Chapter 5: Thermochemistry. Problems: 5.1-5.95, 5.97-98, 5.100, 5.106, 5.108, 5.118-5.119, 5.121, 5.126. Energy: Basic Concepts and Definitions. energy: capacity to do work or to produce heat thermodynamics : study of energy and its transformations from one form to another
Problems: 5.1-5.95, 5.97-98, 5.100, 5.106, 5.108, 5.118-5.119, 5.121, 5.126
energy:capacity to do work or to produce heat
thermodynamics: study of energy and its transformations from one form to another
thermochemistry: study of heat flow accompanying chemical reactions
heat: energy that is transferred from a body at a higher temperature to one at a lower temperature heat always transfers from the hotter to the cooler object!
One becomes hotter by gaining heat.
One becomes colder by losing heat—i.e., when you “feel cold”, you are actually losing heat!
a. You burn your hand on a hot frying pan.
________________ loses heat, and ______________ gains the heat.
b. Your tongue feels cold when you eat ice cream.
_______________ loses heat, and ______________ gains the heat.
Example: A student puts a few drops of rubbing alcohol on her palm then spreads it out with her finger and notices that the area with the alcohol felt cooler. A minute later her palm is dry. Explain what physical change has occurred and why her palm felt cooler.
work (w): w = F d
Work is done whenever a force (F) moves an object a specified distance (d)
kinetic energy (KE):energy associated with an object’s motion
Example: a car moving at 75 mph has much greater KE than the same car moving at 15 mph
Greater damage if the car crashes at 75 mph than at 15 mph
KE = ½ mv2m=mass, v=velocity
potential energy (PE): energy due to position or its composition (chemical bonds)
A 10-lb bowling ball has higher PE when it is 10 feet off the ground compared to 10 inches off the ground
Greater damage on your foot after falling 10 feet compared to falling only 10 inches
positional P.E. = mass force of gravity height
In terms of chemical bonds, the stronger the bond,
more energy is required to break the bond,
the higher the potential energy of the bond
state function: a property that is based only on the physical state and chemical composition of a substance, so it is independent of the path followed to achieve that state or composition.
Potential energy is a state function because the potential energy of two skiers at the top of a given hill would be the same whether they climbed that hill or took a ski lift to the top.
joule (J): 1 J =
Energy is neither created nor destroyed but converted from one form to another.
For example, the kinetic energy of a car can cause considerable damage if the car is stopped suddenly in a crash.
Example (cont’d): At what speed (in miles per hour) must a 7.3102 kg Smart Car move to have the same kinetic energy as the Hummer H2?
system: that part of the universe being studied
surroundings: the rest of the universe outside the system
Systems can be isolated, closedor open.
Let q = heat flow,
endothermic change: a physical or chemical change that requires energy or heat to occur
Boiling water requires energy:
H2O(l) + heat H2O(g)
Electrolysis of water requires energy:
2 H2O(l) + electrical energy 2 H2(g) + O2(g)
exothermic change: a physical or chemical change that releases energy or heat
Water condensing releases energy:
H2O(g) H2O(l) + heat
Hydrogen burning releases energy:
2 H2(g) + O2(g) 2 H2O(g) + heat
For physical changes, consider whether the reactants or products have more kinetic energy.
What causes the Heat of Reaction?
Breaking and Forming Bonds
heat of reaction (qreaction):heat associated with a chemical reaction
If Endothermic Reaction
For chemical changes, observe if the surroundings (including you) feel hotter or colder after the reaction has occurred.
Which of the following are endothermic changes:
freezing vaporizing sublimation
deposition melting condensation
When a student dissolves ammonium chloride in a large test tube, he notices the test tube feels colder. Explain what is releasing heat and what is gaining heat.
The energy of the universe is constant. The energy gained or lost by a system must equal the energy lost or gained by the surroundings.
Essentially, the Law of Conservation of Energy: Energy can neither be created or destroyed but converted from one form to another.
A system’s Internal Energy (E) = kinetic energy (KE) + potential energy (PE) of all the particles in the system.
A system’s internal energy (E) can be changed using heat (q), work (w), or both:
∆E = q + w
In this course, we will focus on work that involves the expansion or compression of gases.
Consider the following sign conventions for the change in volume, ∆V, for the system:
Work is done to the system. w is positive
work is defined as w = –P∆V
Thus, for the expansion or compression of a gas, the change in internal energy (∆E) can be shown as:
∆E = q + w or ∆E = q – P∆V
Consider the combustion of octane (C8H18), a primary component of gasoline.
2 C8H18(l) + 25 O2(g) 18 H2O(g) + 16 CO2(g)
a. Calculate the change in volume (in L) due to the total number of moles of gases produced when 1.000 gal (~3.784 L) of octane undergoes combustion at 1.00 atm and 25.00°C. The density of octane at 25.00°C is 0.703 g/mL.
b. What is the volume of the gases at 475 K (the temperature of a car engine)?
c. Use the equation, w = – P∆V, to calculate the work (in kJ) done by the system (the reaction) during the combustion of 1.000 gal of octane. (Use 1 Latm = 101.325 J).
d. If 1 kJ 1 kWs (kilowatt·second), then calculate the energy (in kJ) in a kilowatt·hour (kWh).
e. Consider a new kind of vehicle that could be powered by an electrical current similar to that used in our homes. Calculate the cost of electricity needed to produce the same amount of energy (determined in part c) as the combustion of 1.00 gal of gasoline if Seattle City light charges 9.55 cents/kWh.
f. The combustion of 1.000 gal of octane produces about 1.2105 kJ of heat. Compare the change in internal energy due to work calculated in part c with the heat of the reaction. How much of the internal energy change is due to the work done by the system due to gas expansion? (Hint: Compare the absolute value for heat versus work.)
For the remainder of the chapter, we will focus mainly on the heat of reaction and assume the change in energy due to gas expansion work is negligible for the reactions considered.
Enthalpy(H): the sum of a system’s internal energy and the product of its pressure and volume
H = E + PV
Thus, the enthalpy change (∆H) is: ∆H = ∆E + ∆(PV)
or at constant atmospheric pressure is: ∆H = ∆E + P∆V
Since the previous definition of internal energy, ∆E = q – P∆V, the equation can be rewritten as∆H = q – P∆V + P∆V = qp
(the subscript p means “at constant pressure”)
Thus, the enthalpy change (∆H) refers to heat flow into and out of a system under constant pressure (usually the case since reactions occur under atmospheric pressure),
qreaction= ∆H = Hproducts – Hreactants
For an endothermic reaction: ∆H = positive
For an exothermic reaction: ∆H = negative
Consider the changes that H2O undergoes when a block of ice is taken from a freezer and heated in a pan until it is completed converted into steam.
specific heat (cs):amount of heat necessary to raise the temperature of 1 gram of any substance by 1°C; has units of J/g°C
molar heat capacity (cp): heat capacity per mole of a substance (in J/mol·°C)
heat capacity: amount of heat necessary to raise the temperature of a given amount of any substance by 1°C; in units of J/°C
Use the following equations to solve calorimetry problems:
q = n cp∆T or q = csm ∆T
where ∆T=change in temperature, n=# of moles, and m=mass.
a. If 279.9 J is required to raise the temperature from 23.0°C to 99.5°C for a 15.5-g sample of silver, what is the specific heat of silver?
b. A beaker with 100.0 g of water is heated from 25.0°C to its boiling point. If the specific heat of water is 4.184 J/g·°C, how much heat is required to heat the water?
c. Determine the final temperature for a 100.0 g sample of iron at 25.0°C heated with the amount of heat calculated in a. given the molar heat capacity of iron is 25.1 J/mol·°C.
calorie (cal): unit of energy used most often in the US. Is equal to the amount of energy required to raise the temperature of 1 g of water by 1˚C
1 cal 4.184 J (Note: This is EXACT!)
But a nutritional calorie (abbreviated Cal) is 1000 cal:
1 Cal = 1 kcal = 4.184 kJ
When consuming an ice-cold drink, one must raise the temperature of the beverage to 37.0°C (normal body temperature). Can one lose weight by drinking ice-cold beverages if the body uses up about 1 calorie per gram of water per degree Celsius (i.e. the specific heat of water = 1.00 cal/g·°C) to consume the drink?
a. Calculate the energy expended (in Cal) to consume a 12-oz beer (about 355 mL) if the beer is initially at 4.0°C. Assume the drink is mostly water and its density is 1.01 g/mL.
b. If the label indicates 103 Cal, what is the net calorie gain or loss when a person consumes this beer? Is this a viable weight loss alternative?
c. Calculate the amount of heat (in kJ) required to heat 1.00 kg (~1 L) of water at 25°C to its boiling point.
Another energy unit is the British thermal unit (abbreviated Btu). A Btu is the energy required to raise the temperature of 1 pound of water by 1°F when water is most dense (at 39°C).
The heating power of many gas cooktops is often given in Btu’s. Calculate the time (in minutes) required to heat 1.00 kg of water at 25°C to boiling using a 12,000 Btu per hour burner. (Assume complete energy transfer from the burner to the water.) Use 1 kWh = 3412 Btu and 1 kJ = 1 kWs.