Transcript Energy and Chemical Change

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Energy and Chemical Change
Measuring and calculating the
energy involved in chemical
changes
Energy
• Energy is defined
as the ability to
do work or
produce heat.
• SI unit for energy
is Joules (J).
• There are 2
types of energy
Potential
Energy
Kinetic
Energy
Energy
1. Potential Energy (PE)
– Energy that is based on an objects
composition (chemical PE) or position;
associated with attractions and repulsions.
– Chemical potential energy
• the energy stored in a substance because
of its composition (ex: plants, gasoline)
• plays an important role in chemical
reactions
Energy
2. Kinetic Energy (KE)
- Energy of motion.
• Kinetic Energy
• Can be measured as temperature or
heat.
• KE of a substance is directly related to
the motion of its atoms and
temperature.
Energy
• * Every object has either potential or
kinetic energy!!!
Energy
• Law of Conservation of Energy:
energy can be converted from one
form to another, but it is neither
created nor destroyed.
Heat
1. Form of energy associated with
changing the temperature of an
object.
2. An object’s temperature increases
because energy is transferred into it
Temperature is a measure of the average
kinetic energy of the particles in a
sample of matter; it does not depend
on the amount of matter in the sample.
Heat
• The heat (energy) transfer occurs when
2 objects of different temperatures
are brought into contact to reach
equilibrium.
• Depends on the mass of the sample.
• Symbol = m
Unit = grams
• Reactions can be endothermic
(heat/energy in) or exothermic
(heat/energy out).
How Can I Measure
Energy/Heat?
There are many units that can
express an amount of energy.
1. calorie (cal) - amount of heat
required to raise the temperature
of 1.00 g of pure water by 1oC.
How Can I Measure
Energy/Heat?
2. kilocalorie (kcal) - 1000 calories or one
Calorie – this is one food Calorie!
3. Joule - SI unit of heat and energy.
(What we will mainly use)
Important conversions:
1 Calorie = 1 kcal
1 kcal = 1000 cal
1 cal = 4.184 J
How Can I Measure
Energy/Heat?
• heat changes that occur during chemical
and physical processes can be measured
accurately and precisely using a
calorimeter
– a calorimeter is an insulated device used
for measuring the amount of heat absorbed
or released during a chemical or physical
process
How Can I Measure
Energy/Heat?
• Inside most calorimeters you will need:
– a known mass of water is placed in an
insulated chamber in the calorimeter
– it is designed to absorb the energy
released from the reacting system
(unknown/known substance) or to provide
the energy absorbed by the system
– the data collected is the change in
temperature of the water inside the
calorimeter
How Can I Measure
Energy/Heat?
– most calorimetric reactions occur at
constant pressure
– can also be used to determine the
specific heat of an unknown metal
– assuming that no heat is lost to the
surroundings, the heat lost by the
metal is equal to the heat gained by
the water
What is Specific Heat? (Or, How
do I Calculate Amount of Energy
Transferred?)
– Specific Heat is the amount of heat
required to raise the temperature of one
gram of any substance by 1oC.
• Units = J/g x OC of kJ/g x oC
– Each substance has its own specific heat.
– Knowing the specific heat of a substance
allows you to calculate the heat
transferred by measuring the mass of the
substance and the temperature change.
What is Specific Heat? (Or, How
do I Calculate Amount of Energy
Transferred?)
–
–
–
–
–
Specific heat of water = 4.184 J/ g x oC
Formula:
q = c x m x ∆T
c = specific heat
m = mass
∆T = Tfinal – Tinitial
• Heat transferred = (specific heat)*
(mass)* (change in T measured in OC )
• q = (g)*(J/g x OC)*(∆T)
What is Specific Heat? (Or, How
do I Calculate Amount of Energy
Transferred?)
--This equation also indicates which direction
the heat is traveling.
**If ∆T is (+) then heat is being
transferred to or absorbed by the
substance (endothermic)
**If ∆T (-) then heat is being released
during the reaction (exothermic)
Keep in mind…
• the specific heat of a
substance is a measure of
how efficiently that
substance absorbs heat
PRACTICE!!!
Problem- How much heat is lost when solid
Al with a mass of 4110 g cools from
660.0 oC to 25 oC?
• (Specific Heat of Al is 0.902 J/g x oC)
• First determine the ΔT = Tfinal -Tinitial
(25oC-660oC = -635oC)
• Remember the equation is q= cxmxΔT
• c = specific heat
• m = mass
PRACTICE!!!
• Plug in the values
• q = (0.902 J/gxoC)*(4110g)*(-635oC)
• Cancel like units
• Multiply all values
• -2354084.7J or divide by 1000 and get
-2354.0847kJ
• Significant figures gives us -2354kJ or
-2.35 x 106J
• Type of reaction: EXOTHERMIC
PRACTICE!!!
Problem- How much heat is required to
raise the temperature of 854 g of H2O
from 23.5 oC to 85.0 oC?
ΔT = Tfinal – Tinitial
ΔT = 85.0oC – 23.5oC = 61.5oC
q = c x m x ΔT
q = (4.184 J/g x oC)*(854g)*(61.5oC)=
219747.9J or 219.7kJ or 2.20 x 105J
Type of reaction: ENDOTHERMIC
So, What is Thermochemistry?
• Thermochemistry is the study of heat
changes that accompany chemical
reactions and phase changes.
***Remember: Energy often takes the
form of HEAT
• As a chemical reaction progresses some
of the energy in the bonds may be
released as heat. These reactions are
called EXOTHERMIC.
So, What is Thermochemistry?
 In an exothermic reaction the products
have less potential energy than the
reactants.
 Exothermic reactions release energy;
more energy is released than is required
to break bonds in the initial reaction.
So, What is Thermochemistry?
 Examples: burning wood, any reaction in which
the test tube or beaker gets warm, formation
of water (explosion of hydrogen and oxygen),
an athletic hot pack.
o the contents of the athletic hot pack is the
system which is the specific part of the universe
that contains the reaction or process you wish to
study
o the universe is defined as the system plus the
surroundings (which is everything in the universe
other than the system)
So, What is Thermochemistry?
 Some chemical reactions absorb
heat as they progress. These
reactions are called
ENDOTHERMIC.
 In an endothermic reaction the
products have more heat energy
(enthalpy) than the reactants.
So, What is Thermochemistry?
 Reaction that needs energy to complete
the reaction; a greater amount of
energy is required to break the initial
bonds than is released from the new
bonds forming
 Examples: cooking eggs, any reaction in
which the beaker or test tube gets
cooler, an athletic cold pack,
decomposition of water (electrolysis).
Quick Quiz!!!
• Identify each of the following as
endothermic or exothermic:
• a.
• b.
• c.
burning natural gas (methane)
exothermic
baking a cake
endothermic
melting ice
endothermic
What is Enthalpy?
 Enthalpy is the chemistry word for heat
content (remember that chemical
compounds have POTENTIAL energy in
their bonds – this is the heat content).
– Enthalpy is represented by the symbol H
– You cannot measure the actual enthalpy of
a substance…instead you measure the
change in enthalpy (the heat absorbed or
released in a reaction)
What is Enthalpy?
• The change in enthalpy is called
the enthalpy heat of reaction
and is represented by ∆H or
∆Hrxn.
• ΔHrxn = Hproducts – Hreactants
Enthalpy of Exothermic
Reactions
 Enthalpy changes for
exothermic
reactions are always
negative. ∆H is (-)
 In an exothermic
reaction the
products have less
potential energy
than the reactants
 Hprod. < Hreact.
Enthalpy of Exothermic
Reactions
• Energy is a product
• Example:
• CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
+ 802 kJ
product
• CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
ΔH = -802 kJ (because it is on the
product side)
Enthalpy of Endothermic
Reactions
• Enthalpy changes for
endothermic
reactions are always
positive. ∆H is (+)
• In an endothermic
reaction the
products have more
heat energy
(enthalpy) than the
reactants.
• Hprod > Hreact
Enthalpy of Endothermic
Reactions
•
•
•
•
•
Energy is a reactant
Example:
2H2O(l) + 572 kJ  2H2(g) + O2(g)
2H2O(l)  2H2(g) + O2(g)
ΔH = 572 kJ (because it is on the
reactant side)
Thermochemical Equations
• Thermochemical equations are balanced
chemical equations that include the
physical states of all reactants and
products and the energy change.
• 2H2O(l)  2H2(g) + O2(g) ΔH = 572 kJ
• CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
ΔH = -802 kJ
Enthalpy of Combustion
• Enthalpy heat of combustion (ΔHcomb) is
the enthalpy change for the complete
burning of one mole of the substance.
-- carried out under standard conditions
which are one atmospheric pressure (1
atm) and 298K (250C)
- C6H1206(s) + 6O2(g)6CO2(g)+ 6H2O(l)
ΔHcomb = -2808kJ
Changes of State
• Molar enthalpy (heat) of vaporization
(ΔHvap) is the heat required to vaporize
one mole of liquid.
-- think of water vaporizing from your
skin after you take a hot shower. Your
skin provides the heat needed to
vaporize the water and as the water
absorbs the heat you feel cool (shiver)
ΔHvap = -ΔHcond (condensation)
Changes of State
• Molar enthalpy (heat) of fusion (ΔHfus)
is the heat required to melt one mole of
a solid substance.
--think of ice in a drink. The drink cools
as it provides the heat for the ice to
melt
ΔHfus = - ΔHsolid (solidification—freezing)
Phase Changes
• Occurs when energy is added or removed
from a system and the substance can go
from one physical phase to another
Endothermic Phase Changes
Melting
• The energy absorbed to melt a solid is
not used to raise the temperature of
that solid
• The energy instead disrupts the bonds
holding the solid’s molecules together
and cause the molecules to move into
the liquid phase
Endothermic Phase Changes
• The amount of energy required to melt
one mole of a solid depends on the
strength of the forces that hold the
solid together
• The melting point of a crystalline solid
is the temperature at which the forces
holding its crystal lattice together are
broken and it becomes a liquid
Endothermic Phase Changes
Vaporization
• Particle that escape from the liquid enter the
gas phase and those liquids at room
temperature the gas phase is called vapor
• Vaporization is the process by which a liquid
changes into a gas or vapor
• Once the solid becomes a liquid then and only
then does the temperature of the substance
begin to increase
Endothermic Phase Changes
• When vaporization takes place only at
the surface of the liquid it is called
evaporation
• Evaporation is the method by which the
human body maintains and controls its
temperature
Endothermic Phase Changes
• The pressure exerted by a vapor over a
liquid is called the vapor pressure
• The temperature at which the vapor
pressure of a liquid equals the external
or atmospheric pressure is called the
boiling point
Endothermic Phase Changes
Sublimation
• Is the process by which a solid
changes directly to a gas
without first becoming a liquid
• Dry ice (CO2) and snow are the
most common examples
Endothermic Phase Changes
• If ice cubes are left in the
freezer for extended periods
of time, they will eventually
sublime and become smaller
– This process is also helpful in
freeze drying foods for hikers
and astronauts
Exothermic Phase Changes
Condensation
• When a vapor molecule loses energy its
velocity is reduced therefore colliding
more with other molecules to form a
liquid
• Condensation is the process by which a
gas or vapor becomes a liquid and it is
the reverse action of vaporization
Exothermic Phase Changes
Deposition
• Is the process by which a substance
changes from a gas or vapor to a solid
without first becoming a liquid
• It is the reverse action of sublimation
• The formation of snow crystals high up
in the atmosphere is an example
Exothermic Phase Changes
Freezing Point
• Is the temperature at which a
liquid is converted into a
crystalline solid
• The same temperature as the
melting point of a given
substance
Phase Change Graph
Phase Change Graph
• Graph shows the energy required to
go from one phase to the other
• Where the graph inclines, potential
energy is at its greatest and
temperature is increasing
• Where the graph plateaus (flat
region) kinetic energy is at its
greatest but the temperature
remains constant
Phase Diagrams
• A phase diagram is a graph of pressure
versus temperature that shows in which
phase a substance exists under different
conditions of temperature and pressure
Phase Diagrams
 The triple point is the point on a phase
diagram
that
represents
the
temperature and pressure at which
three phases of a substance can coexist
 The critical point is the point that
indicates
critical
pressure
and
temperature above which water cannot
exist as a liquid
Phase Diagrams
• Different for
each substance
because of the
different
boiling/freezing
points