Energetics | Chemistrycafe

5.0 Energetics

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Candidates should be able to:

(a) explain that most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds; the reaction can be exothermic (∆H negative) or endothermic (∆H positive)

(b) construct and interpret an energy profile diagram, in terms of the enthalpy change of the reaction and of the activation energy (see also Section 8)

(c) explain and use the terms: (i) enthalpy change of reaction and standard conditions, with particular reference to: formation; combustion; hydration; solution; neutralisation; atomisation (ii) bond energy (∆H positive, i.e. bond breaking) (see also Section 2) (iii) lattice energy (∆H negative, i.e. gaseous ions to solid lattice)

(d) calculate enthalpy changes from appropriate experimental results, including the use of the relationship: heat change = mc∆T

(e) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy

(f) apply Hess’ Law to construct simple energy cycles, e.g. Born-Haber cycle, and carry out calculations involving such cycles and relevant energy terms (including ionisation energy and electron affinity), with particular reference to:

(i) determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion (ii) the formation of a simple ionic solid and of its aqueous solution (iii) average bond energies

(g) explain and use the term entropy (h) discuss the effects on the entropy of a chemical system by the following: (i) change in temperature (ii) change in phase (iii) change in the number of particles (especially for gaseous systems) (iv) mixing of particles [quantitative treatment is not required] (i) predict whether the entropy change for a given process or reaction is positive or negative

(j) state and use the equation involving standard Gibbs free energy change of reaction, ∆G⦵ : ∆G⦵ = ∆H⦵ – T∆S⦵ [the calculation of standard entropy change, ∆S⦵ , for a reaction using standard entropies, S⦵ , is not required] (k) state whether a reaction or process will be spontaneous by using the sign of ∆G⦵ (l) understand the limitations in the use of ∆G⦵ to predict the spontaneity of a reaction (m) predict the effect of temperature change on the spontaneity of a reaction, given standard enthalpy and entropy changes

5.1 Enthalpy Change of Reaction, ΔH

Defined as the energy produced or absorbed in a chemical reaction between molar masses of reagents as indicated by a balanced stoichiometric equation.
Measured experimentally by dividing net energy change over moles of limiting reagent.
ΔH = -Q/no. of moles of limiting reagent

5.2 Standard Enthalpies

Standard Enthalpy of Formation ΔHfo

Enthalpy change when one mole of a compound is formed from its constituent elements under standard conditions of 25°C and 1 bar.
A negative value means products are more energetically stable. (ΔHf<0)
A positive value means products are more energetically unstable. (ΔHf>0)
Zero value is given to enthalpy change of elements. (ΔHf=0)

Standard Enthalpy of Combustion ΔHco

Enthalpy change when one mole of a compound is completely burnt in the presence of excess oxygen from its constituent elements under standard conditions of 25°C and 1 bar.
Value is always negative as burning fuels release energy.
Value of some substances may equal to standard enthalpy of formation of the product of its combustion.(Carbon)
Value helps us compare the usefulness of fuels.

Standard Enthalpy of Hydration ΔHhydo

Enthalpy change when one mole of gaseous ions are dissolved in large amounts of water under standard conditions of 25°C and 1 bar.
Value is always negative as bond forming(ion-dipole) releases energy.
Value depends on the charge density of ions. (Higher charge and low ionic radii will give the large magnitude of hydration energy value)

Standard Enthalpy of Solution ΔHsolno

Enthalpy change when one mole of a substance is completely dissolved in an adequately large amount of solvent such that further addition produces no heat change under standard conditions of 25°C and 1 bar.
ΔHsolno = -L.E. + ΔHhydo
If the value is negative, it means the energy evolved from hydration energy change is enough to break the strong ionic bonds and thus most likely soluble.
If the value is positive, it means the energy evolved from hydration energy change is not enough to break the strong ionic bonds and thus most likely insoluble.

Standard Enthalpy of Neutralisation ΔHno

Enthalpy change when one-mole water is formed from an acid-base reaction carried out in an aqueous solution under standard conditions of 25°C and 1 bar.
Value is always negative.
Weak acid-strong base reactions show a smaller magnitude as some of the energy evolved is used to ionise the weak-acid molecule to release the H+ cation.

Standard Enthalpy of Atomisation ΔHat o

Enthalpy change when one mole of free gaseous atoms is formed from an element or a compound under standard conditions of 25°C and 1 bar.
Value for liquids include enthalpy change of vaporisation and the bonds that are to be broken.
Value for solids includes enthalpy change of fusion and vaporisation.
Value most likely positive as it involves bond breaking.

Bond Energy

Enthalpy change to break one mole of a covalent bond between two atoms in its gaseous state.
Value is always endothermic as it is a bond breaking process

Lattice Energy

Enthalpy change when one mole of an ionic compound is formed from its constituent gaseous ions under standard conditions of 25°C and 1 bar.
Value is always negative as it is a bond forming process.
L.E is directly proportional to |q+ * q-|/(r+ + r-)

Electron Affinity

First electron affinity is the energy change when one mole of atoms gains one mole of electrons to for one mole of singly-charged anions.
First EA is mostly negative as the electron that is gained forms an electrostatic bond that is exothermic in nature.
The more electronegative the element, the higher the EA.
First electron affinity is the energy change when one mole of singly charged anions gains one mole of electrons to form one mole of doubly-charged anions.
Second EA is mostly endothermic as energy is required to overcome the electrostatic repulsion when the second electron is being added.
Successive EAs are increasingly endothermic due to increasing repulsion.

5.3 Hess' Law

Hess's Law states that the enthalpy change of the reaction is always constant and independent of the reaction pathway.

Involving ΔHfo

ΔHf = ∑ΔHf (Products) - ∑ΔHf (Reactant)

Involving ΔHco

ΔHc = ∑ΔHc (Reactants) - ∑ΔHc (Products)

5.4 Born-Haber Cycle

It is an energy cycle which shows all intermediate stages and includes all enthalpy changes in the reactions of ionic compounds.

Atomisation
Ionisation Energy
Electron Affinity
Lattice Energy
Hydration
Solution

5.5 Entropy and Entropy Change

Entropy is the degree of disorder or randomness in a system. Symbol: S Unit: J mol^-1 K^-1
Entropy change is the change in entropy (Final -Initial) and hence a measure of change in disorder.
ΔS>0 when the system becomes less orderly
ΔS<0 when the system becomes more orderly

5.6 Gibbs Free Energy

Every chemical reaction involves enthalpy change and entropy change and the combined effects of these is represented by the Gibbs Free Energy, G.
A measure of spontaneity and feasibility of reaction.
Most spontaneous reaction is exothermic forming more energetically stable products accompanied by the release of energy to surroundings.

Standard Gibbs Free Energy ΔGo = ΔHo - TΔSo

ΔGo < 0, the reaction is feasible and spontaneous.
ΔGo > 0, the reaction is not feasible and not spontaneous.
ΔGo = 0, the reaction is at equilibrium.
ΔGo is temperature is dependent.

Limitations of Standard Gibbs Free Energy ΔGo

ΔG < 0 does not indicate that a reaction will proceed. It only indicates the feasibility of the reaction.
ΔG does not indicate the speed of the reaction.
ΔG < 0 may not show any changes due to the rate of the reaction being slow.

CHEMISTRYCAFE

Build 18 December 2018

5.0 Energetics

NOTE: It may appear that there are gaps on your phone, ensure you close each table section for a better scrolling experience!

Candidates should be able to:

(a) explain that most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds; the reaction can be exothermic (∆H negative) or endothermic (∆H positive)

(b) construct and interpret an energy profile diagram, in terms of the enthalpy change of the reaction and of the activation energy (see also Section 8)

(d) calculate enthalpy changes from appropriate experimental results, including the use of the relationship: heat change = mc∆T

(e) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy

(f) apply Hess’ Law to construct simple energy cycles, e.g. Born-Haber cycle, and carry out calculations involving such cycles and relevant energy terms (including ionisation energy and electron affinity), with particular reference to:

(i) determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion (ii) the formation of a simple ionic solid and of its aqueous solution (iii) average bond energies

5.1 Enthalpy Change of Reaction, ΔH

5.2 Standard Enthalpies

Standard Enthalpy of Formation ΔHfo

Standard Enthalpy of Combustion ΔHco

Standard Enthalpy of Hydration ΔHhydo

Standard Enthalpy of Solution ΔHsolno

Standard Enthalpy of Neutralisation ΔHno

Standard Enthalpy of Atomisation ΔHat o

Bond Energy

Lattice Energy

Electron Affinity

5.3 Hess' Law

Involving ΔHfo

Involving ΔHco

5.4 Born-Haber Cycle

5.5 Entropy and Entropy Change

5.6 Gibbs Free Energy

Standard Gibbs Free Energy ΔGo = ΔHo - TΔSo

Limitations of Standard Gibbs Free Energy ΔGo

5.0 Energetics

NOTE: It may appear that there are gaps on your phone, ensure you close each table section for a better scrolling experience!

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Candidates should be able to:

(a) explain that most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds; the reaction can be exothermic (∆H negative) or endothermic (∆H positive)

(b) construct and interpret an energy profile diagram, in terms of the enthalpy change of the reaction and of the activation energy (see also Section 8)

(d) calculate enthalpy changes from appropriate experimental results, including the use of the relationship: heat change = mc∆T

(e) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy

(f) apply Hess’ Law to construct simple energy cycles, e.g. Born-Haber cycle, and carry out calculations involving such cycles and relevant energy terms (including ionisation energy and electron affinity), with particular reference to:

(i) determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion (ii) the formation of a simple ionic solid and of its aqueous solution (iii) average bond energies

5.1 Enthalpy Change of Reaction, ΔH

5.2 Standard Enthalpies

Standard Enthalpy of Formation ΔHfo

Standard Enthalpy of Combustion ΔHco

Standard Enthalpy of Hydration ΔHhydo

Standard Enthalpy of Solution ΔHsolno

Standard Enthalpy of Neutralisation ΔHno

Standard Enthalpy of Atomisation ΔHat o

Bond Energy

Lattice Energy

Electron Affinity

5.3 Hess' Law

Involving ΔHfo

Involving ΔHco

5.4 Born-Haber Cycle

5.5 Entropy and Entropy Change

5.6 Gibbs Free Energy

Standard Gibbs Free Energy ΔGo = ΔHo - TΔSo

Limitations of Standard Gibbs Free Energy ΔGo