Enthalpy and Hess Law

Enthalpy

The measurement of energy in a thermodynamic system. The quantity of ____ equals to the total content of heat of a system, equivalent to the system's internal energy plus the product of volume and pressure.

• H = E + PV

Hess Law

States that the enthalpy change for a chemical reaction is independent of the route taken.

• This means that the enthalpy change for the overall process will be identical regardless of how many steps are taken.

Calculating Enthalpy for an Overall Process

Step 1.) Start with the target reaction:

• Look at the overall reaction you are trying to achieve. Compare it to the given reactions and identify:

  • Which substances appear in the reactants or products in the target reaction.

  • The stoichiometric coefficients (amounts of each substance) in the target reaction.

Step 2.) Determine the position of substances:

• If a substance is a reactant in the target reaction but appears as a product in a given reaction, you’ll need to flip that reaction.

  • Target: A + B → C

  • Given: C → A + B ΔH= 100kJ

  • Flip to match target: A + B → C ΔH= - 100kJ

• If the amount of a substance in a given reaction doesn’t match the target, multiply the reaction (and its ΔH) by the correct factor.

  • Target: 2A + 2B → 2C

  • Given: A + B → C

  • Multiply given by 2

Step 3.) Adjust to cancel out unwanted substances

• If there are intermediate species in the given reactions that are not in the target reaction, you need to manipulate the reactions so these intermediates cancel out.

  • Target: A + D → C

  • Given: A → B & B + D → C

  • Add them directly to cancel B

    • [A → B] + [B + D → C] = A + D → C

Standard Enthalpy of Formation

• Formation of 1 mole of a compound from elements in standard states.

• Use known Δ𝐻𝑓∘ values from a table or data source

  • Write the balanced equation for the formation reaction. Ensure that all reactants are elements in their most stable states.

• Example: Formation of water

  • H2 (g) + ½ O2 (g) → H2O (l)

    • Δ𝐻𝑓∘ = -285.8 kJ/mol

Drawing Enthalpy Diagrams

• Exothermic: Products are at a lower energy level than reactants (ΔH<0).

  • Downward arrow for exothermic reactions. (Reactant → Product)

• Endothermic: Products are at a higher energy level than reactants (ΔH>0)

  • Upward arrow for endothermic reactions (Reactant → Product)

Exothermic Enthalpy

• ΔH is negative (heat released).

• Temperature increase in surroundings.

• If heat is a product (right side) ΔH < 0

Endothermic Enthalpy

• ΔH is positive (heat absorbed).

• Observe a temperature decrease in surroundings.

• If heat is a reactant (left side) ΔH > 0

Enthalpy Change of an Aqueous Reaction

• Use the heat equation: q = mCΔT

• Calculate ΔH:

  • q / moles

Using Hess’s Law to Calculate Enthalpy

Step 1.) Write the target reaction.

Step 2.) Manipulate the given reactions:

• Flip equations as needed and reverse ΔH.

• Scale equations to match target coefficients

Step 3.) Combine the given reactions: Ensure unwanted species cancel out

Step 4.) Add the adjusted ΔH values to find the overall ΔH

Calculating Enthalpy of Reaction from Standard Enthalpy of Formation

Step 1.) Write the balanced chemical equation for the reaction.

Step 2.) Use the formula for ΔH∘rxn

• ΔH∘rxn = ∑ΔH∘f (products) - ∑ΔH∘f (reactants)

Step 3.) Identify Δ𝐻𝑓∘ values from a data table

Step 4.) Multiply each Δ𝐻𝑓∘ by its coefficient in the balanced equation

Units - kJ/mol