Chapter 16
Theories of Energy Changes
Section 16.1A Temperature change and Heat
THERMODYNAMICS - the study of energy and energy transfer
THERMOCHEMISTRY - the study of energy involved in chemical reactions
THE LAW OF CONSERVATION OF ENERGY = THE FIRST LAW OF
THERMODYNAMICS
- the total energy of the universe is constant
- energy is neither created nor destroyed; it only changes forms
SECOND LAW OF THERMODYNAMICS
- when energy is transformed some is changed into an unusable form (usually heat)
SYSTEM - the part of the universe being studied and observed (usually the reactants and products)
SURROUNDINGS - everything else in the universe (eg flask, lab desk, the students)
Universe = system + surroundings
△ E universe = △ E system + △ E surroundings = 0 (because energy is neither being created
nor destroyed so E universe = 0)
So, △ E system = - △ E surroundings
Systems are defined in 3 different ways:
1. OPEN system - the reaction is open to its surroundings Both energy and matter may be
exchanged between the system and its surroundings. eg open beaker.
2. CLOSED system - matter can not be transferred between the system and its surroundings
BUT energy can be transferred. eg stoppered flask
3. ISOLATED SYSTEM = the system is completely insulated from the surroundings. Neither energy nor
matter is exchanged eg thermos
Heat (q) vs. Temperature (page 628)
A. All energy is classified into 2 fundamental types:
1. KINETIC ENERGY = energy of motion
2. POTENTIAL ENERGY = energy that is stored
B. TEMPERATURE is defined as the average KE of the particles of a substance or a system
- the faster the molecules move the ____________ the temperature
- measured in 0C or K
- 0 K = absolute zero = all motion stops therefore no kinetic energy
- K = 0C + 273.15
- △T = Tf - Ti
Where f = final and i = initial
*a positive △T indicates and increase in temperature
*a negative △T indicates a decrease in temperature
C. HEAT = the transfer of KE between objects with different temperatures
- think of it as “energy in transit”
- represented be the letter, q
- unit = joule
- heat only moves from hot to cold (remember that heat can NOT be detected by senses; only changes in
heat can be detected by looking at temperature changes)
Since the first law of thermodynamics states that the energy released by a system is gained by its
surroundings, then:
q system = - q surroundings
Calculating Heat Transfer (Specific heat capacity problems) p.633
Heat transfer depends on three factors:
1. Mass of the substance, m
2. Change in temperature, T
3. Specific heat capacity, c
SPECIFIC HEAT CAPACITY = the quantity of energy required to change one gram of a specific
substance by one degree Celsius.
- unit = J/ g x 0C
- is a characteristic property of a pure substance (like melting pt, boiling pt and density) and depends on the
type of substance and the state of the substance
- a large specific heat capacity means the substance can release and absorb more energy than a substance with
a low specific heat capacity.
Specific Heat Capacity Problems
q=mxcx △T
1. A chunk of ice was added to 60.0 grams of liquid water. The initial temperature of the water
was 26.50C. The final temperature of the mixture was 9.70C. How much heat was lost by the water?
2. 100. grams of ethanol at 250C is heated until it reached 50.0C. How much heat does the ethanol gain?
3. A solid substance has a mass of 250.00 grams. It is cooled BY 25.00 0C and LOSES 4937.50 J
of heat. What is the specific heat capacity? Identify the substance.
Water has a high specific heat capacity compared to most substances due to its hydrogen bonds (strong intermolecular
bonds). This accounts for smaller variations in year-round temperatures in areas such as PEI which are surrounded by
water be able to explain why!!!
Water has a high specific heat capacity (due to strong intermolecular bonds) whereas metals have low specific heat
capacities. What does this mean for each?
Heat Capacity
HEAT CAPACITY, C = the quantity of energy required to change the temperature of the
substance one degree Celsius.
- unit = kJ/ 0C
-C=m. c
-so, q = C △T
C = heat capacity (kJ/ 0C)
m =mass (kg)
c = specific heat capacity (kJ/kg x 0C)
Heat Capacity Problems – to be assigned
Enthalpy Changes
Phase changes (ie changes in states of matter) involve changes in the potential energy of a
system (not the kinetic energy)
ENTHALPY CHANGE ( △H)
- the result of the change in potential energy of a system during a process such as a
chemical reaction or a physical change
- measured at constant pressure
- units are kJ/mol
- also called HEAT OF REACTION or even internal energy
Enthalpy Changes in Chemical Reactions
- The enthalpy change in a chemical reaction is the difference between the potential energy of
the products and the potential energy of the reactants.
- The potential energy changes result from chemical bonds being broken and formed
- Remember: chemical bonds are sources of potential energy BUT it still requires energy
to break a bond and creating a bond releases energy.
- Chemists use different subscripts to represent enthalpy changes for specific kinds of
reactions. For example △ H comb represents the enthalpy change of a combustion reaction
- Enthalpy of reaction is dependent on temperature and pressure so we use STANDARD
ENTHALPY OF REACTION =
△H0. It occurs at SATP (250C and 100 kPa)
- CATALYSTS do NOT have any effect on enthalpy
You can represent an EXOTHERMIC REACTION three ways:
1. As a thermochemical equation with the energy stated as a product
N2 + 3H2 –> 2 NH3 + 93 kJ
2. As a separate expression. H0 in an exothermic reaction is always negative.
N2 + 3H2 –> 2 NH3
△ H0 = -93 kJ/mol
3. As an enthalpy diagram.
You can represent an ENDOTHERMIC REACTION three ways:
1. As a thermochemical equation with the energy stated as a reactant.
MgCO3 + 117 kJ –> MgO + CO2
2. As a separate expression. △ H0 in an endothermic reaction is always positive.
MgCO3 –> MgO + CO2
△ H0= + 117 kJ/mol
3. As an enthalpy diagram.
** The enthalpy of reaction is linearly dependent on the unit of substances that react. So if yo u
multiply the coefficients in the thermochemical equation by any factor, you must also multiply
the enthalpy by the same factor.
Enthalpy Change Problems p.642-649
** all types use equation:
q = n∆H
Type 1 - Standard molar enthalpy of formation
Standard Molar Enthalpy of Formation ∆ H0f =
- the quantity of energy that is absorbed or released when one mole of a compound is
formed directly from its elements in their standard states
- when writing a formation equation:
1. Always write the elements in their standard states
2. Always show the formation of exactly one mole of the compound of interest
(may want to do practice problems here)
**Refer to Appendix E for Molar Enthalpies of Formation
1. How much heat is released when 50.00 grams of methane is formed from its elements?
2. How much heat is released when 0.534 grams of liquid water is formed from its elements?
Type 2- Standard molar enthalpy of combustion
*** the values in Table 16.3 are not accurate -get correct values from teacher
Standard Molar Enthalpy of Combustion ∆H0comb =
- the quantity of energy that is associated with the combustion of one mole of a given
substance.
1. How much heat is released when 50.00 grams of methane undergoes comp lete combustion?
2. Determine the heat released by the combustion of 57.68 grams of pentane.
Draw an Enthalpy Diagram:
Type 3 - Standard molar enthalpy and phase changes
phase changes (states of matter changes) always involve energy changes, but do not
involve a change in temperature. Therefore they are associated with potential energy
not kinetic energy.
note: although the temperature of the system remains constant, the temperature of the
surroundings often changes because of energy being absorbed or released as heat.
molar heat of vaporization, ∆ Hvap = the enthalpy change for the state change of one
mole of the substance from a liquid to a gas.
molar heat of condensation, ∆ Hcond = the enthalpy change for the state change of
one mole of the substance from a gas to a liquid.
molar heat of melting, ∆ Hmelt = the enthalpy change for the state change of one mole
of the substance from a solid to a liquid.
molar heat of freezing, ∆ Hfreez = the enthalpy change for the state change of one mole
of the substance from a liquid to a solid.
Since vaporization and condensation are opposite processes:
∆Hvap = - ∆ Hcond
Since melting and freezing are opposite processes:
∆ Hmelt = - ∆ Hfreez
1. A teacup containing 0.100 kg of water is at its freezing point. If the water freezes solid, how
much heat is released to the surroundings?
Draw an Enthalpy Diagram:
2. A sample of liquid mercury vaporizes. If the mercury is at its boiling point and has a mass of
0.325 grams, how much heat is released from/absorbed to its surroundings?
Type 4 - Standard molar enthalpy of solution
Molar enthalpy of solution, ∆Hsoln occurs when one mole of a solute dissolves in a solvent.
1. If 25.3 grams of sodium chloride were added to a glass of water and allowed to dissolve,
calculate the heat absorbed or released by the process. The molar enthalpy of solution for
sodium chloride is 3.9 kJ/mol.
Draw an Enthalpy Diagram
Section 16.3 Heating and Cooling Curves
Chemists use heating curves and cooling curves to represent temperature and phase changes in
substances
- a COOLING CURVE shows how the temperature of a substance changes as heat is
removed from it.
- a HEATING CURVE shows how the temperature of a substance changes as heat is
added to it.
These curves show temperature on the y-axis and either heat transfer or time on the x-axis.
** note: phase changes are the flat (horizontal) parts of the curve while warming/cooling changes
are the slanted parts of the curve.
(Read stages on p. 651 and 652 to understand when KE is changing versus potential energy is
changing)
{practice drawing heating/cooling curves before moving on to problems}
Problems - Total heat absorbed or released by a system
These problems require 3 types equations:
1. q = m x c x ΔT (Determines the heat when the KE is changing - the slanted part of the curve)
2. q = nΔH
(Determines the heat when the PE is changing - the horizontal part of the curve)
3. q total = q1 + q2 + q3 ........
Note: the equation order for 1 and 2 depends on the shape of the curve.
Examples:
1. 1.451 kg of water at 25.20C is added to a kettle and the water is completely vaporized.
A. Draw a heating curve for this process.
B. How much heat is required for this process.
2. A metal bucket containing 532.1 grams of ice at 0.00 0C is placed on a wood burning stove.
The ice is melted and then heated to 45.210C.
A. Draw a heating curve for this process.
B. How much heat is required to melt the ice and heat the water to 45.210C?
3. A sample of 36.8 grams of ethanol gas at 300.00C is cooled to liquid ethanol at 25.50C. The
specific heat capacity of ethanol gas is 1.43 J/g x 0C, and the specific heat of liquid ethanol is
2.45 J/g x 0C. The boiling point of ethanol is 78.50C.
A. Draw a cooling curve for this process.
B. Calculate the heat released during this process