Topic 5/15 Energetics and Thermochemistry

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enthalpy

represented by H

the total energy of a system, some of which is stored as potential energy in chemical bonds; aka “heat content”

standard enthalpy change of reaction

represented by ∆H°

heat transferred in a reaction under **standard conditions** (101.3 kPa, 298 K)

exothermic reaction

a chemical reaction in which heat energy is **released** to the surroundings

endothermic reaction

a chemical reaction in which energy is **absorbed** from the surroundings

calorimetry

a technique used to measure the enthalpy associated with a particular change

heat capacity

The amount of energy that a given mass of substance (either 1 g or 1 kg) can absorb that produces a 1ºC increase in temperature

Hess’s Law

states that the enthalpy change for a reaction is the same, regardless of the reaction route. This means that the enthalpy changes of multiple reaction steps can be added in order to determine the enthalpy change for the overall reaction

average bond of enthalpy

represented by D

the amount of energy that must be supplied to break **1 mole** of chemical bonds in an **isolated molecule** in the **gaseous** state

* It’s also the amount of energy released when this bond forms
* They are a measure of the strength of the covalent bond
* They are also called bond dissociation energies

standard state

the standard thermodynamic conditions chosen for substances when listing or comparing thermodynamic data

* 1 atm pressure (760.0 mmHg or 101.3 kPa) and 25ºC (298 K)

standard enthalpy change of formation

represented by ΔH_f°

the enthalpy change for the formation of one mole of a substance in its standard state from its elements in their reference form and in their standard states

standard enthalpy of combustion

represented by ∆H_cΘ

the enthalpy change when 1 mole of the compound undergoes complete combustion in excess oxygen under standard conditions

* It’s always exothermic
* Enthalpies of formation of many oxides are equivalent to the enthalpies of combustion of the element

lattice enthalpy

the energy required to convert one mole of a substance from the solid state as a compound to the gaseous state as ions

enthalpy change of atomization

the enthalpy change to convert an element in its standard state to one mole of gaseous atoms

electron affinity

the enthalpy change for one mole of gaseous atoms to gain electrons and form a mole of anions

* For most atoms this is exothermic
* Gaining a 2nd electron is endothermic because of the repulsion between the electron and the negative ion

ionization energy

the enthalpy change for one mole of gaseous atoms to lose electrons and form a mole of cations

enthalpy of formation

the enthalpy change for the formation of one mole of a substance in its standard state from its elements in their reference form and in their standard states

Construct a Born-Haber cycle and use it to calculate an enthalpy change.

enthalpy of formation = enthalpy of atomization(s) + ionization energies + electron affinities + lattice enthalpy

solvation (dissolution)

any stabilizing interaction of a solute and the solvent or a similar interaction of solvent with groups of insoluble material

entropy (S)

the amount of randomness or disorder of a system

* The natural tendency of the universe is towards greater disorder
* A positive entropy change indicates a system is more disordered
* A negative entropy change indicates a system is less disordered (more ordered)

standard molar entropy ( S°) a.k.a. absolute entropies

the entropy change of one mole of the pure substance at 1 atm pressure being heated from 0 K to 298 K (in J/K mol)

∆S° = ΣS° (products) - ΣS° (reactants)

First Law of Thermodynamics

in any process, spontaneous or non-spontaneous, the total energy content of a system and its surroundings is constant

Second Law of Thermodynamics

in any spontaneous process, the total entropy of a system and its surroundings always increases