Chapter 17: Thermochemistry

Thermochemistry is the study of

______= a type of chemical energy, sometimes referred to as “heat content”, ΔH (the heat of reaction for a chemical reaction)

  • Endothermic reactions: absorbs heat from surroundings (+) H______

If you touch an endothermic reaction it feels ______

  • Exothermic reactions: release heat to the surroundings (-) H ______

If you touch an exothermic reaction it feels______

  • Units of Heat Energy:

1 kcal = 1,000 cal = 1 Cal (nutritional)

1 kJ = 1,000 J

1 calorie = 4.184 J

1 kcal = 4.184 kJ

A chemical equation that shows the enthalpy (H) is a ______

Rule #1: The magnitude (value) of H is directly proportional to the amount of reactant or product. H2 + Cl22HCl H = - 185 kJ

* meaning there are 185 kJ of energy RELEASED for every:

Example 1: H2 + Cl22HClH = - 185 kJ
Calculate H when 2.00 moles of Cl2 reacts.
Example 2: Methanol burns to produce carbon dioxide and water. What mass of methanol is needed to produce 1820 kJ?
2CH3OH + 3O2 2CO2 + 4H2O H = - 1454 kJ
Rule #2:H for a reaction is equal in the magnitude but opposite in sign to H for the reverse reaction. (If 6.00 kJ of heat absorbed when a mole of ice melts, then 6.00 kJ of heat is given off when 1.00 mol of liquid water freezes)
Example 1:
Given: H2 + ½O2 H2OH = -285.8 kJ
Reverse: H2O  H2 + ½O2H = +285.8 kJ
Example 2:CaCO3 (s)  CaO (s) + CO2 (g) H = 178 kJ
What is the H for the REVERSE RXN?
CaO (s) + CO2 (g)  CaCO3 (s) H = ______
Alternate form of Thermochemical equation - Putting the heat content of a reaction INTO the actual thermochemical eq.
H2 + ½O2 H2OH = -285.8 kJ
Alternate form:
Example: 2 NaHCO3 + 129 kJ  Na2CO3 + H2O + CO2
Put in the alternate form:
Put the following in alternate form:
  • H2 + Cl2 2 HCl H = -185 kJ
  • 2 Mg + O2 2 MgO + 72.3 kJ
  • 2 HgO  2 Hg + O2H = 181.66 kJ

Enthalpies of Formation

  • usually exothermic
  • see table for Hf value (table given)
  • enthalpy of formation of an element in its stable state = 0
  • these can be used to calculate H for a reaction

To calculate Heat of Formation:

Basically:

Heat of reaction = (add upHf for all the products) – (add up Hf for all the reactants)

Example: elements in their standard states can be omitted:

2 Al(s) + Fe2O3(s)  2 Fe(s) + Al2O3(s)

Why was Al(s) & Fe(s) crossed off?

Example 2: The coefficient of the products and reactants in the thermochemical equation must be taken into account:

O2(g) + 2SO2 (g)  2SO3 (g)

Example 3: Calculate the standard heat for formation of benzene, C6H6, given the following thermochemical equation:

C6H6(l) + 15/2 O2(g)  6CO2(g) + 3H2O(l) H = -3267.4 kJ

Specific Heat

Heat (q) is ______that transfers from one object to another because of ______

  • Heat ALWAYS flows from a ______object to a ______one.
  • Heat moves between the system (reaction) and the surroundings
  • Must obey the law of conservation of energy (______)
  • Thermochemical equations tell you the direction of heat flow by the “sign”, + or –
  • Endothermic reactions: absorbs heat from surroundings (+).
  • If you touch an endothermic reaction it feels COLD
  • Exothermic reactions: release heat to the surroundings (-)
  • If you touch an exothermic reaction it feels HOT
  • UNIT of energy = JOULE (J)

Heat capacity is

*** heat capacity of an object depends on both its mass and its chemical composition

  • The greater the mass, the greater its heat capacity
  • Iron Bar: ______SPECIFIC HEAT. Heats up ______, only a ______of energy is needed to raise the temperature of the iron bar.
  • Wooden Bar: ______Specific Heat. Heats up ______. ______is needed to raise the temperature of the wooden bar.

SPECIFIC Heat Capacity (specific heat, c)

  • specific to a substance
  • Units = ______

Specific heat values (in J/gC):

  • CO2(g) = 0.843 J/gC
  • Cu(s) = 0.382 J/gC
  • Fe(s) = 0.446 J/gC
  • H2O (l) = 4.184 J/gC

CALCULATING HEAT - You can calculate how much heat is needed to raise the temperature of a given amount of substance

Where:

q = heat (Joules)

m = mass (grams)

c = specific heat (J/gC) T = change in temperature (C) Tf-Ti

Example 1: How much heat is given off by a 50.0 g sample of copper when it cools from 80.0 to 50.0C?

Example 2: Iron has a specific heat of 0.446 J/gC. When a 7.55 g piece of iron absorbs 10.33 J of heat, (A) what is the change in temperature? (B) If it was originally at room temp. (22.0C), what is the final temperature?

(A) T = ?(B) Tf = ?

Example 3: A metal plate originally at 25.0oC and a mass of 135.5 g absorbs 9,870 kJ of heat when placed in a 215.6oC oven. Calculate the metal’s specific heat.

Using the list of specific heats, what metal is this?

Calorimetry

The enthalpy change associated with a chemical reaction or process can be determined experimentally. By measuring the heat gained or lost during a reaction at CONSTANT pressure

  • A calorimeter is a device used to measure the heat absorbed or released during a chemical or physical process

Picture:

If you leave your keys and your chemistry book sitting in the sun on a hot summer day, which one is hotter?______Why is there a difference in temperature between the two objects?

A calorimeter takes advantage of a difference in temperature between two objects due to different specific heats.

What happens in a calorimeter:

  • System loses heat to surroundings = EXO = -q
  • System absorbs heat from surroundings = ENDO = +q
  • ______= ______(ALWAYS!)

Make a chart when doing calorimetry problems:

Water / Object/Reaction
** / **
4.184

** The numbers in these two boxes are always the same, but with different signs (+/-). What heat one lost, the other gained.

EXAMPLE 1: A small pebble is heated and placed in a foam cup calorimeter containing 25.0 g of water at 25.0 C. The water reaches a maximum temperature of 26.4 C. How many joules of heat were released by the pebble?

Water / Object/Reaction
Heat (q)
Mass (m)
Specific heat (c) / 4.184
Initial Temp (Ti)
Final Temp (Tf)

Example 2: Suppose that 100.00 g of water at 22.4 °C is placed in a calorimeter. A 75.25 g sample of Al is removed from boiling water at a temperature of 99.3 °C and quickly placed in a calorimeter. The substances reach a final temperature of 32.9 °C . Determine the SPECIFIC HEAT of the metal.

Example 3: A lead mass is heated and placed in a foam cup calorimeter containing 40.0 g of water at 17.0C. The water reaches a temperature of 20.0 C.

How many joules of heat were released by the lead?

Calculating Heat During a Change of Phase

Heating/Cooling curve for water:

(Label solid, liquid, gas, melting/freezing, boiling/condensing, heat of fusion, heat of vaporization, boiling point, melting point) see p. 523 for help

As a substance heats up (or cools down) each phase & phase change needs to be calculated separately.

  • ______is the amount of heat that must be added to a stated mass of a substance to raise its temperature by 1°C, with ______.
  • Specific heat of water = 4.184 J/g°C
  • Specific heat of ice =2.09 J/g°C
  • Specific heat of steam = 2.03 J/g°C
  • CHANGING STATE requires “heat of vaporization” and “heat of fusion”
  • Heat of vaporization = amount of heat that must be added to 1 g of a liquid at its boiling point to convert it to vapor with NO temp. change

Heat of vaporization =

  • Heat of fusion = amount of heat needed to melt 1 g of a solid at its melting point

Heat of fusion =

Use your heating/cooling curve to help with calculations:

Example 1: How much heat is released by 250.0 g of water as it cools from 85.0 °C to 40.0 °C?

Example 2: How much heat energy is required to bring 135.5 g of water at 55.0 °C to its boiling point and then vaporize it?

Example 3: How much heat energy is required to convert 15.0 g of ice at -12.5 °C to steam at 123.0 °C?