Download Copy of Ch6-Energy in Chemical Reactions

Louisiana Tech University, Chemistry 100. Group Activity on Chapter
6. Energy and Chemical Reactions: Forms of energy and Enthalpy
Why?
What is first law of thermodynamics? What are the forms of energy? How we
produce them? How is chemistry useful understanding energy produced? Why heat is an
important form of energy? How heat is measured using calorimetry? What is enthalpy?
How does an enthalpy change come from chemical bonds? What are the definitions of
two forms of enthalpy: enthalpy of a reaction (Hrex) and standard enthalpy of formation
(Hf)? What is Hess's Law and how it is applied to calculate enthalpy change for a
reaction. How would you evaluate chemical fuels for home and industries? How would
we evaluate foods and their contributions to caloric intake?
Learning Objectives
The students should be able to understand the concepts and have working knowledge in
the following:
6.7 Where does the Energy come from?
1. Understand the origin of the enthalpy change for a chemical reaction in terms of bond
enthalpies.
6.8 Measuring Enthalpy Changes: Calorimetry
1. Describe how calorimeters can measure the quantity of thermal energy transferred
during a reaction.
Success Criteria
Understand the first law of thermodynamics, the forms of energy in the
Universe, the definition of system and surrounding, use of chemistry to understand
energy production, heat is a important form of energy, calorimetry, enthalpy, and
enthalpy of a reaction (Hrex) and standard enthalpy of formation (Hf)
Resources
Chemistry: The Molecular Science 1st Edition, John W. Moore, Conrad L. Stanitski and
Peter C. Jurs.
Prerequisites
High school chemistry: Definition of matter. Scientific method, Dalton’s Atomic theory,
Components of atoms: nucleus (proton and neutrons) and electrons. Atoms, molecules
and ions.
New Concepts
The forms of energy in the Universe
Scientists define energy as the ability to do work. People have learned how to change
energy from one form to another so that we can do work more easily and live more
comfortably. Energy is found in different forms, such as light, heat, sound and motion.
There are many forms of energy, but they can all be put into two categories: Macroscopic
and microscopic, and kinetic and potential.
Macroscopic Energy: Energy associated with large moving objects.
Microscopic Energy: Invisible internal energy associated with microscopic energy on
the atomic and molecular scale.
Kinetic Energy: Due to motion––of waves, electrons, atoms, molecules, substances, and
objects.
Potential Energy: Due to stored energy and the energy of position.
Stored Mechanical Energy: Energy in compressed springs and stretched rubber bands.
Chemical Energy: Energy stored in the bonds of atoms and molecules.
Nuclear Energy: Energy of nucleus released as fission and fusion energy.
Gravitational Energy: Energy of position or place. A rock resting at the top of a hill and
water in a reservoir behind a dam has gravitational potential energy.
Thermodynamics
The study of the laws that govern the conversion of energy from one form to another, the
direction in which heat will flow, and the availability of energy to do work and move an
object.
First law of thermodynamics
The total energy of the system is constant. Energy is neither created nor destroyed.
For thermodynamics it is specifically
 = q + w;  = internal energy, q=heat change to the system and w= work involved by
the system. w= -PV (volume expansion work, V= Vf-Vi)
Thermodynamic Standard State
Temperature:25C or 298 K 0, and pressure: 1 atm or 760 torr
Measuring Energy Changes:
System and surrounding
System – particular portion of the universe on which we wish o focus our attention
Surroundings – everything that’s not the system.
Universe = system + surrounding
Open System: matter and energy can be exchanged between system and surroundings
Closed system: matter cannot be exchanged, energy can.
Isolated system: neither matter nor energy can be exchanged.
Sign Conventions: Watch for -/+ signs, they mean exothermic/endothermic, sometimes
they’re not given, must find out yourself.
– sign
+ sign
is heat ADDED TO a system FROM
-q is heat REMOVED FROM a
+q
system TO the surroundings
the surroundings
is work done BY a system ON the
-w is work done ON a system BY +w
the surroundings
surroundings
Heat Energy
Heat may be defined as energy that’s internal and kinetic in nature when two objects one
at a high temperature transfer heat to a lower temperature object. The internal energy may
be increased by transferring kinetic energy to the object from a higher temperature
(hotter) object - this is properly called heating.
Specific Heat: Specific Heat is amount of heat needed to increase temperature of 1g of a
material by 1oC.
Heat capacity = specific Heat x mass.
Heat gain = [Specific heat x mass x t] = - Heat loss= - [Specific heat x mass x t]
Temperature Drop/gain (t): Is the change in temperature in a calorimetric experiment.
Usually written as t: t = tfinal - tinitial; t = + for a gain and - for a drop.
Temperature
Temperature measures only the kinetic energy part of the internal energy. Hotter an
object more kinetic energy it has. When two objects are in thermal equilibrium they are
said to have the same temperature. This is a statement relating to zeroth law of
thermodynamics.
Calorimetry
Calorimetry is the branch of chemistry/physics that deals with the measurement and
calculation of heat changes. Calorimeter is the instrument to used measure heat changes.
According to Law of Conservation of Energy, heat absorbed
{Heat gain:= [Specific heat x mass x t] }by
a material should come from another material releasing
{-Heat loss: = - [Specific heat x mass x t]} it.
Enthalpy(H)
For systems that involve heat change (qp) and volume expansion work (PV) at
constant pressure, qp is called enthalpy (H).
E = qp + w; w = -PVE = qp -PV + w; qp = E + PV; H = E + PV
Enthalpy change = sum of the internal energy and volume expansion work.
Internal energy (E)
For systems that involve heat change (qv) at constant volume (V =0) is called internal
energy: E = qv + w; qv = E + 0; w =-PV = 0 since V =0;
E = qv = E (internal energy )
Internal energy = heat change (qv) at constant volume (V =0).
Chemistry to understand energy production: Thermochemistry
Chemical reactions release/absorb energy in the form of heat when potential enegey
stored in chemical bonds change. Chemical reactions like burning fossil fuels are used to
obtain energy for many purposes. When a regular chemical equation includes the change
of heat the equation is called a thermochemical equation.
E. g. 2H2 (g) + O2 (g) ---> 2H2O(l) ;H = - 572 kJ; H is called the enthalpy of
reaction. If H is + or - reaction is called endothermic and exothermic, respectively.
Enthalpy of a reaction (Hrex)
Enthalpy of a chemical reaction (Hrex)is the heat change of a chemical reaction at
constant pressure.
E. g. 2H2 (g) + O2 (g) ---> 2H2O(l); Hrex = - 572 kJ
Standard enthalpy of formation (Hf)
Enthalpy change that results from one mole of a substance being formed in a formation
reaction from its elements at their standard states.
E. g. H2 (g) + ½O2 (g) ---> H2O(l); H°f = -286 kJ/mole
H°f are tabulated in the appendix. Note H°f for an element is zero. Why?
Thermal Stoichiometry
E.g. Calculate the heat evolved (kJ) for the reaction:
CH4(g) + 2 O2(g) --->CO2(g) + 2 H2O(l); Horex = -890.3 kJ/mole
When 2.40 g of CH4 is burned in excess oxygen.
M. W. (CH4) = 16 g/mole
= -133.55 kJ
2.40 g CH4 1 mole CH4 -890.3 kJ
16 g CH4
1 mole CH4
Key Questions (relatively simple to answer using the Focus Information)
1) What is energy?
2) What forms of energy are available in the Universe?
a)
e)
b)
f)
c)
g)
d)
h)
3) Explain the differences between following categories of energy.
a) Kinetic Energy:
b) Potential Energy:
c) Macroscopic Energy:
e) Microscopic or Internal energy:
4) What is thermodynamics?
4) What is first law of thermodynamics?



5) Given  = q + w; [ = internal energy, q=heat change to the system and w= work
involved by the system. w= -PV (volume expansion work, V= Vf-Vi)].
Show this equation follows first law of thermodynamics.
6) Consider following chemical reactions.
a) Taken place at room temperature, 25C and atmospheric pressure:
2H2 (g) + O2 (g) ---> 2H2O(l); Hrex(1) = - 572 kJ
b) Taken place at room temperature, 200C and higher pressure:
2H2 (g) + O2 (g) ---> 2H2O(g); Hrex(2) = ? kJ
Compare the values of Hrex(1) to Hrex(2). Which one would be larger?
7) What is thermodynamic standard state?
8) A acid/base chemical reaction is carried out in 250 ml flask with a thermometer and
stirrer by mixing 50 mL 0.1 M HCl (aq) and 50 mL of 0.1 M NaOH(aq) solutions and
exotermic reaction observed.
a) Draw a diagram for this set up:
b) What items constitute the system?
c) What items constitute the surrounding?
d) What is releasing heat?
e) What is absorbing heat?
9) Describe the following systems:
a) Open System:
b) Closed system:
c) Isolated system:
10) What are the sings of following changes of the system?
a) Heat REMOVED FROM a reaction mixture system TO the flask:
System:
Surrounding:
Sign:
b) Heat ABSORBED a reaction FROM the surroundings:
System:
Surrounding:
Sign:
c) Work is performed ON a gas trapped inside a piston by the surroundings:
System:
Surrounding:
Sign:
d) Work done by a heat generated inside a gas in piston ON the surroundings:
System:
Surrounding:
Sign:
11) What is heat energy?
12) Why is Heat gain = [Specific heat x mass x t] = - Heat loss= - [Specific heat x mass
x t]?
13) When 5.8 g of potassium persulfate, K2S2O8, dissolve in 48.6 g of water (specific heat
= 4.18 J/g°C),the temperature is raised from 22.6°C to 29.0°C; how much energy is
released by the dissolution of this sample?
14) What is the specific heat of ethyl alcohol if 700.0 J of heat are required to raise the
temperature of an 80.0-g sample from 30.0°C to 45.0°C?
15) The reaction: S(s) + 3F2 (g) -----> SF6 (g) is studied in a bomb calorimeter. If 6.40 g
of sulfur is reacted with excess fluorine gas in a calorimeter whose heat capacity is
32.5 kJ/°C, the temperature inside the calorimeter rises from 21.3°C to 28.7°C.
Determine the heat produced if one mole of sulfur would react similarly.
16) What is the quantity of heat evolved when 100.0 g H2O(l) is formed from the
combustion of H2(g) and O2(g)?
a) Chemical reaction: 2H2 (g) + O2 (g) ---> 2H2O(l); Hrex(1) = - 572 kJ
b) Reaction exotermic or endothermic:
c) Moles of H2O:
d) Conversion factor moles to heat:
e) Heat:
17) Look up in the appendix for heat of formation ((Hf) values:
a) H2O(l) =
b) Na(s) =
c) O2(g) =
d) Cl-(aq) =
18) Which of the following substances has a heat of formation ((Hf) equal to zero at
25°C and 1 atm?
a) H2O(g)
b) Na(g)
c) O2(g)
d) Cl-(aq)