CHAPTER 6:
PRINCIPLES OF
REACTIVITY
ENERGY AND CHEMICAL
REACTIONS
6.0 Objectives:







Describe various forms of energy and energy transfer.
Understand the terms reactant –favored, product-favored,
and thermodynamics.
Differentiate between kinetic and potential energy and know
the SI unit used to measure thermal energy.
Understand the term specific heat capacity and know how to
calculate amount of thermal energy transferred from one
object to another.
Use heat of fusion and heat of vaporization to solve simple
thermal problems.
Recognize and correctly use vocabulary related to
thermodynamics: system, surroundings, endothermic,
exothermic, enthalpy, first law of thermodynamics, and
calorimetry.
Calculate the enthalpy change of a system in several ways:
graphically, experimentally, and using Hess’s Law and the
summation equation.
Homework:
 Homework

Energy and Specific Heat
 Homework

#4 – 51, 53, 83, Hess’s Law Worksheet
Hess’s Law
 Homework
worksheet

#3 – 39, 41, 43, 45, 47, 79
Calorimetry
 Homework

#2 – 27, 29, 31, 33, 35, 37
Phase Changes and Enthalpy
 Homework

#1 – 13, 15, 17, 19, 21, 69, 71, 73
#5 – 55, 57, 61, 63, Heats of Formation
Heats of Formation
6.1 ENERGY: SOME BASIC
PRINCIPLES

1. Definitions
 Thermodynamics
the study of the effects of work, heat, and energy on a
system
 “energy changes that occur during chemical reactions”

 Energy
 “capacity
to do work”
 Kinetic
Energy
 “energy
of motion”
 KE = ½ mv2
 Includes Mechanical, Thermal, Electric, Sound
 Types of Motion

translational, vibrational, rotational
 Potential
 “energy
Energy
of position or arrangement”
 Examples: gravitational potential,
spring potential,
chemical,
electrostatic
2.
Law of Conservation of Energy, or First
Law of Thermodynamics
 “Energy
cannot be created nor destroyed”
 Energy of the universe is constant
3.
Heat vs. Temperature
 definitions
 Heat
and Temperature ARE NOT the same!
 Heat
(q)-

Total kinetic energy of the atoms of a substance; HEAT IS A
FORM OF ENERGY
 Temperature


(T)-
a measure of the average kinetic energy of the molecules in a
system
Measure temperature in degrees Kelvin
 4.
System vs. Surroundings

System- what is being studied
Surroundings- everything else

 5.
Direction of Heat Flow – 3 Principles
1. Heat transfers from a hotter object to a cooler one
2. K.E. transfer occurs until both objects are at the
same temp.
3. Heat lost by hotter object = heat gained by cooler
object
 6.
Exothermic vs. Endothermic
processes


Exothermic- heat is transferred from the
system to the surroundings
Endothermic- heat is transferred from the
surroundings to the system

7. Units of Energy
 Joule
(J) = SI unit of energy
 kg*m2/s2
= Force x distance = Work
 1 KJ = 1000 J
 calorie
(cal)
 amount
of energy required to raise 1 g of H2O 1 oC
 1 cal = 4.184 J
 Calorie
1
(Cal)
Cal = 1000 cal = 1 kcal
6.2 SPECIFIC HEAT CAPACITY AND
HEAT TRANSFER

1. Specific heat capacity, definition, units,
examples, c of water
 Specific
heat capacity
 “c”
- quantity of heat required to change the temp of the
system by 1oC
 Units = J/g.oC or J/g.K
 c-H2O = 4.184 J/g.oC
Examples:
 Which
heats up faster? An empty pan or a pan full of water?
 2.
Calculating heat
changes:
 equation
q
= m c T

Ex6.1 How many kilojoules of heat are required to increase
the temperature of 35.5g of iron from 23.6oC to 434oC?
The specific heat of iron is 0.451J/g.K.

(6.57 KJ)

Ex6.2 What is the specific heat of benzene if 3,450J of heat is
added to a 150.0g sample of benzene and its temperature
increases from 22.5oC to 35.8oC?

C = 1.73 J/g.K

5. Sign conventions
q
= (+)
 Heat
is being put into the system
 Endothermic
q
= (-)
 Heat
is being released from the system
 Exothermic
 6.
Heat transfer – definition and equation
 Heat
is transferred from one substance to
another
 Can measure this using q = mcT
 -q lost = q gained (conservation of energy)

Ex6.3 A 358.11g piece of lead was heated in boiling water in Salt Lake
City to 94.1oC. It was removed from the water and placed into 100.0mL
of water in a Styrofoam cup. The initial temperature of the water was
18.7oC and the final temperature of the lead and water was 26.1oC.
What is the specific heat of lead according to this data?
6.3 ENERGY AND CHANGES OF
STATE
 1.

Hfus & Hvap q=m H
Hfus – heat required to covert a substance from a
solid to a liquid at its melting point (MP)
 333

J/g for H2O
Hvap – heat required to covert a substance from a
liquid to a gas at its boiling point (BP)
 2256
 cH2O
J/g for H2O
solid = 2.09 J/g oC
 cH2O liq. = 4.184 J/g oC
 cH2O gas = 1.84 J/g oC

2. Graph – Equations, PE and KE, Phase
changes

Ex6.4 Calculate the quantity of heat needed to convert
125 g of ice at –25.0oC to steam at 175oC.
6.4 FIRST LAW OF THERMODYNAMICS

1. E = q + w
 1st
Law of Thermo:
 “energy
change for a system is the sum of heat transferred
and the work transferred to or from the system”
E
= internal energy
 Sum
of the potential and kinetic energy of all atoms &
molecules

2. Enthalpy

Enthalpy: H = E + PV (sum of Internal E and work)
 Work
is represented by PV
 Like Internal E, enthalpy cannot be directly measured
 Enthalpy
Change: H = E + P V
 Heat
of Rxn at constant pressure (most rxns open to
atmosphere)

Standard Temperature and Pressure (STP)
 1 atm, 273 K (0oC)

3. State functions
 Depends
only on present conditions (doesn’t matter how it
got there)
State Function
Path Function
Independent of path taken to Dependent on path taken to
establish property or value.
establish property or value.
Can integrate using final and
initial values.
Need multiple integrals and
limits of integration in order to
integrate.
Multiple steps result in same
value.
Multiple steps result in different
value.
Based on established state of
system (temperature, pressure, Based on how state of system
amount, and identity of
was established.
system).
Normally represented by an
uppercase letter.1
Normally represented by a
lowercase letter.1
6.5 ENTHALPY CHANGES FOR CHEMICAL RXNS
1.
Measuring H
 a.
Graphically
 b.
Experimentally
 Determine
how much heat is evolved or
absorbed for a given reaction
 Calorimetry (next topic)

c. Mathematically
 Hrxn
= Energy (Products) - Energy (Reactants)
 2.
Principles of Enthalpy
 Identity
and state are important
 H2O (l)
is different from H2O (g)
 Different energies for the different states
 Positive
or negative values
 (-)
= Exothermic; heat added to products side
 (+) = Endothermic; heat added to reactants side
 Stoichiometric
 Coefficients
quantity
balance the chemical equation
 May not always have whole numbers or reduced coefficients
 Changing the coefficients changes the H!
 Reverse
reactions
 Reactions
are cyclic in nature
 If Rxn is reversed, must change sign
 Ex.
3.

2HgO (s) --> 2Hg (l)+ O2 (g) H= +181.66 kJ
2Hg (l)+ O2 (g) --> 2HgO (s) H= -181.66 kJ
Thermochemical equation –
H2 (g) + ½ O2 (g)  H2O (l)
H = -285.8 kJ
 Ex6.5 When 5.00g of sucrose, C12H22O11, is burned to
produce gaseous carbon dioxide and liquid water, 82 kJ of
energy are produced. Write the correct thermochemical
equation for this reaction.
6.6 CALORIMETRY

1. Definition and assumptions
 Calorimetry
- Technique used to measure a quantity of heat
based on conservation of energy
 Heat flows from the hotter to the cooler object resulting in a
temp change
 Heat lost by hot solids is gained by water in calorimeter
 q = mcT

2. Coffee cup calorimeter, equation
qH2O = -qobject
qobject = amt. Of heat given off by
substance as it cools to a
common T with H2O
cm∆T = cm ∆T
qH2Oor Soln. = -qrxn
qH2O = amt. Of heat given off by
rxn
qrxn = cm ∆T

Ex6.6 A student placed 0.500g of Mg(s) in a coffee cup calorimeter
along with 100.0mL of 1.00M HCl. The reaction produces aqueous
magnesium chloride and hydrogen gas. The temperature of the
solution rises from 22.2 to 44.8oC. What is the enthalpy of this reaction?
Assume that the specific heat capacity of the solution is 4.2 J/g K and
the density of the solution is 1.00g/mL.
 Advantages and disadvantages:


Easy to do; simple instrumentation; relatively
robust/accurate
Sig figs are limited to tools used; some heat is lost to cup

4. Bomb Calorimeter
 qrxn
+ qbomb + qwater = 0
 Initiate
rxn by ignition wire
 Treat whole metal assembly as one Sp. Heat capacity
 Use for reactions involving gases

Ex. Combustion of Gasoline (Octane)
Cbomb T
 qwater = cmT
 qrxn= -(qbomb+qwater)
 qbomb =

Ex6.7 When a 1.00g sample of naphthalene, C10H8, is burned in
a bomb calorimeter with a heat capacity of 13.24kJ/K and
1.20kg of water, the temperature rises from 22.42 to 25.46oC.
Write the thermochemical equation for this reaction.
6.7 HESS’S LAW

1. Statement and usage
 “If
a reaction is the sum of two or more other
reactions, H for the overall process is the sum of the
H values of those reactions”

2. Ex6.8 From the following information calculate the
enthalpy of the formation of one mole of SO3 from its
elements.
S(s) + O2(g)  SO2(g)
H = -296.8 kJ
2SO2(g) + O2(g)  2SO3(g) H = -197.8 kJ

3.Ex6.9 Use the following equations and a Hess’s Law process to
determine the enthalpy of formation of 1.00 mole of NO(g) from its
reactants.
N2(g) + 3H2(g)  2NH3(g)
4NH3(g) + 5O2(g)  4NO (g) + 6H2O(g)
H2(g) + 1/2O2(g)  H2O(g)
H = -91.8kJ
H = -906.2kJ
H = -241.8kJ

1. Definition
 “Enthalpy
change for the formation of 1 mole of a
compound directly from its elements in their
standard states (1 atm; 25oC)”
 Hfo
 2.
Examples
 CO2 (g)
 C(s)
 N2 (g)

+ O2 (g)  CO2 (g)
Hfo= -393.5 kJ
Hfo = 0 kJ!!
3. Reference book values
 See
Reference booklet for Standard
enthalpy of formation values (Hfo )
 4.
Summation equation
Hrxn
o
= H
products
o
- H
reactants

5. Ex6.10 Use the summation equation to determine
the enthalpy of the following reaction:
4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)

Ex6.11 Use the table of heats for formation and the
equation in example 6.7 to calculate the standard heat of
formation for naphthalene, C10H8.