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Name: __________________________________ Date: ____________________
Unit 6: Energy
Aim: What is Energy?
Energy:
__________________________________________
Energy is required to bring about changes in matter (atoms, ions, or molecules).
Energy is measured in _____________ or _______________.
Conversion Practice:
It takes 2260 Joules to vaporize 1 gram of water.
Convert this quantity to kilojoules.
Law of Conservation of Matter & Energy:
______________________________________________
______________________________________________
Physical Changes
Example:
Chemical Changes
Example:
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Other forms of energy include potential & kinetic energy:
Which phase of matter has the highest kinetic energy? _________________
What is the relationship between heat and kinetic energy? _______________
___________________________________________________
Heat energy must be measured indirectly.
__________________________________________________
__________________________________________________
__________________________________________________
Temperature vs. Heat
What is the difference between temperature and heat?
_________________________________________________
__________________________________________________
__________________________________________________
Potential Energy
Kinetic Energy
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HEAT is a form of ENERGY.
_________________________________________________
_________________________________________________
__________________________________________________
Initial conditions Final conditions
How is heat energy transferred in each of the following situations?
In terms of energy transfer, what happens when you hold a snowball in your hands?
___________________________________________________________________
In terms of energy transfer, what happens when you place an ice pack on an injury?
___________________________________________________________________
TEMPERATURE is a measurement of the AVERAGE KINETIC
ENERGY of the particles in a sample. Units: _______ or ________ Lab tool: ________________
Which of the following has the highest kinetic energy?
a) Liquid mercury at 95oC
b) Solid mercury at -42oC
c) Gaseous mercury at 225oC
Water
90oC Water
50oC Water
50oC
Water
50oC
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Measuring Average Kinetic Energy: Celsius vs. Kelvin
Converting between Celsius & Kelvin
Formula (Table T): _____________
1. Convert the melting/freezing point of water from Celsius to Kelvin.
2. Convert the boiling/condensation point of water from Celsius to Kelvin.
3. 42oC = ____K
4. 450 K = ____oC
5. -45oC = ____K
Celsius Scale:
______________
______________
______________
______________
______________
______________
______________
______________
_____________
m.p./f.p = _____
c.p./b.p. = ____
Kelvin Scale:
______________
______________
______________
______________
______________
______________
______________
______________
_____________
m.p./f.p = _____
c.p./b.p. = ____
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Name: __________________________________ Date: ____________________Period:________
Aim:
What energy changes occur as substances are heated and cooled?
Do Now: Use Reference Table I to write the thermochemical equation for the
following reactions, including the heat energy value as a reactant or product:
a) Sodium hydroxide dissolves into its ions.
Thermochemical Equation:
____________________________________________________________
Circle the phrases that correctly describe the reaction above:
∆H is [positive / negative].
Heat is a [reactant / product].
Reaction is [endothermic / exothermic].
Reaction is a [physical / chemical] change.
b) The synthesis of nitrogen dioxide from its elements.
Thermochemical Equation:
____________________________________________________________
Circle the phrases that correctly describe the reaction above:
∆H is [positive / negative].
Heat is a [reactant / product].
Reaction is [endothermic / exothermic].
Reaction is a [physical / chemical] change.
c) The decomposition if nitrogen dioxide into its elements
Thermochemical Equation:
____________________________________________________________
Circle the phrases that correctly describe the reaction above:
∆H is [positive / negative].
Heat is a [reactant / product].
Reaction is [endothermic / exothermic].
Reaction is a [physical / chemical ] change.
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I. The System vs. The Surroundings
In chemistry, the world is divided into the system and the surroundings.
The _______________ is the part of the world we want to study.
The _____________________ consist of everything else outside the system.
II. Endothermic vs. Exothermic Processes
Exothermic Processes Endothermic Processes A process whereby the system
_________________________ to the surroundings.
A process whereby the system
_________________________ from the surroundings.
The heat value (∆H) is ____________.
The heat value (∆H) is ____________.
Heat is a _________________.
Heat is a ___________________.
Thermochemical Equation Example:
Thermochemical Equation Example:
Exothermic reaction Endothermic Reaction
III. Heat & Changes of State
_________________ is a required part of any phase change.
Determine whether the following phase changes are endothermic or exothermic.
1. Melting: ___________ 3. Evaporation:___________ 5. Sublimation: _________
2. Freezing: __________ 4. Condensation: ___________ 6. Deposition: __________
IV. Interpreting Phase Change Diagrams States of matter exist on the sloped lines. Kinetic energy (temperature) is changing
while potential energy is unchanging.
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Phase changes exist on the plateaus (unsloped lines). Potential energy is changing while
kinetic energy (temperature) is unchanging.
Heating Curve of Water
Analyze your phase change diagram by answering the following questions:
1. Is this heating curve endothermic or exothermic? _________________
2. On which line segments is potential energy increasing? _________________
3. On which line segments is kinetic energy increasing? _________________
4. At which point does the highest kinetic energy exist? _________________
5. Which point represents the melting point? _________________
6. Which point represents the evaporation point? _________________
7. Identify the line segments that represent more than one phase of matter existing at once.
_________________
IV. Interpreting Phase Change Diagrams States of matter exist on the sloped lines. Kinetic energy (temperature) is changing
while potential energy is unchanging.
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Phase changes exist on the plateaus (unsloped lines). Potential energy is changing while
kinetic energy (temperature) is unchanging.
Cooling Curve of Water
Analyze your phase change diagram by answering the following questions:
1. Is this cooling curve endothermic or exothermic? _________________
2. On which line segments is potential energy decreasing? _________________
3. On which line segments is kinetic energy decreasing? _________________
4. At which point does the highest kinetic energy exist? _________________
5. Which point represents the freezing point? _________________
6. Which point represents the condensation point? _________________
7. Identify the line segments that represent more than one phase of matter existing at once.
_________________
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Name: __________________________________ Date: ____________________
Aim:
How can we calculate the heat changes that occur during physical
and chemical changes?
Do Now: Interpret the following heating/cooling curves by answer the questions that follow.
1. The following graph is a heating curve showing the addition of heat at a constant rate
of 500.0 joules/minute to a 3.00 gram sample of ice at –20.0°C. The final temperature
of the vapor is 140.0°C.
a. During which segments is kinetic energy increasing? _________________________
b. During which segments does kinetic energy remain the same? ___________________
c. During which segments is potential energy increasing? ________________________
d. During which segments does potential energy remain the same? _________________
e. At what time does the liquid phase first appear? ____________________________
f. During which segment is the substance entirely in the solid state? _______________
g. How long is the solid state in equilibrium with the liquid state? _________________
h. How long does it take to raise the temperature form the melting point to the boiling
point? _________________
i. At what time does the gas phase first appear? ____________________________
j. How long does it take to vaporize the sample at its boiling point? ________________
k. How long does it take to completely melt the sample at its melting point? __________
l. How many joules are required to melt the sample at its melting point?** ___________
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2. The following is a cooling curve showing the release of heat at a constant rate of 500.0
joules/minute from a 3.00 gram sample of water vapor at 140.0°C. The final temperature of the ice
is –20.0°C.
a. During which segments is kinetic energy decreasing? _________________________
b. During which segments does kinetic energy remain the same? ___________________
c. During which segments is potential energy decreasing? ________________________
d. During which segments does potential energy remain the same? _________________
e. During which segments is one phase only present? ___________________________
f. During which segments are two phases present? _______________________
g. At what time does the liquid phase first appear? _______________________
h. At what time does the solid phase first appear? _______________________
i. At what time do the particles have the highest average kinetic energy? ___________
j. At what time does the gas phase first appear? ____________________________
k. How long does it take to condense the sample at its condensation point? ___________
l. How long does it take to completely freeze the sample at its freezing point? _______
m. During which segment is the substance entirely in the solid state? _____________
n. During which segment is the substance entirely in the liquid state? ___________
o. During which segment is the substance entirely in the gas state? ____________
p. How long is the solid state in equilibrium with the liquid state? ______________
The temperature of the sample at this point is _________________ Kelvin.
q. 18. How long is the liquid state in equilibrium with the gas state? ______________
The temperature of the sample at this point is _______________________ Kelvin.
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Aim:
How can we calculate the heat changes that occur during physical
Specific Heat Capacity, C
___________________________________________________________________
___________________________________________________________________
The specific heat capacity for H2O(l) is ______________ (Table B).
Calculating how much energy is involved with changing the temperature
of a given sample at a given mass:
Try This:
How much heat is needed to bring 25 grams of water from an initial temperature of 20.0oC to
65.0 oC? Convert your final answer to kJ.
q =
m =
C =
∆T =
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LET’S PRACTICE!
1. What is the specific heat of silver if a 93.9 g sample cools from 215.0 oC to 196.0 oC with
the loss of 428 J of energy?
q =
m =
C =
∆T =
2. What is the change in heat energy when 64.82 g of aluminum metal at 100.0oC is cooled to
82.0 oC? The specific heat of aluminum is 0.897 J /g °C.
q =
m =
C =
∆T =
3. An unknown substance with a mass of 140 g is cooled from 98.4oC to 62.2oC with a release
of 1,137 J of heat. From this data, calculate the specific heat of the unknown substance.
q =
m =
C =
∆T =
4. If 100.0 J are added to 20.0 g of water at 30.0 oC, what will be the final
temperature of water?
q =
m =
C =
∆T =
5. A 1.6 g sample of a metal that looks like gold requires .0058 kJ of energy to change the
temperature from 23 0 C to 41 0 C. Is the metal pure gold?
q =
m =
C =
∆T =
Final Thought:
Why can’t be used to calculate heat change during phase changes (plateaus)?
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Heat of Fusion, Hf
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
The heat of fusion for water is ______________ (Table B).
Calculating how much energy is required to melt a given mass of a given sample:
Examples:
1. How much heat is required to melt 4 g of ice at 0oC?
q =
m =
Hf =
2. How much heat is required in kJ to melt 55.0 g of nickel at its melting point? The heat of
fusion of nickel is .296 kJ/g.
q =
m =
Hf =
3. How much heat is transferred when 105 g of ethanol is frozen from a liquid to its solid
state? The heat of fusion of ethanol is 108.99 J/g.
q =
m =
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Hf =
Heat of Vaporization, Hv
__________________________________________________________
__________________________________________________________
The heat of vaporization for water is ______________ (Table B).
Calculating how much energy is required to melt a given mass of a given sample:
Examples:
1. How much heat is required to turn 12 g of water at 100oC into steam at the same
temperature?
q =
m =
Hv =
2. How much energy is released to the environment by 50.0 grams of condensing water vapor?
q =
m =
Hv =
3. You have 250,000 J of heat available to turn 0.12 kg of water at 100oC into steam. Is it
enough?
q =
m =
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Hf =
FINAL PRACTICE – Mixing It Up! 1. How many kJ are required to completely vaporize a 25.0 gram sample of water at 100°C?
2. How many kJ are required to change 10.0 grams of ice to water at 0°C?
3. If 11.9 kJ are used to heat a sample of water the temperature increases from 20.0°C to
65.0°C. Calculate the mass of the water.
4. What is the maximum mass of water that can be heated from 25oC to 30o by the addition
of 1254 joules of heat?
5. The heat of vaporization of isopropyl alcohol is 714 J/g. How many kJ must be added to an
80.0 gram sample to completely vaporize it at its boiling point?
6. How many kJ must be added to an 80.0 gram sample of water to completely vaporize it at
100. 0°C.
7. Compare your answers to questions 5 and 6. Which substance, isopropyl alcohol or water,
has stronger intermolecular forces of attraction? Explain.
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