IB1 Chemistry HL Energetics Why do chemical reactions happen?

Topic 15: Energetics (8 hours) 15.1 Standard enthalpy changes of reaction Define and apply the terms standard state, standard enthalpy change of formation (¬H ) f Ö and standard enthalpy change of combustion (¬H ) c Ö Determine the enthalpy change of a reaction using standard enthalpy changes of formation and combustion Born–Haber cycle Define and apply the terms lattice enthalpy and electron affinity Explain how the relative sizes and the charges of ions affect the lattice enthalpies of different ionic compounds Construct a Born–Haber cycle for group 1 and 2 oxides and chlorides, and use it to calculate an enthalpy change Discuss the difference between theoretical and experimental lattice enthalpy values of ionic compounds in terms of their covalent character Entropy State and explain the factors that increase the entropy in a system Predict whether the entropy change (ΔS) for a given reaction or process is positive or negative Calculate the standard entropy change for a reaction (¬S Ö ) using standard entropy values (S Ö ) Spontaneity 2.5 hours Predict whether a reaction or process will be spontaneous by using the sign of ¬GÖ Calculate ¬GÖ for a reaction using the equation ¬GÖ = ¬H Ö− T¬S Ö and by using values of the standard free energy change of formation, ΔGf Ö Predict the effect of a change in temperature on the spontaneity of a reaction using standard entropy and enthalpy changes and the equation ¬GÖ= ¬H Ö− T¬S Ö.

Standard enthalpy change of reaction in standard state at 101kPa, 298K (PV = nRT)

Standard enthalpy of combustion:  H c  When a substance is fully burned in oxygen for example CH 4 + 2O 2  CO 2 + 2H 2 O

Standard enthalpy of formation:  H f   H f  : The energy absorbed or evolved when 1 mol of the substance is formed from its elements in their standard states. The enthalpy of formation of any element is zero.  H 2(g) + ½O 2(g)  H 2 O (l)  H f  = -285 kJ/mol   H =  H f(products) -  H f(reactants)

Ionic compound consist of ions arranged in a lattice

Two or more electrons can be transferred  Different sized atoms give different mineral structures as they pack in a different way Hexagonal Beryl crystal; Image Wikipedia

Lattice enthalpy,  H lattice  Relates to the endothermic process MX(s)  M + (g) + X - (g) in which the gaseous ions of a crystal are separated to an infinite distance from each other.  NaCl (s)  Na + (g) + Cl - (g)  H lattice = 771kJ/mol Endothermic reactions.

Lattice enthalpy depends on  charge on an ion  size of an ion  packing arrangement MgO -3791kJ/mol NaCl -790kJ/mol

Standard enthalpy of atomization:  H at   H at  : The energy required to atomize one mole of an element.  Na (s)  Na (g) + e -  H at  = +108 kJ/mol (physics: related to latent heat of vaporization)

Electron affinity The enthalpy change per mole when an atom gains one electron in the gaseous phase Cl (g) + e - (g)  Cl - (g)  H ea = -351 kJ/mol.  Electron affinity can be both exothermic and endothermic depending on element.

How does the electron affinity change across a period?

Born-Haber cycle

Born-Haber cycles used to calculate lattice enthalpies  Can be used to find out if a bond is more or less ionic

Born Haber cycle for sodium chloride  Enthalpy of formation of NaCl= -411kJ/mol  Enthalpy of atomisation of Na = +103 kJ/mol  Enthalpy of atomisation of Cl = +121 kJ/mol (½ energy of Cl-Cl bond )  Electron affinity of Cl= -364 kJ/mol  Ionisation energy of Na= kJ/mol

Lattice enthalpy for sodium chloride Energies of atomisation + Electron affinity + Ionisation energy = = Enthalpy of formation + Lattice enthalpy Lattice energy = 771 kJ/mol

How ionic is the ionic lattice? Chemistry Data Booklet gives lattice enthalpies as:  Experimental values (obtained from Born-Haber cycle)  Theoretical values (calculated using electrostatic calculations) Greater difference between theoretical and experimental values  more covalent character of the bond.

Decomposition of ammonium nitrate NH 4 NO 3(s)  N 2 O (g) + 2 H 2 O (l)  NH 4 NO 3(s)  H f  = -366 kJ/mol  S  = 151 J/K*mol  N 2 O (g)  H f q = +82 kJ/mol  S  = 220 J/K*mol  H 2 O (l)  H f q = -285 kJ/mol  S   = 70 J/K*mol  H = [  H f (N2O(g)) +  H f (H2O(l)) ] – [  H f(NH4NO3(g) ] = =[82 + 2(-285)] - [-366] = -122 kJ/mol

Entropy: disorderdisorder

Which is more likely if the particles are in constant random motion?

Entropy Entropy, S = DisorderUnit: JK -1 mol -1   S = change in disorder   S = S p - S r  Absolute value of S can be measured

Entropy and changes of stae SolidLiquidGas H2OIceWaterSteam JK -1 mol Increasing entropy 

Will a reaction happen? Spontaneity

All reactions involve changes in H and S  S is probably positive if moles of gas increase and moles of solid or liquid decrease. NH 4 Cl (s)  NH 3(g) + HCl (g)  S = + 285JK -1 mol -1 Pb 2+ (aq) + 2 I -  PbI 2(s)  S = - 70 JK -1 mol -1

Spontaneity of a reaction Nature likes low internal energy (  H to decrease) and high disorder (  S to increase)  A reaction will occur if the final state is more probable than the initial state.  Decrease in  H  Increase in  S

Gibbs free energy,  G  =  H  - T  S  Temperature dependent Spontaneous:  G negative (  G  < 0) Equilibrium:  G  = 0 Not spontaneous:  G  positive(  G  > 0)

HH SS GG Spontaneity Negative (Exothermic) Positive (More random)  G < 0 Always negative Always spontaneous Positive (Endothermic) Negative (More order)  G > 0 Always positive Never spontaneous Negative (Exothermic) Negative (More order) Depends on T Spontaneous at low Temp Positive (Endothermic) Positive (More random) Depends on TSpontaneous at high Temp

Activation energy is also important  Just the fact that a reaction is spontaneous doesn’t mean that it will occur at once.  It also depends on activation energy. And we will deal with that later on in topic 7.

Links  Ionic bonding ng/ ng/  Covalent bonding nd/ nd/