Energy Diagrams

Law of Conservation of Energy

• The Law of Conservation of Energy tells us that energy is neither created nor destroyed, only converted into other forms.

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• This means that in a isolated system energy is converted back and forth between kinetic and potential, but the total energy of the system doesn’t change.

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ΔE_universe = ΔE_system + ΔE_surroundings

System vs. Surroundings

• Virtually every chemical reaction and change of state either releases or absorbs heat. This is because in the real world, reactions are never isolated.
• We will thus look at chemical reactions in terms of system vs. surroundings.
• The system is the specific part (chemical reaction) that you are interested in.
• The surroundings are everything else.

Endothermic vs. Exothermic

• An exothermic reaction occurs when more energy is released in forming new bonds than is required to break bonds in the initial reactants.
• A good example of an exothermic reaction is a hot pack. Heat produced by the reaction in the heat pack (the system) moves to your hands (surroundings).
• Other Examples: Freezing, condensation, rusting iron, candle burning

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Endothermic vs. Exothermic

• An endothermic reaction occurs when a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the products molecules

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• Examples: Melting, evaporation, photosynthesis

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Enthalpy

• Enthalpy is the heat content of a system at a constant pressure (usually atmospheric pressure). The symbol for enthalpy is H.

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• The enthalpy (heat) of a reaction is written as ΔH_rxn

ΔH_rxn = ΔH˚_f (products) - ΔH˚_f (reactants)

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• The enthalpy (heat) of a reaction measures the change in heat (in other words energy).
• Heat of formation (ΔH˚_f) is the change in enthalpy that accompanies the formation of one mole of the compound in its standard state from its constituent elements in their standard states.

Examples

• 4Fe + 3O₂ 🡪 2Fe₂O₃ + 1625 kJ or

4Fe + 3O₂ 🡪 2Fe₂O₃ ΔH_rxn = -1625 kJ

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• 27 kJ + NH₄NO_{3 (s)} 🡪 NH₄⁺ + NO₃^- or

NH₄NO_{3 (s)} 🡪 NH₄⁺ + NO₃^- ΔH_rxn = 27 kJ

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Exothermic Reactions

In an exothermic reaction:

• The energy of the products is less than the energy of the reactants.
• Heat is released.
• Heat is a product.
• The surroundings get warmer.
• ΔH = negative

Endothermic Reaction

In an endothermic reaction:

• The energy of the products is greater than the energy of the reactants.
• Heat is absorbed.
• Heat is a reactant.
• The surroundings get colder.
• ΔH = positive

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Collision Theory

• For a reaction between two substance to take place, the molecules of those two substances must collide (hit each other).

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• If they collide with enough energy, they will react. The minimum amount of kinetic energy required to react when they collide is known as activation energy.

Energy Diagrams

• Energy diagrams show the potential energy of a reaction over the course of time.

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• The peak on the graph represents the activated complex, which is the point when the two molecules are reacting.

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• The activation energy is the difference between the energy of the activated complex and the energy of the reactants.

Catalysts

• Catalysts lower the activation energy for a given reaction. They do not change the overall enthalpy of the reaction.

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• Catalysts are not consumed by the reaction – this means that they can be used over and over, and don’t affect the products of a reaction.

Labeling an Energy Diagram

Drawing an Energy Diagram

• For the theoretical reaction A + B 🡪 C + D, the energy of the reactants is 30 kJ/mol. The energy of the products is 4 kJ/mol, while the activation energy is 12 kJ/mol.
• Draw the graph and label 1) reactants, 2) products, 3) activation energy, 4) enthalpy of the reaction, ∆H.