Download Heat - Humble ISD

What is Energy?
• Energy is the ability to do work or produce
heat.
• Potential energy is energy due to
composition or position.
• Kinetic energy is energy of motion.
Law of Conservation of Energy
•The total energy of the universe is constant and
can neither be created nor destroyed; it can only
be transformed.
•Energy can’t be destroyed, but it can:
• Move or transfer to other locations
• Transform into a different type
Heat (represented by q)
• Form of energy
• Flows from a warmer object to a cooler object
• Heat is transferred until both objects are at the
same temperature
Measuring Heat
• A calorie is defined as the amount
of energy required to raise the
temperature of one gram of water
one degree Celsius.
• The energy content of food is
measured in Calories, or 1000
calories. Also known as a
kilocalorie.
• A joule is the SI unit of heat and
energy, equivalent to 0.2390
calories.
Relationships among
Energy Units
1 J = 0.2390 cal
1 cal = 4.184 J
1 Cal = 1 kcal
1 Cal = 4.184 kJ
Heat Conversions
• A breakfast of cereal, orange juice, and milk might
contain 230 Calories. Convert this energy to joules.
• A fruit and oatmeal bar contains 142 nutritional
Calories. Convert this energy to calories.
• An exothermic reaction releases 86.5 kJ of energy.
How many calories of energy are released?
Specific Heat
• The specific heat of any
substance is the amount of
heat required to raise one
gram of that substance one
degree Celsius.
Specific Heats at 298 K (25⁰C)
Substance
c in J/g°C
Aluminum
0.897
Bismuth
0.123
Copper
0.386
Brass
0.380
Gold
0.126
Lead
0.128
Silver
0.233
Tungsten
0.134
Zinc
0.387
Mercury
0.140
Alcohol(ethyl)
2.44
Water
4.186
Ice (-10 C)
2.05
Granite
0.803
Concrete
0.840
Q=mc∆T
• Tells us the heat that is gained or lost by a
system
• Q = heat gained or lost
• Measured in Joules
• M = mass of sample
• C = specific heat of sample
• ∆T = change in temperature (final temp-initial
temp)
• Must be measured in Celsius!
Example 1
A 500.g piece of Iron increases its temperature
7oC when heat energy is added. How much heat
energy produced this change in temperature?
Example 2
When 300. calories of energy is lost from a 125g
object, the temperature decreases from 45 oC to
40oC. What is the specific heat of this object?
What was the most likely identity of the object?
Example 3
A scientist wants to raise the temperature of a
10.0 g sample of aluminum from -45oC to 15oC.
How much heat energy is required to produce this
change in temperature?
System and Surroundings
• study of heat changes that accompany
chemical reactions and phase changes.
• Between the system and the surroundings
• The system is the specific part of the universe
that contains the reaction or process you wish
to study.
• The surroundings are everything else other
than the system in the
universe
Heat Transfer
• Heat is neither created nor destroyed; it is
transferred
•
•
Endothermic – system absorbs heat from the
surroundings
Exothermic – system
releases heat to the
surroundings
Endo (absorb heat?) or Exo (release
heat?)
Endothermic or exothermic process?
evaporating alcohol
leaves burning
boiling water
water cooling
melting ice
freezing water
Endo vs. Exo
• Endothermic
• Heat is a reactant
• Q is positve
• Exothermic
• Heat is a product
• Q is negative
Enthalpy
• Total heat content of a
system
• Heat of fusion, is the
amount of energy required
to melt one gram of a
substance
• Heat of vaporization, is
the amount of energy
required to boil one gram
of a substance
Calorimeter
• insulated device used
for measuring the
amount of heat
absorbed or released
in a chemical reaction
or physical process.
Phase Change Diagrams
Heat of Fusion
Heat of Vaporization
Increasing Temperature
Condensation
Evaporation
Freezing
Melting
Increasing Energy
Heat of fusionH2O = 334 J/g
Heat of vaporizationH2O = 2260 J/g
Ex. How much heat, in joules, is
needed to heat 150.0 g liquid water
from 0.0°C to 20.0OC?
• Happens in one phase
• Q=mcT
• Q=(150.0g)(4.184 J/g*°C)(20°C)
Ex. How much heat, in joules, is
needed to melt 150.0 g of ice at 0oC?
• Happening at one temp
• Phase change
• Q=mHf
• Q=(150.0g)(334 J/g)
Ex. How many joules are needed to
convert 5.0 g of ice at -15oC to steam at
130oC?
• Changing temps
• Changing phases
• Q=mcT
• Q=(5.0 g)(2.09 J/g°C)(15°C) – why 15°C???
• Q=(5.0 g)(334 J/g)
• Q=(5.0 g)(4.184 J/g°C)(100°C) – why 100°C???
• Q=(5.0 g)(2260 J/g)
• Q=(5.0 g)(2.01 J/g°C)(35°C)
Add
these
together
to get
total Q
Exothermic vs. Endothermic Recap
• Which type of reaction needs heat/energy?
•
Endothermic
• Which type of reaction releases heat/energy?
•
Exothermic
• Which reaction would you see the reactants
have a lower total energy than the products?
•
Endothermic
• Which reaction would you see the reactants
have a higher total energy than the products?
•
Exothermic
Energy
Diagrams
• Activation Energy - the minimum quantity of energy
that the reacting species must possess in order to
undergo a specified reaction
• Activated complex - the maximum energy point
along the reaction path
• H – difference in energy between the products and
the reactants (H = Hproducts – Hreactants)
Endothermic and Exothermic
Endothermic
Exothermic
Energy
Diagram
H
H is positive
H is negative
CH4(g) + 2O2(g)  CO2(g) +
Sample 2H2O + 241.8 kJ  2H2(g) + O2(g)
2H2O(g) + 890 kJ
Reaction ΔH = 241.8 kJ (energy is required) ΔH = -890 kJ (energy is released)
Where’s
the
energy?
Energy is a reactant
Energy is a product
Using Heat in Stoichiometry
• Can be used in reaction as reactant or product
(kJ/ mole)
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + 890 kJ
• In this equation, 890 kJ of heat are produced per the
amount of each reactant listed. How much energy is
produced if 5.0 L of oxygen were consumed?
1 𝑚𝑜𝑙𝑒 𝑂2
890 𝑘𝐽
5.0 L O2 x
x
= 99 kJ
22.4 𝐿 𝑂2 2 𝑚𝑜𝑙𝑒 𝑂2
Example 2
1049 kJ of energy are released when aluminum reacts
with hydrochloric acid
2 Al (s) + 6 HCl (aq)  2 AlCl3 (aq) + 3 H2 (g) + 1049 kJ
How many grams of aluminum are needed when 616 kJ
of energy are produced?
2 𝑚𝑜𝑙𝑒𝑠 𝐴𝑙 26.98 𝑔 𝐴𝑙
616 kJ x
x
= 31.7 g Al
1049 𝑘𝐽
1 𝑚𝑜𝑙𝑒 𝐴𝑙
Example 3
In the combustion of C2H6, 3120 kJ of energy are
produced. How much energy is produced in the
combustion of 21.50 grams of C 2H6?
2 C2H6 (g) + 7 O2 (g)  4 CO2 (g) + 6 H2O (l) + 3120 kJ
1 𝑚𝑜𝑙𝑒 𝐶2𝐻6
3120 𝑘𝐽
21.50 g C2H6 x
x
= 1115 kJ
30.08 𝑔 𝐶2𝐻6 2 𝑚𝑜𝑙𝑒 𝐶2𝐻6
Calculating H
H = Hproducts – Hreactants
Look up the enthalpies of each reactant and product
Sum of the products – sum of reactants
Ex. What is the H in the combustion of methane?
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g)
H = (-394 kJ + (2 x -286 kJ)) – (-75 kJ + (2 x 0 kJ)) = -891 kJ
Is the reaction endothermic or exothermic?
Example 2
How much energy is involved in the reaction between
magnesium and oxygen?
2 Mg (s) + O2 (g)  2 MgO (s)
H = (2 x -601.24 kJ) – ((2 x 0 kJ) + 0 kJ) = -1202.48 kJ
Example 3
How much energy is involved in the decomposition of
solid magnesium hydroxide?
Mg(OH)2 (s)
 H2O + MgO (s)
H = (-286 kJ + -601.24 kJ) - (-924.54 kJ) = 37.1 kJ
Hess’s Law
Hess’s law states that if you can add two or more
thermochemical equations to produce a final equation for a
reaction, then the sum of the enthalpy changes for the
individual reactions is the enthalpy change for the final
reaction.
Example 1
How much energy is involved in the decomposition of
hydrogen peroxide?
2 H 2O 2  2 H 2O + O 2
a
b
2 H 2 + O 2  2 H 2O
H2 + O 2  H2O 2
ΔH = -572 kJ
ΔH = -188 kJ
Since H2O2 is a reactant, we need to flip reaction b and
change the sign. What else do we have to do with equation
b?
Example 1
2 H 2O 2  2 H 2O + O 2
a
b
2 H 2 + O 2  2 H 2O
H2 + O 2  H2O 2
ΔH = -572 kJ
ΔH = -188 kJ
B becomes 2 H2O2  2 H2 + 2 O2
and ΔH = 188 kJ (x2) = 376 kJ
Add b to a to get total ΔH
-572 kJ + 376 kJ = -196 kJ
Example 2
Calculate the enthalpy of the following chemical reaction:
CS2 (l) + 3 O2(g)  CO2 (g) + 2 SO2 (g)
a
b
c
C (s) + O2  CO2 (g)
S (s) + O2  SO2 (g)
C (s) + 2 S (s)  CS2 (l)
ΔH = -393.5 kJ
ΔH = -296.8 kJ
ΔH = 87.9 kJ
Will we have to do anything to any of the reactions? Which
ones and what do we have to do?
Example 2
Calculate the enthalpy of the following chemical reaction:
CS2 (l) + 3 O2(g)  CO2 (g) + 2 SO2 (g)
a
b
c
C (s) + O2  CO2 (g)
2 S (s) + 2 O2  2 SO2 (g)
CS2 (l)  C (s) + 2 S (s)
Now add them all together!
ΔH = -393.5 kJ
ΔH = -593.6 kJ
ΔH = -87.9 kJ
Entropy (S)
• A measure of
randomness or
disorder
• Associated with
probability (There are
more ways for
something to be
disorganized than
organized.)
• Entropy increases going
from a solid to a liquid to
a gas.
• Entropy increases when
solutions are formed.
• Entropy increases in a
reaction when more
atoms or molecules are
formed.
• The entropy of a
substance increases with
temperature.
Entropy
Law of Thermodynamics
In any spontaneous
process there is always
an increase in the
entropy of the universe.
The energy of the
universe is constant but
the entropy of the
universe is increasing.
Spontaneous Process
• occurs without outside intervention
• may be fast or slow
When the
weather warms,
the melting of ice
is spontaneous.
The 3rd Law of Thermodynamics
• The entropy of a
perfect crystal at
0 K is zero.
• Soreaction =
Soprod- Soreact
• They are given in
joules/K.mol.