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LATTICE ENTHALPY
WARNING There can be two definitions - one is the opposite of the other! Make sure you know which one is being used.
Lattice Formation Enthalpy ( ∆LEΗ ) OCR preferred AQA can use either definition
Definition The enthalpy change when ONE MOLE of an ionic crystal lattice is formed from its isolated gaseous ions.
Values • highly exothermic - strong electrostatic attraction between ions of opposite charge
- a lot of energy is released as the bond is formed
- relative values are governed by the charge density of the ions.
Example Na+(g) + C l ¯(g) -——> Na+^ C l ¯(s)
Notes • one cannot measure this value directly; it is found using a Born-Haber cycle
- the greater the charge densities of the ions, the more they attract each other and the larger the lattice enthalpy.
- the more exothermic the lattice enthalpy, the higher the melting point
ENDOTHERMIC EXOTHERMIC
Lattice Dissociation Enthalpy ( ∆LEΗ ) Option for
Definition The enthalpy change when ONE MOLE of an ionic lattice dissociates into isolated gaseous ions.
Values • highly endothermic - there is a strong electrostatic attraction between ions of opposite charge
- a lot of energy must be put in to overcome the attraction
Example Na+^ C l ¯(s) -——> Na+(g) + C l ¯(g)
Knockhardy Publishing
M X¯
(s)
(g) (^) (g)
M + X¯
LATTICE
DISSOCIATION
ENTHALPY
M X¯
(s)
(g) (^) (g)
M + X¯
LATTICE
FORMATION
ENTHALPY
Consequences
HIGHER CHARGE DENSITY IONS LOWER CHARGE DENSITY IONS GREATER ATTRACTION LESS ATTRACTION LARGER LATTICE ENTHALPY SMALLER LATTICE ENTHALPY
Thermal stability and Lattice Enthalpy
Oxides • thermal stability of Group II oxides decreases down the group
Mg2+^ O2-^ Ca2+^ O2-^ Sr2+^ O2-^ Ba2+O2- Lattice Enthalpy (kJ mol-1) -3889 -3513 -3310 - Melting Point (°C) 2853 —— decreasing values ——>
MgO • magnesium oxide is used to line furnaces - REFRACTORY LINING
- this is because of its high melting point (2853°C)
- the high melting point is a result of the large (highly exothermic) lattice enthalpy
- high lattice enthalpy due to the attraction between ions of high charge density
Carbonates • thermal stability of Group II carbonates increases down the group
- MgCO 3 decomposes much easier than BaCO 3
- BUT the lattice enthalpy of MgCO 3 is HIGHER!
MgCO 3 CaCO 3 SrCO 3 BaCO 3 Decomposes at 350°C 832°C 1340°C 1450°C
Lattice Enthalpy (kJ mol-1) -3123 ———————> -
- Mg2+^ ions are SMALLER and have a HIGHER CHARGE DENSITY
- this makes them MORE HIGHLY POLARISING
- they DISTORT THE CO 3 2-^ ion
- this WEAKENS THE ATTRACTION BETWEEN IONS
- the LATTICE IS NOT AS STRONG
Knockhardy Publishing
Mg2+^ O2-^
Na+^ Cl¯
Q.1 Which substance in the the following pairs has the larger lattice enthalpy?
a) NaCl or KCl b) NaF or NaCl
c) MgCl 2 or NaCl d) MgO or MgCl 2
First Ionisation Energy
Definition The energy required to remove one mole of electrons (to infinity) from one mole of gaseous atoms to form one mole of gaseous positive ions.
Values Always endothermic need to overcome the pull of the nucleus on the electron
Example(s) Na(g) -——> Na+(g) + e¯ and Mg(g) -——> Mg+(g) + e¯
Notes • There is an ionisation energy for each successive electron removed.
e.g. SECOND IONISATION ENERGY Mg+(g) -——> Mg2+(g) + e¯
- Look back in your notes to refresh your memory about the trends in I.E.‘s
Electron Affinity
Definition The enthalpy change when ONE MOLE of gaseous atoms acquires ONE MOLE of electrons (from infinity) to form ONE MOLE of gaseous negative ions.
Values Always exothermic - a favourable process due to the nucleus attracting the electron
Example C l (g) + e¯ -——> C l ¯(g)
Notes • Do not confuse electron affinity with electronegativity.
Knockhardy Publishing
Q.5 Write equations representing the following electron affinity (EA) changes;
1st EA of bromine
1st EA of oxygen
2nd EA of oxygen
Q.4 Write equations representing the following ionisation energy changes;
1st IE of calcium
2nd IE of calcium
1st IE of lithium
1st IE of aluminium
BORN-HABER CYCLES
Theory • involve the application of Hess’s Law
- used to outline the thermodynamic changes during the formation of ionic salts
- used to calculate Lattice Enthalpy
- Lattice Enthalpy cannot be determined directly by experiment
BORN-HABER CYCLE FOR SODIUM CHLORIDE
According to Hess’s Law, the enthalpy change is independent of the path taken. Therefore...
STEP 6 = - (STEP 5) - (STEP 4) - (STEP 3) - (STEP 2) + (STEP 1)
- (-364) - (+500) - (+121) - (+108) + (-411) = - 776 kJ mol-
Knockhardy Publishing
1
2
3
6
4
5
Na(s) + ½C l 2 (g)
NaC l (s)
Na(g)+ ½C l 2 (g)
Na(g) + C l (g)
Na+(g) + C l (g)
Na+(g) + C l ¯(g)
STEPS (values are in kJ mol-1) ¬ Enthalpy change of formation of NaC l Na(s) + ½Cl 2 (g) ——> NaC l (s) – 411 Enthalpy change of sublimation of sodium Na(s) ——> Na(g) + 108 ® Enthalpy change of atomisation of chlorine ½C l 2 (g) ——> C l (g) + 121 ¯ Ist Ionisation Energy of sodium Na(g) ——> Na+(g) + e¯ + 500 ° Electron Affinity of chlorine C l (g) + e¯ ——> C l ¯(g) – 364 ± Lattice Enthalpy of NaC l Na+(g) + C l ¯(g ) ——> NaC l (s)
Q.6 Construct a similar Born-Haber cycle for NaCl 2.
If the Lattice Enthalpy of NaCl 2 is -3360 kJ mol -1, what is its enthalpy of formation? What does this tell you about the stability of NaCl 2?
Will an ionic salt dissolve in water?
Introduction If a pair of oppositely charged gaseous ions are placed together, they will attract each other. The energy change (LATTICE ENTHALPY) is highly exothermic.
If the ions were put in water, they would be attracted to polar water molecules. the resulting energy change (HYDRATION ENTHALPY) is highly exothermic.
In both; the greater charge density of the ions = a more exothermic reaction
The missing stage of the cycle is known as the ENTHALPY OF SOLUTION.
The size and value of the enthalpy of solution depends on the relative values of the lattice enthalpy and the hydration enthalpy.
If HE >> LE then the salt will
probably be soluble
Enthalpy Change of Hydration (∆ hyd Η)
Definition The enthalpy change when ONE MOLE of a gaseous ion dissolves in (an excess of) water (to give an infinitely dilute solution).
Values Exothermic
Example Na+(g) ——> Na+(aq) C l ¯(g) ——> C l ¯(aq)
Notes The polar nature of water stabilises the ions.
The greater the charge density of the ion, the greater the affinity for water and the more exothermic the process will be.
Na+^ -390 Mg2+^ -1891 Cl¯ -384 (all in kJ mol-1) K+^ -305 Ca2+^ -1561 Br¯ -
Enthalpy Change of Solution (∆ sol Η)
Definition The enthalpy change when ONE MOLE of solute dissolves in (an excess of) solvent (to give an infinitely dilute solution).
Values Mainly exothermic
Example NaC l (s) ——> NaC l (aq) [for ionic compounds, the ions will be dissociated]
Knockhardy Publishing
LATTICE ENTHALPY
HYDRATION ENTHALPY
M X¯^ + (s)
(g) (g) M+ + X¯
(aq) M^ + + X¯ (aq)
ENTHALPY OF SOLUTION
SOME USEFUL VALUES FOR THERMODYNAMIC CHANGES
Enthalpy change
- Values, which may be slightly different in other books, are in kJ mol- - Na+ -411 -381 -288 - of formation Cl¯ Br¯ I¯ O2- - K+ -437 -398 -328 - - Mg2+ -641 -524 -364 - - Ca2+ -796 -683 -534 -
- Enthalpy of atomisation H +218 Na +108 F + - C +716 K +89 Cl + - N +472 Mg +148 Br + - O +249 Ca + - Na +496 + Ionisation Energy 1st I.E. 2nd I.E. - Mg +738 + - Ca +590 + - K +419 +
- Electron Affinity F -348 C l -349 Br -342 I¯ - - O -
- 2nd Electron Affinity O +
- Check which definition Na+ -780 -742 -918 - Lattice Enthalpy Cl¯ Br¯ F¯ O2-
- is being used and use K+ -711 -679 -817 -
- appropriate sign for ∆ H Rb+ -685 -656 - - Mg2+ -2256 - - Ca2+ -
- Hydration Enthalpy Li+ -499 Be2+ -2385 F¯ - - Na+ -390 Mg2+ -1891 Cl¯ - - K+ -305 Ca2+ -1561 Br¯ - - Al3+ -4613 I¯ - - OH¯ -
