Transcript Energy & Chemical Change

Energy & Chemical
Change
Chapter 16
16.1 Energy
• What is Energy?
• What are some types of energy
that you are familiar with?
– Kinetic
– Potential
– Thermal
– Electrical
– Chemical
– Nuclear
A. The Nature of Energy
• Energy: ability to do work
or produce heat
• Potential Energy: energy
due to composition
(chemical) or position of an
object (gravitational)
• Kinetic Energy: energy of
motion
• Kinetic energy of a
substance is directly related
to the constant random
motion of it particles and is
proportional to
temperature.
• Chemical Potential Energy
of a substance depends
upon its composition
–Type of atoms
–# & type of chemical
bonds
–How atoms are arranged
Law of conservation of energy
• States that in any chemical
reaction or physical
process, energy can be
converted from one form to
another, but it is neither
created nor destroyed
Chemical potential energy
• Energy stored in a substance
because of its composition
• Ex. Gasoline – when burned
chemical potential energy is
converted to useful mechanical
energy.
• Heat (q): process of energy
flowing from a warmer object
to a cooler object
How is Heat Transferred?
1. Conduction – transfer of heat within
solid objects by direct contact
2. Convection – transfer of heat within
fluids (liquids & gases)
3. Radiation – transfer of heat by
electromagnetic radiation (like the sun)
Measuring Heat - Units
• calorie: amount of heat
required to raise temperature
of one gram of pure water one
degree Celsius
• Joule: SI unit of heat & energy
– 1 calorie = 4.184 joules
• 1000 calories = 1 Kcal or 1
nutritional Calorie
• Practice Problems p.492
Specific Heat
• Amount of heat required to raise the
temperature of one gram of that
substance by one degree Celsius
• Remember water has a high specific
heat – it takes lots of energy to
change it’s temperature
Calculating heat evolved & absorbed
q = m x c x T
q = heat absorbed or released (the value
is positive if heat is absorbed and
negative if heat is released)
m = mass of sample in grams
c = specific heat of substance (can be
determined or looked up in a table)
T = difference between final
temperature & initial temperature
Phase Changes – know for your TEST!
Which phase changes require energy? (endothermic)
• Melting, evaporation, sublimation
Which phase changes release energy? (exothermic)
• Freezing, condensation, deposition
Energy and Phase Change
• Heat of vaporization - energy required to
change one gram of a substance from liquid
to gas.
• Heat of condensation - energy released
when one gram of a substance changes from
gas to liquid.
• For water 540 cal/g
Energy and Phase Change
• Heat of fusion - energy required to
change one gram of a substance from
solid to liquid.
• Heat of solidification - energy released
when one gram of a substance changes
from liquid to solid.
• For water 80 cal/g
Heating Curve for Water
120
Steam
Water and
Steam
Temperature in Celsius
100
80
60
Water
40
20
0
Ice
Water
and Ice
-20
0
40
120
220
Pressure in mmHg or torr
760
800
Heating Curve for Water
Temperature in Celsius
120
Heat
of and
Water
Steam
Vaporization
100
Steam
80
60
Water
40
20
0
Ice
Water
and Ice
-20
0
40
120
220
Pressure in mmHg or torr
760 800
Heating Curve for Water
Temperature in Celsius
120
Steam
Water and
Steam
100
80
60
Water
40
20
Heat ofWater
0
IceFusionand Ice
-20
0
40
120
220
Pressure in mmHg or torr
760 800
Heating Curve for Water
Temperature in Celsius
120
Steam
Steam
Water and
Steam
100
80
Water
60
Water
40
20
0
IceIce
Slope =
Specific Heat
Water
and Ice
-20
0
40
120
220
Pressure in mmHg or torr
760 800
Heating Curve for Water
Temperature in Celsius
120
BothWater
Water
and
Steam
and Steam
100
Steam
80
60
Water
40
20
0
Ice
Water
and Ice
-20
0
40
120
220
Pressure in mmHg or torr
760 800
Heating Curve for Water
Temperature in Celsius
120
Steam
Water and
Steam
100
80
60
Water
40
20
Ice and Water
0
IceWater and Ice
-20
0
40
120
220
Pressure in mmHg or torr
760 800
Heating Curve for Water
Temperature in Celsius
120
Steam
Water and
Steam
100
80
60
Water
40
20
0
Ice
Water
and Ice
-20
0
40
120
220
Pressure in mmHg or torr
760 800
Phase Diagram
1. Phase diagram – a graph of pressure vs.
temperature that shows in which phase a
substance exist under different conditions of
temperature and pressure.
2. Triple Point – the point on the phase diagram that
represents the temperature and pressure at which
all three phases can coexist
3. Critical Point – the point that indicates the critical
temperature and pressure. Critical temperature is
the temp. above which the sub. can’t exist in the
liquid state. Critical pressure is the lowest
pressure at which the substance can exist at the
critical temperature.
Phase Diagram of Water
16.2 Heat in Chemical
Reactions & Processes
Measuring Heat
• Calorimeter is an insulated
device used for measuring
the amount of heat
absorbed or released during
a chemical or physical
process
Chemical Energy & Universe
• System: specific part of universe
that contains reaction or process
• Surroundings: everything in the
universe other than the system
• Universe = system + surroundings
1. Enthalpy & Enthalpy changes
• Enthalpy: (H) heat content of a
system at constant pressure
• Can’t measure actual energy or
enthalpy of a substance, you can
measure change in enthalpy, which
is heat absorbed or released in a
chemical reaction
Enthalpy
• ∆Hrxn = Hproducts – Hreactants
• When ∆Hrxn is negative the reaction is
exothermic – Hproducts < Hreactants
• When ∆Hrxn is positive the reaction is
endothermic - Hproducts > Hreactants
2. Sign of enthalpy reaction
• Exothermic Reactions: heat pack
• Endothermic Reactions: cold pack
Practice Problem
Thermodynamic Heats of Formation
for one mole at 298K and 1 atmosphere pressure
• Substance (form) Enthalpy of
• Formation ΔHf (kJ)
• NaCl(s) -411.15
• Na+(aq) -240.12
• Cl-(aq) -167.16
•
Na+(aq) + Cl- (aq)  NaCl(s)
What is the change in enthalpy (ΔH) for this reaction?
• ∆Hrxn = Hproducts – Hreactants
• ∆Hrxn = [-411.15] – [(-240.12)+(-167.16)] = -3.87 kJ
•
So is the reaction Exothermic or Endothermic?
Endothermic & Exothermic Graphs
Hess’ Law
States that if you can add two or more thermochemical
equations to produce a final equation then the sum of the
enthalpy changes of the individual reactions is the
enthalpy change for the final reaction
Example:
Calculate the change in enthalpy for the following
reaction:
2S(s) + 3O2(g)  2SO3(g)
Given:
S(s) + O2(g)  SO2(g)
2SO3(g)  2SO2(g) + O2(g)
ëH = -297 kJ
ëH = 198 kJ