic energy and the average
potential energy of the molecules in the system.
● For the Universe: ΔE = 0
, ● There are two different ways to change the energy of the system
○ Heat: random motion of particles in multiple directions
○ Work: motion against an opposing force and causes matter to move
● Work is creating a force acting over a distance.
● Heat is when energy is lost/dissipated and there is no work.
● Work and heat represents the transfer of energy.
ΔE = q+w
● We track ENERGY transfer between a system and its surroundings
Work Heat
Work done by the surroundings Heat into the system
W﹥0 Q﹥0
Energy of the system increases Energy of the system increases
Work done by the system Heat out of the system
W﹤0 Q﹤0
Energy of the system decreases Energy of the system decreases
Video C: Pressure-Work Volume
Review: Ideal Gas Postulates
1) Gas molecules have no volume of their own.
2) Gas molecules are perfectly elastic collisions and are in constant linear motion.
3) Gas molecules experience no intermolecular forces acting between one another.
4) The Kinetic Energy of an ideal gas is directly proportional to its temperature measured in
degrees K.
Equations: PV = nRT
R = 0.082 L-atm/mol-K
Internal Energy to Temperature
KE = (3/2) nRT
KEMolar = (3/2) RT
Any collection of gases at same temperature have the same KE molar.
● Work due to the expansion or compression of gas (PV work) is equal to the negative
external pressure of gas (Pexternal) times the change in volume of gas (ΔV).
WPV = - (Pexternal) ((ΔV)
, ● For any gas compression: ΔV ﹤0
○ When a gas expands (more moles of gas on product side) its volume increases.
It pushes against an external pressure and its energy goes down.
○ Work is being done to the gas; work must be positive (w > 0).
● For any gas expansion: ΔV﹥0
○ When a gas is compressed (less moles of gas on the product side) its volume
decreases. An external force pushes on the gas and its energy goes up.
○ Work is being done by the gas; work must be negative (w < 0).
Video D: Calorimetry Calculation (1 system problem)
● Calorimetry: measures heat of a reaction at constant temperature in ℃.
○ Heat: process for the transfer of energy
○ Temperature: internal kinetic energy of the system.
● Specific Heat Capacity: energy required to raise the temperature of 1 gram of a
substance by 1 ℃ or in other words, how much energy is required to change the
temperature of a substance.
C = J/(g x deg) or Cal/(g x deg)
Equation for Calorimetry:
Q = mcΔT
T = temperature
m = mass of a substance
c = specific heat capacity
● Substances with small heat capacity values heat up very quickly (e.g metals)
● Substances with high heat capacity values will need a lot of heat to raise the
temperature of that substance.
● The specific heat capacity of water is 4.184 J/(g・deg)
● When energy is absorbed by the system as heat, q is positive.
● When energy is released by the system as heat, q is negative.
HOW TO APPROACH 1-SYSTEM CALORIMETRY CALCULATION
● List out the variables given, m (mass), c (specific heat), ΔT ( Tf - Ti), q (heat in Joules)
● Isolate the variable that is being asked for and solve.
Sample Problem:
1. If I have 32 grams of unknown metal ( c= 3.68 J/(℃-g) ) at 25 ℃ and I transfer 120 of
energy to the metal, what is the new temperature?
m= 32 grams
c= 3.68 J/(℃-g)
ΔT= ?
