The Nature of Energy

1
Section 15-1
The Nature of Energy

• Energy is the ability to do work or produce heat.
• weightless, odorless, tasteless

• Two forms of energy exist, potential and kinetic.
• Potential energy is due to composition or
  position.
• Chemical potential energy is energy stored in a
  substance because of its composition.
• Kinetic energy is energy of motion.

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Section 15-1
The Nature of Energy (cont.)

• The 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.
• Heat is energy that is in the process of flowing
  (transferring) from a warmer object to a cooler
  object.
• q is used to symbolize heat.

3
Endothermic and Exothermic Processes

• Essentially all chemical reactions and changes in
  physical state involve either
• release of heat, or
• absorption of heat

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Endothermic and Exothermic Processes

• In studying heat changes, think of defining these
  two parts
• the system - the part of the universe on which
  you focus your attention
• the surroundings - includes everything else in
  the universe
• Together, the system and its surroundings
  constitute the universe

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Endothermic and Exothermic Processes

• Heat flowing into a system from its
  surroundings
• defined as positive
• q has a positive value
• called endothermic
• system gains heat (gets warmer) as the
  surroundings cool down

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Endothermic and Exothermic Processes

• Heat flowing out of a system into its
  surroundings
• defined as negative
• q has a negative value
• called exothermic
• system loses heat (gets cooler) as the
  surroundings heat up

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Section 15-1
Measuring Heat

• A calorie is defined as the amount of energy
  required to raise the temperature of one gram of
  water one degree Celsius.

• Food is measured in Calories, or 1000 calories
  (kilocalorie).
• A joule is the SI unit of heat and energy,
  equivalent to 0.2390 calories.
• 1 calorie 4.184 J or 1 J 0.2390 calories

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Section 15-1
Measuring Heat (cont.)
9

• Example
• A candy bar has 245 Calories. Convert this to
  calories and then to Joules of energy.

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Section 15-1
Specific Heat

• The specific heat of any substance is the amount
  of heat required to raise one gram of that
  substance one degree Celsius.

• Some objects require more heat than others to
  raise their temperature.

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Section 15-1
Specific Heat (cont.)

• Calculating heat absorbed and released

• q c m ?T
• q heat absorbed or released (in Joules)
• c specific heat of substance
• m mass of substance in grams
• ?T change in temperature in Celsius

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Specific Heat (cont.)

• Examples
• How much heat does a 20.0 g ice cube absorb as
  its temperature increases from (-27.0oC) to
  0.0oC? Give your answer in both joules and
  calories.
• q c m ?T
• Specific Heat of Ice 2.03 J/goC
• 1 calorie 4.184 J

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Specific Heat (cont.)

• Example Cont.
• q ?
• c 2.03 J/goC
• m 20.0 grams
• ?T FinalTemp(0.0oC) InitialTemp (-27.0oC)
  Change (27.0oC)
• q c m ?T
• q (2.03 J/goC)(20.0g)(27.0oC)

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• Example 2
• A 5.00 gram sample of a metal is initially at
  55.0 ºC. When the metal is allowed to cool for a
  certain time, 98.8 Joules of energy are lost and
  the temperature decreases to 11.0º C. What is the
  specific heat of the metal? What metal is it?
• q c m ?T

To make the problem easier, solve for the unknown
BEFORE you plug in the numbers.
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Measuring Heat

• For Water during a phase change
• The Heat of Fusion (melting) is 334 j/g
• The Heat of Solidification (freezing) is 334 j/g
• They are the same value (energy in or out)
• The Heat of Vaporization is 2260 j/g
• The Heat of Condensation is 2260 j/g
• They are the same value (energy in or out)

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• Example Phase change
• Calculate the amount of energy needed to convert
  55.0 grams of ice to all liquid water at its
  normal melting point.
• Using the same amount of water calculate the
  energy needed to completely vaporize the water at
  its normal boiling point.
• Why is there such a large difference in energy
  needed?

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120
The liquid is boiling at 100o C no temperature
change (use q mass x ?Hvap.)
The gas temperature is rising from 100 to 120 oC
(use q mass x ?T x C)
The Heat Curve for Water, going from -20 to 120
oC,
The liquid temperature is rising from 0 to 100 oC
(use q mass x ?T x C)
The solid is melting at 0o C no temperature
change (use q mass x ?Hfus.)
The solid temperature is rising from -20 to 0 oC
(use q mass x ?T x C)
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Section 15-2
Calorimetry

• Calorimetry - the measurement of the heat into or
  out of a system for chemical and physical
  processes.
• A calorimeter is an insulated device used for
  measuring the amount of heat absorbed or released
  in a chemical reaction or physical process.
• Based on the fact that the heat released the
  heat absorbed

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Section 15-2
Chemical Energy and the Universe (cont.)

• Chemists are interested in changes in energy
  during reactions.

• Enthalpy is the heat content of a system at
  constant pressure.
• Enthalpy (heat) of reaction is the change in
  enthalpy during a reaction symbolized as ?Hrxn.
• ?Hrxn Hfinal Hinitial
• ?Hrxn Hproducts Hreactants

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Changes in enthalpy ?H q ?H These terms
will be used interchangeably in this
textbook Thus, q ?H m x C x ?T ?H is
negative for an exothermic reaction ?H is
positive for an endothermic reaction
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Section 15-2
Chemical Energy and the Universe (cont.)
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Section 15-2
Chemical Energy and the Universe (cont.)
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Exothermic

• The products are lower in energy than the
  reactants
• Thus, energy is released.
• ?H -395 kJ
• The negative sign does not mean negative energy,
  but instead that energy is lost.

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Endothermic

• The products are higher in energy than the
  reactants
• Thus, energy is absorbed.
• ?H 176 kJ
• The positive sign means energy is absorbed

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Chemistry Happens in

• MOLES
• An equation that includes energy is called a
  thermochemical equation
• CH4 2O2 ? CO2 2H2O 802.2 kJ
• 1 mole of CH4 releases 802.2 kJ of energy.
• When you make 802.2 kJ you also make 2 moles of
  water and 1 mole of CO2

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Thermochemical Equations

• The heat of reaction is the heat change for the
  equation, exactly as written
• The physical state of reactants and products must
  also be given.
• Standard conditions (SC) for the reaction is (1
  atm.) and 25 oC (different from STP)

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CH4(g) 2 O2(g) ? CO2(g) 2 H2O(l) 802.2 kJ
1

• If 10.3 grams of CH4 are burned completely, how
  much heat will be produced?

Convert moles to desired unit
Convert to moles
Start with known value
1 mol CH4
802.2 kJ
10. 3 g CH4
16.05 g CH4
1 mol CH4
514 kJ
Ratio from balanced equation
?H -514 kJ, which means the heat is released
for the reaction of 10.3 grams CH4