Unit 6:Thermochemistry.

6.1 Endothermic and Exothermic Processes.

Determine if a process is endothermic or exothermic and assign the proper sign to ∆H.

  1. Is melting ice endothermic or exothermic? Explain your reasoning.
  2. A + B → AB ΔH = −50 kJ·molrxn-1 Is the reaction endothermic or exothermic. Justify how you know.

6.2 Energy Diagrams.

Identify reactants and products as being higher in energy on an energy diagram.


    energy diagram with 3 hills

    Image credit: libretexts

  1. Does the energy diagram above show a endothermic or exothermic reaction. Justify your answer.
  2. Label the reactants, intermediates, and products on the energy diagram.

Determine if a process is endothermic or exothermic from an energy diagram.

  1. Given an energy diagram, identify if the reaction absorbs or releases energy.
  2. Explain how the difference in energy between reactants and products indicates endothermicity or exothermicity.

Calculate ∆H from values on an energy diagram.

  1. Using a diagram with energy levels of reactants and products, calculate ∆H.
  2. If reactants are at 150 kJ/mol and products at 80 kJ/mol, what is ∆H?

6.3 Heat Transfer and Thermal Equilibrium.

Interpret a particle diagram using arrows to show heat transfer.

  1. Draw arrows to show heat flow when a hot metal rod is placed in cooler water.
  2. Explain the direction of heat transfer between two particle systems at different temperatures.

Justify energy transfer as resulting from collisions between warmer and cooler particles.

  1. Why does heat flow from a hot object to a cold object at the particle level?
  2. Describe how particle collisions increase the energy of cooler particles.

Assign the direction of heat flow based on temperature data for two system components.

  1. If a 50°C metal is placed in 20°C water, which way does heat flow?
  2. Explain how temperature differences determine the direction of energy transfer

6.4 Heat Capacity and Calorimetry.

Calculate any variable in q = mc∆T for heating, cooling, and phase changes.

  1. Calculate the heat required to raise the temperature of 100 g of water by 25°C (c = 4.18 J/g°C).
  2. Determine the final temperature if 500 J is added to 50 g of water starting at 20°C.

Use conservation of energy to calculate temperature changes based on heat lost/gained by another substance.

  1. A 100 g metal at 80°C is placed in 200 g of water at 20°C. What is the final temperature?
  2. Explain how energy lost by one object is equal to energy gained by another.

6.5 Energy of Phase Changes.

Differentiate between energy changes due to heating/cooling and phase transitions in a pure substance.

  1. How does heating water differ from melting ice in terms of energy changes?
  2. Explain why temperature remains constant during a phase change even when heat is added.

Calculate the heat of a phase change given temperature, mass, and enthalpy of phase change data.

  1. Calculate the energy required to melt 50 g of ice (∆Hfus = 334 J/g).
  2. If 2000 J is absorbed, how much ice will melt at 0°C?

6.6 Introduction to Enthalpy of Reaction.

Describe energy exchange between reaction systems and surroundings.

  1. Explain why an exothermic reaction warms its surroundings.
  2. Describe energy flow in an endothermic reaction at the molecular level.

Calculate the amount of heat absorbed/released in a reaction using ∆H in a stoichiometric calculation.

  1. How much energy is released when 2 moles of H2 react with O2 to form water (∆H = -286 kJ/mol)?
  2. Given ∆H, calculate heat change for a reaction involving 3 moles of reactant.

Compare and contrast the chemical potential energy of reactants and products in a system.

  1. Which has higher potential energy in an exothermic reaction, reactants or products?
  2. Explain how bond energies affect the chemical potential energy difference.

6.7 Bond Enthalpies.

Justify the sign of ∆H by comparing bond strengths of reactants and products.

  1. How does stronger product bonding than reactants affect ∆H?
  2. Explain why breaking bonds requires energy and forming bonds releases energy.

Calculate ∆H using bond energies and the equation ∆H = ∑[bonds broken] - ∑[bonds formed].

  1. Calculate ∆H for H2 + Cl2 → 2 HCl using bond energies.
  2. If bonds broken total 500 kJ and bonds formed total 600 kJ, what is ∆H?

6.8 Enthalpy of Formation.

Calculate ∆H using enthalpies of formation and the equation ∆H°rxn = ∑∆Hf°prod - ∑∆Hf°react.

  1. Given ∆Hf° of reactants and products, calculate ∆H°rxn.
  2. Explain why the sum of formation enthalpies can determine reaction enthalpy.

6.9 Hess’s Law.

Identify a valid series of steps for an overall process, considering stoichiometry and energy change.

  1. Determine which sequence of reactions correctly represents an overall reaction.
  2. Explain how ∆H of steps combines to give overall ∆H.

Manipulate reaction steps to match an overall process and calculate ∆H.

  1. Reverse or multiply reactions to match a target equation and calculate ∆H.
  2. Show how multiplying a reaction by 2 affects its enthalpy change.

Use Hess’s law results to justify the overall sign of a process, even if component steps have different signs.

  1. Given reactions with mixed ∆H signs, determine overall ∆H using Hess’s law.
  2. Explain why the overall process can be exothermic even if some steps are endothermic.
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