Transcript Heat - Midway ISD

(11) Science concepts. The student understands the energy changes
that occur in chemical reactions. The student is expected to:
(A) understand energy and its forms, including kinetic, potential,
chemical, and thermal energies;
(B) understand the law of conservation of energy and the processes of
heat transfer;
(C) use thermochemical equations to calculate energy changes that
occur in chemical reactions and classify reactions as exothermic or
endothermic;
(D) perform calculations involving heat, mass, temperature change,
and specific heat; and
(E) use calorimetry to calculate the heat of a chemical process.
Ch. 16
Energy and Chemical
Change
16.1 Energy
 Energy-
the ability to do
work or produce heat
 2 Forms:
 Potential
Energy
 Kinetic Energy
Potential Energy

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
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Potential Energy -energy due to the
composition or position of an object.
Ex: water stored behind a dam
depends on composition:
1. the type of atoms
2. the number and type of chemical
bonds joining the atoms
3. the way the atoms are arranged.
Kinetic Energy
 Kinetic
Energy – is the
energy of motion
 Ex: water flows from the
dam
 Chemical
systems contain
both potential and kinetic
energy
Potential
Kinetic

Heat- represented by symbol Qenergy that is in the process of
flowing from a warmer object to
a cooler object

Chemical Potential Energy -
the energy stored in a substance because
of its composition.
Composition is the type, number, and
arrangement of atoms and bonds.
Thermal energy


the energy created by moving
particles inside a substance.
more movement of particles =
more thermal energy
Heat is Thermal energy that is
transferred
Heat is Transferred in 3 ways
 Conduction – the way heat moves
through solids. (direct transfer)
Vibrating molecules pass on heat
from molecule to molecule.
Convection – the way heat moves
through gases and liquids.
Heated molecules move AWAY from the
heat and cooler molecules take their
place.
Ex: Hot air rises and cool air sinks

Radiation


Radiation – the way heat moves
through empty space.
Does not need atoms or molecules to
work.
Electromagnetic radiation –
light and heat from the sun,
visible light, microwaves, X-rays,
etc.
Forms of Energy
Wednesday
Phase Changes

http://www.youtube.com/watch?fea
ture=player_embedded&v=YG77v1
PwQNM
 Specific
Heat –is the
amount of heat required to
raise the temperature of one
gram of that substance by
one degree Celsius.
 each substance has its own
specific heat
 Table 16-2 pg 492
Heat of Vaporization

The amount of heat required to
convert unit mass of a liquid into
the vapor without a change in
temperature.
Heat of Fusion

The amount of heat required to
convert unit mass of a solid into
the liquid without a change in
temperature.
Measuring HEAT!!!
Two units for measuring heat
 calorie
- the amount of heat
required to raise the
temperature of one gram of pure
water by one degree Celsius
 Joule
energy
- SI unit of heat and
1 calorie = 4.184 joules
 1000 calorie = 1 Calorie



1J = 0.2390 calories
Table 16-1 Conversion factors and
relationships pg 491
 Calories
are nutritional or food
Calories
 1 Calorie = 1000 calories
 1Calorie = 1 kilocalorie
 approximates
the energy
needed to increase the
temperature of 1 kilogram of
water by 1 °C.
Calculating Specific Heat
Q = m x c x ΔT
Q
= heat absorbed or released
 m = mass of the sample in
grams
 c = specific heat of the
substance
 ΔT = difference between final
temperature and initial
temperature, or Tfinal- Tinitial
16.2 Heat in Chemical Reactions
and Processes
 Measuring

Heat
Heat changes are measured with a
calorimeter
Lab and worksheet
The temperature of a sample of iron
has a mass of 10.0g changed from
50.4oC to 25.0oC with the release of
114 J of heat. What is the specific
heat of iron?
Q = mc∆T
114 = 10 x c x (50.4-25)
114 = 254c
C = 114/254 = 0.449 J/goC


Calorimeter – an insulated device
used for measuring the amount of
heat absorbed or released during a
chemical or physical process.
Data is the change in temperature
of this mass of the substance.
Determining Specific Heat
Place a hot metal into water.
 Heat flows from the hot metal to
the cooler water until the
temperature of the metal and
water are equal.
 The heat gained by the water is
equal to the heat lost by the
metal

Calculating Heat
Example
125 g water with an Initial temperature of 25.60C
50 g metal at 1150C is placed in the water.
Heat flows from the hot metal to the cooler water
until the temperature of the metal and water are
equal. Both have a final temperature of 29.3 0C.
Calculate the Heat gained by the water.
Example Part A:
q
q
q
q
=cxm
water =
water =
water =
x /\T
4.184 J/(g x0C) x 125 g x (29.30C – 25.60C)
4.184 J/(g x0C) x 125 g X 3.7 0C
1900 J
Calculating Specific Heat



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

Example
50 g metal at 1150C is placed in the water.
Heat flows from the hot metal to the cooler water until the
temperature of the metal and water are equal. Both have a final
temperature of 29.3 0C.
Water absorbed 1900 J of heat.
Example Part B: Calculate the Specific Heat of the Metal
c=
q___
m x /\T
c metal = 1900 J
m x /\T
c metal = _______1900 J_________
(50.0 g)(1150C – 29.3 0C)
c metal = ____1900 J_____
(50.0 g)(85.700C)
c metal = 0.44 J/(g x 0C) specific heat of the metal
Look at pg 492 at the table. What is this metal?
Thursday- Lab
Friday- Practice worksheet
Monday
16.3 and 16.4 Enthalpy and
Enthalpy Changes
 Enthalpy-
(H) the heat
content of a system at a
constant pressure

A thermochemical equation is a
balanced chemical equation that
includes the physical states of all
reactants and products and the energy
change expressed as the change in
enthalpy, ∆H.





You can’t measure actual enthalpy, but
you can measure change in enthalpy,
which is called enthalpy (heat) of
reaction (ΔH rxn)
Use the table on pg. 510 in your
textbook
ΔH rxn = H final – H initial or
ΔH rxn = H products – H reactants
Example:

What is the heat of reaction for the
following reaction? H2S + 4F2  2HF + SF6
Endothermic Reaction


If the ∆H is shown on the reactants
side, it is endothermic (gaining
energy)
The heat of the reaction will be
positive.



(energy) 27 kJ + NH4NO3  NH4 + NO3
NH4NO3  NH4 + NO3 ΔH = +27 kJ
Energy required to break the bonds in a
reactant is less than released after the
bonds in the product is formed
Exothermic Reaction


If the ∆H is shown on the products
side, it is exothermic (losing
energy)
The heat of the reaction will be
negative.




4 Fe + 3O2  2 Fe2O3 + 1625 kJ (energy)
4 Fe + 3O2  2 Fe2O3
ΔH = -1625 kJ
Energy needed to break the bond in the
reactant is more than energy released
after the bonds in the products are
formed
http://www.youtube.com/watch?v=ksN-t2mmpvM&feature=related
END
Sign of the Enthalpy of Reaction
 Exothermic
reactions have a
negative enthalpy
 Hproducts < Hreactants
 Endothermic reactions have
a positive enthalpy
 Hproducts > Hreactants
16.3 Thermochemical Equations
 Enthalpy
(heat) of
combustion- enthalpy
change for the complete
burning of one mole of the
substance
 ΔHcomb
Entropy
 Measure
of the disorder or
randomness of the particles
that make up a system
 Symbolized by S
Molar Enthalpy (heat) of
Vaporization
 Heat
required to vaporize
one mole of a liquid
 ΔHvap
 Endothermic (positive
enthalpy)
Molar Enthalpy (heat) of Fusion
 The
heat required to melt
one mole of a solid
substance
 ΔHfus
 Endothermic (positive
enthalpy)
16.5 Reaction Spontaneity
 Spontaneous
processphysical or chemical change
that occurs with no outside
intervention
Law of Disorder
 States
that spontaneous
processes always proceed in
such a way that the entropy
of the universe increases
Chemical Energy and the Universe

Thermochemistry – the study of
heat changes that accompany
chemical reactions and phase
changes.


system – the specific part of the
universe that contains the reaction
or process you wish to study.
surroundings – everything in the
universe other than the system


universe – the system
plus the surroundings
universe = system +
surroundings

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Example: Using a heat pack to warm your
hands
Heat flows from the heat pack (the
system) to your cold hands (surroundings)
Exothermic - If energy is shown as a
product it means that heat is released.
The heat of the reaction will be
negative.
4 Fe + 3O2  2 Fe2O3 + 1625 kJ (energy)
4 Fe + 3O2  2 Fe2O3 Heat of rxn = -1625 kJ
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Example: Using a cold pack on an injured
knee
Heat flows from the knee (the
surroundings) to the cold pack (the
system)
Endothermic – If energy is shown as a
reactant it means that energy is absorbed.
The heat of the reaction will be
positive.
(energy) 27 kJ + NH4NO3  NH4 + NO3
NH4NO3  NH4 + NO3 Heat of rxn = 27 kJ