Thermochemistry Part 1

Thermochemistry Overview

• Definition of Energy:

  • Ability to do work or produce heat; everything possesses it.

  • Law of Conservation of Energy: Energy cannot be created or destroyed; it can only be transformed from one form to another.

Types of Energy

Potential Energy

• Energy due to position or composition.

Kinetic Energy

• Energy due to the motion of an object.

• Depends on:

  • Mass (m)

  • Velocity

Heat vs. Temperature

• Temperature: Reflects random motion of particles in a substance.

• Heat: Energy transfer between objects due to a temperature difference.

Chemical Energy Transfer

• The universe during a reaction can be divided into:

  • System: The part we focus on, including reactants and products.

  • Surroundings: Everything else, often the reaction container and ambient conditions; usually water

• energy gained by surroundings = energy lost by the system

Heat Transfer in Reactions

• Exothermic: Heat flows out of the system (surroundings gain energy); results in a rise in temperature.

• Endothermic: Heat flows into the system (surroundings lose energy); results in a decrease in temperature.

Laws of Thermodynamics

Zeroth Law

• States that if A=B and B=C, then A=C (transitive property of equality).

First Law

• Restates the conservation principle in the context of energy: energy cannot be created nor destroyed, only transformed.

Things to know:

• (aq) = inside water, water temp is tracked (ΔT)

• we cannot monitor the temperature of a chemical reaction directly, so we monitor temperature of the surroundings

Calorimetry

• Calorimeter: Device used to determine heat change during a chemical reaction by observing temperature changes.

• Heat Capacity:

  • Defined as C = heat absorbed / temperature increase.

  • Depends on the quantity of the substance.

• Specific Heat Capacity:

  • Energy required to raise the temperature of 1 g of a substance by 1°C (J/K•g).

  • formula = q = mcΔT where q is the amount of heay (in joules), m is the mass of the substance (in grams), c is the specific heat capacity (in J/K•g), and ΔT is the change in temperature (in °C).

Heat Transfer Calculations

Molar Heat Calculations

molar heat capacity: energy required to raise the temperature of 1 gram of a substance by 1°C (J/K•mole)

• Use:

  • q (heat) = n × ΔH (change in enthalpy) or other forms based on phase changes or reactions.

  • n = moles

  • q = kJ

  • ΔH = kJ/mole

• Types of Heat:

  • Heat of Fusion (melting) solid to liquid:

    • q = n × ΔH_fus (energy needed to melt).

    • q = mass/molar mass

    • (+) heat needed to melt ice into liquid water.

    • (-) heat released when freezing.

  • Heat of Vaporization (boiling) liquid to gas:

    • q=n × ΔH_vap (energy needed to vaporize).

    • (+) heat needed to convert liquid water to steam.

    • (-) heat released when condensing steam.

  • Heat of Dissolution solid to aqueous solution:

    • q = n × ΔH_auss (heat associated with dissolving solids into aqueous solutions).

    • (+) denotes heat absorbed; (-) denotes heat released.

  • Chemical Reactions ΔH_rxn

    • heat of reaction

    • q = n ΔH_rxn

    • balanced equation coefficients will impact ΔH_rxn values

Hess's Law

• Establishes that the change in enthalpy for a reaction is the same regardless of whether the reaction occurs in one step or multiple steps.

• Characteristics of ΔH in a reaction:

  1. If the reaction is reversed, the sign of ΔH also reverses.

  2. The magnitude of ΔH is proportional to the quantities of reactants and products involved in the reaction.

• + ΔH = reaction needs energy = endothermic

• - ΔH = reaction releases energy = exothermic

• H2O gas and H2O liquid are DIFFERENT, so they must be treated as separate species.

Enthalpy Changes

• Enthalpy (ΔH): Measured as the heat of a reaction; unless one species in the reaction is being analyzed can be calculated as:

  • ΔH = H_products - H_reactants

• Units: kJ/mole preferred in calculations.

• to obtain proper ΔH units of kJ/mole, we need to know with respect to which species

• Standard States: Reference states (g, l, s) used for comparability; conditions include pure substances at 1 atm pressure and 25°C.

• standard enthalpy of formation ( ΔH_f°): the change in enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states

• the enthalpy change for a given reaction is calculated by subtracting the enthalpies of formation of the reactants from the enthalpies of formation of the products

  • formula: ΔH° = Sum of VΔHf°(product) - Sum of VΔHf°(reactants).

• elements require no change to form so ΔH = 0

Bond Energy Calculations

• chemical bonds: forces that cause a group of atoms to behave as a unit; they occur when atoms are more stable (lower in energy)

• bonds are helpful with dealing with the energies of chemical reactions

• as the number of shared electrons increase, the bond length shortens

• Bond Enthalpy: Used to estimate energy changes during reactions.

• Energy must be added to break bonds (endothermic) and released during bond formation (exothermic).

• Approximates energy involved in breaking/forming reaction bonds in terms of reactant and product bonds.

• formula Δ H = Σ Δ H ( reactant bond energy values ) − Σ Δ H ( product bond energy values)