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 = cal = 1 Cal (nutritional)1 kJ = J1 calorie = J1 kcal = 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 kJCalculate 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 kJAlternate 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 upHf 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 –o Endothermic reactions : absorbs heat from surroundings (+).

If you touch an endothermic reaction it feels COLDo 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/gC):o CO2(g) = 0.843 J/gCo Cu(s) = 0.382 J/gCo Fe(s) = 0.446 J/gCo H2O (l) = 4.184 J/gC

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/gC)

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 50.0C?

Example 2: Iron has a specific heat of 0.446 J/gC. 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. (22.0C), 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?

CalorimetryThe 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 processPicture:

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/ReactionHeat (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 17.0C. 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 ____________________.

o Specific heat of water = 4.184 J/g°Co Specific heat of ice =2.09 J/g°Co 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?