Thermochemistry

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Thermochemistry
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Definitions

• Energy capacity for doing work or supplying
  heat.
• Thermochemistry study of energy changes that
  occur during phase changes and chem. rxns.
• Chem. Potential Energy energy stored in
  chemical bonds.

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Example
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Heat

• Represented by q.
• Energy that transfers from one object to another
  because of a Temp. difference between them.
• Heat flows from warm ? cool until the two objects
  are at the same Temp.

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Exothermic vs. Endothermic

• In exothermic processes, the system loses heat as
  its surroundings warm up.
• q has a negative value b/c the system is losing
  heat.
• In endothermic processes, the system gains heat
  as its surroundings cool down.
• q has a positive value b/c the system is gaining
  heat.

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Potential Energy Diagram of Ice Melting at 0ºC.
Is the melting of ice an endothermic or an
exothermic process? How can you tell?
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Measuring Heat Flow

• SI Unit of heat flow Joule (J)
• Common unit used in chemistry calorie (cal)
• Amt. of heat needed to raise 1 gram of water by
  1ºC.
• 1 cal 4.184 J
• Food Calorie (capital C) 1000 cal, or 1
  kilocalorie 4184 J

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What do Calories mean in food?

• 10 grams of sugar has 41 Calories.
• When 10 grams of sugar are burned, 41 kcal (170
  kJ) of energy are released.
• Your body burns food for energy.
• In order to use the energy available in 10 grams
  of sugar, you must do 41 kcal worth of work.

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Heat Capacity

• Amount of heat needed to raise an objects
  temperature by 1C.
• Depends on the chemical composition and the mass
  of the object.
• EXAMPLE 1 gram of water requires 1 cal to raise
  its temperature by 1C.
• 100. g of water require 100. cal to raise the
  temp. by 1C.

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Heat Capacity
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Specific Heat (c)

• Amt. of heat needed to raise 1 gram of a
  substances temperature by 1ºC.
• Expressed in J/g ºC, or cal/g ºC
• The higher a substances specific heat, the more
  energy it takes to heat it.
• Substances with low specific heats heat up and
  cool down quickly (most metals, e.g.)

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Some Specific Heats
Substance Specific Heat Specific Heat
Substance J/gºC cal/g ºC
Water 4.18 1.00
Grain alcohol 2.4 0.58
Ice 2.1 0.50
Steam 1.7 0.40
Chloroform 0.96 0.23
Aluminum 0.90 0.21
Iron 0.46 0.11
Silver 0.24 0.057
Mercury 0.14 0.033
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Specific Heat (c)

• c heat / (mass x change in Temp.)
• c q / (m x ?T)
• q m x c x ?T

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Example Problem

• The temperature of a 95.4-g piece of Cu increases
  from 25.0ºC to 48.0ºC when the Cu absorbs 849 J
  of heat. What is the specific heat of Cu?
• SOLUTION q m x c x ?T
• 849 J (95.4 g) c (48.0ºC 25.0ºC)
• 849 J (95.4 g) c (23.0ºC)
• 849 J (2190 gºC) c
• c 0.388 J/gºC
• Based on what you know about metals, does this
  answer make sense?

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Example Problem

• When 435 J of heat is added to 3.4 g of olive oil
  at 21ºC, the temperature increases to 85ºC. What
  is the specific heat of olive oil?
• SOLUTION q m x c x ?T
• 435 J (3.4 g) c (85ºC 21ºC)
• 435 J (3.4 g) c (64ºC)
• 435 J (220 gºC) c
• c 2.0 J/gºC

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Example Problem

• How much heat is required to raise the
  temperature of 250.0 g of mercury by 52ºC? The
  specific heat of mercury is 0.14 J/gºC.
• SOLUTION q m x c x ?T
• q (250.0 g)(0.14 J/gºC)(52ºC)
• q 1800 J 1.8 kJ

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Enthalpy Changes

• Enthalpy (H) the heat content of a system at
  constant pressure.
• Enthalpy change (?H) the heat that enters or
  leaves a system at constant pressure.
• q ?H
• Neg. ?H exothermic process
• Pos. ?H endothermic process

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Thermochemical Equations

• Enthalpy change can be written as a reactant or a
  product.
• Reactant ? endothermic
• Product ? exothermic
• Example The reaction of calcium oxide with water
  is exothermic.
• It produces 65.2 kJ of heat per mole of CaO
  reacted.
• CaO(s) H2O(l) ? Ca(OH)2(s) 62.5 kJ

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An Exothermic Reaction
CaO(s) H2O(l) ? Ca(OH)2(s) 62.5 kJ
CaO(s) H2O(l)
?H -65.2 kJ
Ca(OH)2(s)
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Thermochemical Equations

• 2NaHCO3(s) 129 kJ ? Na2CO3(s) H2O(g) CO2(g)

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Thermochemical Equations and Stoichiometry

• You can use thermochemical equations in
  stoichiometry.
• How much heat energy is produced when 55.0 grams
  of ethanol is burned completely?
• C2H5OH(l) 3O2(g) ? 2CO2(g) 3H2O(g) 1300. kJ
• Given 55.0 g C2H5OH
• Want kJ
• Conversion factors
• 1 mol C2H5OH produces 1300. kJ when burned
• 1 mol C2H5OH 46.07 g/mol

55.0 g C2H5OH
1550 kJ
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Thermochemical Equations and Stoichiometry

• 0.500 grams of methane gas are burned completely
  beneath a container that holds 100. grams of
  water, originally at 20.0º. If all of the heat
  from the combustion reaction goes into the water,
  what will the waters final temperature be?
• CH4(g) 2O2(g) ? CO2(g) 2H2O(g) 803 kJ
• First find out how much total heat is released.
• Given 0.500 g CH4(g)
• Want kJ
• Conversion factors
• 1 mol CH4 16.05 g CH4
• 1 mol CH4 produces 803 kJ when completely burned

0.500 g CH4
25.0 kJ
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Thermochemical Equations and Stoichiometry

• The combustion of 5.00 grams of methane releases
  250. kJ of heat.
• Now well calculate how hot the water in the
  container will get if it absorbs all of the heat.
• First convert 25.0 kJ to J
• q m x c x ?T
• 2.50x104 J (100. g) (4.18 J/gºC) ?T
• 2.50x104 J (418 J/ºC) ?T
• ?T 59.8ºC
• The water will get 59.8ºC warmer.
• The final temperature will be 20.0ºC 59.8ºC
  79.8ºC.

25.0 kJ
2.50x104 J