Chapter 6 Principles of Reactivity: Energy and Chemical Reactions

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Chapter 6Principles of Reactivity Energy and
Chemical Reactions
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Definitions Class Lecture Notes OWL
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1. INTRODUCTION
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• CHEMISTRY - The study of the properties,
  composition, and structure of matter, the
  physical and chemical changes it undergoes, and
  the ENERGY liberated or absorbed during those
  changes.
• THERMODYNAMICS - Derived from the Greek words
  for heat and power it is the study of all forms
  of energy and the interconversions among the
  different forms.
• THERMOCHEMISTRY - The study of the energy
  liberated (released) or absorbed by chemical or
  physical changes of matter.

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SYSTEM AND SURROUNDINGS - 1) System - The
part of the universe a scientist is
interested in. 2) Surroundings - Everything
else in the universe that is outside of
the system.
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3) Boundary - A real or imaginary barrier
between the system and its surroundings through
which THERMAL ENERGY may flow, work may appear or
disappear, and matter may or may not be
exchanged. 4) Closed System - A system that
does not exchange matter with its surroundings.
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• 5) Open System - A system that may exchange
  THERMAL ENERGY and MATTER with its
  surroundings.
• 6) Isolated System - A system that does not
  exchange MATTER or THERMAL ENERGY with its
  surroundings.

2. CHANGE - WHY DOES IT HAPPEN? A.
SPONTANEOUS CHANGE - A change that takes place
by itself. B. NON-SPONTANEOUS CHANGE -
The opposite of a spontaneous change.
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C. Factors Affecting Change - 1) Energy - The
ability to do work. a) Thermal b)
Electrical c) Radiant d) Chemical e)
Mechanical f) Nuclear g) Kinetic h)
Potential 2) Entropy - A measure of
disorder. 3) The Rate of Change - (Kinetics)
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D. Predicting Change - Requires a knowledge
of how energy and entropy of a system will
change as a result of the change and whether the
change will take place at a practical speed.
3. Temperature, Thermal Energy, Heat A.
Temperature - An indirect measure of the average
kinetic energy of the molecules, atoms, or ions
in the material A measure of the hot- ness or
coldness of a material an INTENSIVE property
of matter. B. Thermal Energy - The energy that
is transferred from hotter objects to colder
objects due to the kinetic energy of the
molecules, atoms, or ions an EXTENSIVE property
of matter.
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C. Heat - The transfer of thermal energy
that results from a difference in temperature
the process of transferring thermal energy
from hotter objects to colder objects. D.
Potential Energy - Stored energy that is a
related to an objects relative position. E.
Kinetic Energy - The energy of motion. K.E.
½ mv2
K.E. ½(2 kg)(1 m/s)2 1kg-m2/s2 1 Joule
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4. The First Law of Thermodynamics Energy is
neither created nor destroyed. Energy given up
by the system must be absorbed by the
surroundings. This type of change
is EXOTHERMIC. Energy absorbed by the system
must come from the surroundings. This type of
change is ENDO- THERMIC. A. Internal Energy, E
- The total energy of a system.
D E q w Efinal - Einitial
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Constant Volume Calorimetry

• E q w
• (w 0 at
• constant V)

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B. Enthalpy, H - The Thermal Energy gained or
lost by a system when the system under- goes a
change under constant pressure.
DH Hfinal - Hinitial
We can only measure The Change in enthalpy, not
the absolute enthalpy. Enthalpy is a state
function.
Exothermic Change - Changes during which the
system gives off thermal energy and DH lt 0
(negative).
Endothermic Change - Changes during which the
system absorbs thermal energy and DH gt 0
(positive).
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D H
Constant Pressure Calorimetry
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Examples of Exothermic Changes - H2O (l) H2O
(s) DH lt 0
H2O (l)
This is an exothermic process because the final
energy state is lower than the initial energy
state.
DH -
H2O (s)
2 H2 (g) O2 (g) 2 H2O (l) DH lt 0
2 H2 (g) O2 (g)
Water is lower in enthalpy than hydrogen and
oxygen are. Exothermic
DH -
2 H2O (l)
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Examples of Endothermic Changes - H2O (l) H2O
(g) DH gt 0
H2O (g)
This is Endothermic because the final state is
higher in enthalpy than the initial state.
DH
H2O (l)
CaCO3 (s) ? CaO (s) CO2 (g)
DH gt 0
CaO (s) CO2 (g)
CaO (s) and CO2 (g) are higher in enthalpy than
CaCO3 (s).
DH
CaCO3 (s)
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Class Exercises - Endothermic or Exothermic ?
3 H2 (g) N2 (g) 2 NH3 (g)
DH - 46.1 kJ/mol
N2 (g) 2 O2 (g) 2 NO2 (g)
DH 33.2 kJ/mol
C3H8 (g) 5 O2 (g)
3 CO2 (g)
4 H2O (g)
Sign of DH ?? Negative
NH4Cl dissolves in water with a decrease
in temperature.
Sign of DH ?? Positive
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Thermochemical Equations - A chemical equa- tion
that includes an enthalpy change explicitly.
2 H2 (g, 1 atm) O2 (g, 1 atm) 2 H2O
(l) ?H298 K - 571.7 kJ
A. ?H subscript indicates temperature of rxn. B.
?H represents thermal energy evolved when 2 moles
of H2O (l) are formed!
H2 (g, 1 atm) 1/2 O2 (g, 1 atm) H2O
(l) ?H298 K - 285.8 kJ
C. The sign of ?H is reversed when the
equation is reversed!
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H2O (l) ? H2 (g, 1 atm) 1/2 O2 (g, 1
atm) ?H298 K 285.8 kJ
D. The enthalpy change during a reaction can be
considered as a reactant (endothermic) or as
a product (exothermic).
H2 (g, 1 atm) 1/2 O2 (g, 1 atm) ? H2O (l)
285.8 kJ
285.8 kJ H2O (l) ? H2 (g, 1 atm) 1/2 O2
(g, 1 atm)
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Practice Problem When 2.0 moles of isooctane
are burned, 10 930.9 kJ of thermal energy are
liberated under constant temperature conditions.
How many kJ will be liberated when 369 g of
isooctane are burned?
Problem 369 g Equation 2 C8H18 (l) 25 O2
(g) ? 16 CO2 (g) 18 H2O (l) D H -10
930.9 kJ Molar Masses 114.231 ? kJ
When in doubt, calculate MOLES!
(369 g C8H18)(1 mol C8H18/114.231 g C8H18)
3.230 mol C8H18
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(3.230 mol C8H18)(-10 930.9 kJ/2 mol C8H18)
- 1.77 x 104 kJ
Practice Problem The DHrxn for the burning of
H2 (g) to form H2O (l) is 285.83 kJ/mol H2O.
How many grams of H2 (g) are needed to produce
539.63 kJ of thermal energy?
Problem ? g - 539.63
kJ Equation H2 (g) ½ O2 (g) ? H2O (l)
DH -285.83 kJ Molar Masses 2.015 88 u
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(- 539.63 kJ)(1 mol H2/-285.83 kJ) 1.887 94
mol H2
(1.887 94 mol H2)(2. 015 88 g H2/mol H2)
3.8059 g H2
Molar Mass
Avogadros Number
MOLES
MASS
PARTICLES
Molarity
PV nRT
Molar Heat of Reaction
VOLUME
P, V, T
Thermal Energy
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• Standard Enthalpy Changes - A standard enthalpy
  change, DH0, is the enthalpy change for a
  reaction in which each reactant and each product
  is in its Standard State.
• Solids - Pure solid at 1 atm
• Liquids Pure liquid at 1 atm
• Gas Ideal gas at 1 atm partial pressure
• Solute Ideal solution at 1 M conc.

H2 (g, 0.5 atm) Cl2 (g, 0.5 atm) ? 2 HCl
(g, 1 atm) D Hrxn, NOT DHorxn
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Standard Enthalpy of Formation - DHf The
enthalpy change accompanying the formation of one
mole of a substance in its standard state from
its elements, each in their standard states and
most stable form.
Hg (l) Cl2 (g) ? HgCl2 (s) DHfo -221.3
kJ S (s) O2 (g) ? SO2 (g) DHfo -296.8
kJ H2 (g) ? H2 (g) DHfo 0 kJ
Elements in their standard states have zero
enthalpies of formation!
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Standard Enthalpy of Combustion - DHcomb The
enthalpy change accompanying the combustion of
one mole of a substance in O2 in its standard
state from its elements, each in their standard
states and most stable form.
H2 (g, 1 atm) ½ O2 (g, 1 atm) ? H2O (l, 1
atm) DHcombo -285.8 kJ C8H18 (l)
25/2 O2 (g) ? 8 CO2 (g) 9 H2O
(l) DHcombo -5455.6 kJ C2H5OH (l)
7/2 O2 (g) ? 2 CO2 (g) 3 H2O
(l) DHcombo -1366.8 kJ