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system = part of the universe that contains the reaction or process being studied

surroundings = rest of the universe that interacts with the system

Thermochemistry is the study of heat changes (energy!) that accompany chemical reactions and phase changes• Heat is exchanged between the system and

surroundings

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Enthalpy (H) is the heat content of a system at constant pressure. (Enthalpy is potential energy)

Change in enthalpy for a reaction is called the enthalpy (or heat) of reaction ( Hrxn).

Hrxn = Hfinal - Hinitial

Hrxn = Hproducts - Hreactants

The second equation shows you how you can think about enthalpy of a reaction as the difference in energy of the products and reactants.

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Exothermic Reaction

H<0, - HEnth

alpy

Reactants

Products

Energy is

released

Energy required to break bonds

Energy released to form bonds<

What happens to the surrounding?

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Endothermic Reaction

H>0, + H

Enth

alpy

Reactants

Products

Ener

gy is

abso

rbed

Energy required to break bonds

Energy released to form bonds>

What happens to the surrounding?

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Is this reaction exothermic or endothermic?

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Thermochemical equations: balanced chemical equation that includes the physical states of all reactants and products and the energy change, usually expressed as the change in enthalpy ( H).

Example:

CH4(g) + 2O2 (g) CO2(g) + 2H2O(g) Hrxn= -802.3kJ

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Coffee cup calorimeter

Reaction happens in water

Bomb calorimeter

-qsystem = qwater

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Calorimetry Lab Calculations

Step 1: Finding how much heat was transferred from metal to water (remember...heat flows from higher temperature to lower temperature)

Step 2: Using heat transferred, calculate specific heat of metal.

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Mass of metal = 50.5g

Initial temperature of metal = 100ºC

Final Temperature of metal = 25.5ºC

Sample Calculations

Mass of water = 75g

Initial Temperature of water = 21.0ºC

Final Temperature of water = 25.5ºC

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Hess's Law• Used when it is impractical to measure ΔH using a

calorimeter• Hess's Law states if you can add two or more

thermochemical equations to produce a final equation, the sum of the enthalpy changes of each individual reaction is the enthalpy for the final reaction

Imagine that chemical reactions happen in small steps. You add all of the steps together along with their enthalpy changes

Hoverall=)))H1)+)))H2)+)))H3).).).∆ ∆∆ ∆

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Example 17:

Calculate ΔH for the formation of sulfur trioxide

2S(s) + 3O2(g) !�2SO3(g) ΔH =?

Given:

S(s) + O2(g) ! SO2(g) ΔH = -297kJ

2SO3(g) ! 2SO2(g) + O2(g) ΔH = 198 kJ

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Hess's Law Rules1. If the reaction is reversed, the sign of ΔH is also

reversed2. If the coefficients are multiplied by a factor,

then the ΔH is multiplied by the same factor3. The state of matter (s, l, g, aq) of reactants and

products affects whether they cancel out or not.

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Example 18:

Calculate the enthalpy of reaction of

2C(s) + O2(g) + 2H2O(l) !��H2CO2(l) ΔH=?

Given:

C(s)+1/2O2(g) + H2O(l)!CO(g)+ H2O(l) ΔH=-110.5kJ

CO(g)+ H2O(l)!H2CO2(g) ΔH=33.7 kJ

H2CO2(g)! H2CO2(l) ΔH=-62.9 kJ

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Example 19:

Calculate the enthalpy of reaction of

2H2O2(l) !��H2O(l) + O2 ΔH=?

Given:

2H2(g) + O2(g)!2H2O(l) ΔH=-572 kJ

H2(g) + O2 !H2O2(l) ΔH=-188 kJ

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Not practical to calculate every ΔH to use in Hess's Law.

Use Standard Enthalpy of Formation!

Standard Enthalpy of Formation: Change in enthalpy when one mole of a compound forms from it's constituent elements in their standard states.• standard state: normal physical state at 1 atm and

25ºC• For a pure elements in its standard state ΔH = 0

C(s) + 2H2(g) CH4(g) H= -74.6 kJ/mol∆

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Using Standard Enthalpies of formation• You can use standard enthalpies of formation to

calculate enthalpy of reaction under standard conditions (ΔHºrxn) using Hess's Law

Example 20:

H2S(g) + 4F2(g) 2HF(g) + SF6(g) ΔHºrxn =?

Given:• 1/2H2(g) + 1/2F2(g) HF(g) ΔHºf=-273kJ• S(s) + 3F2(g) SF6(g) ΔHºf= -1220kJ• H2(g) + S(s) H2S(g) ΔHºf= -21kJ

ΔHºrxn= ΔHºf(product) - ΔHºf(reactants)

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ΔHºrxn= ΔHºf(product) - ΔHºf(reactants)

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Example 21: Using the chart given to you, calculate ΔHºrxn for the following reactions:

a) CH4(g) + 2O2(g) CO2(g) + 2H2O(l)

b) CaCO3 CaO(s) + CO2(g)

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Calculating ΔH using bond energy (BE) data• you can estimate ΔH for gaseous molecules• Bond formation is ____________• Breaking bonds is ____________• ΔH = BEbreaking - BEforming

Example 21: Determine ΔH for the following reaction given the bond energies

N2(g) + 3H2 2NH3(g)N=N 941 kJ/mol

H-H 436 kJ/mol

N-H 391 kJ/mol

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