A summary of topic 13, organised so the notes are easy to understand. The notes are on slides, so they can be printed out and used as revision cards or posters, for revision on the go. The notes cross-reference the specification so it is easy to see where each bit of information has come from. They...
Enthalpy definitions: Factors that affect the magnitude of lattice enthalpy of formation
Enthalpy of atomisation = enthalpy Larger -ve value = ↑ strength of ionic bond
change when 1 mole of gaseous atoms - Charge size - - Mg carries 2x more charge than Na ∴ more -ve (larger)
is made from the element in its SS
under SC (endothermic) e.g. ⤷ ↑ charge size = ↑ ES attraction = large -ve value = ↑ strength of ionic bond
First electron affinity = enthalpy change - Cation to anion interactions - - more interactions because 2x more Cl-
when 1 mole of electrons is added to 1 per cation ∴ more -ve (larger)
mole of gaseous atoms under SC ⤷ ↑ interactions (no. of atoms) = ↑ ES attraction = large -ve value = ↑ strength
(exo for most) e.g. of
Second electron affinity = enthalpy ionic bond
change when each ion in 1 mole of - Sum of ionic radii’s - - Mg^2+ smaller than Na^+ ∴ ↓ sum of ionic radii ∴
gaseous 1- ions gains 1 electron to form more -ve (larger)
1 mole of 2- ions (endo as -ve e- added ⤷ ↓ sum of ionic radii = closer packing of ions = stronger electrostatic attraction
to -ve ∴ requires energy) e.g. =
First ionisation = the enthalpy change large -ve value = ↑ strength of ionic bond
when each atom in 1 mole of gaseous Born haber cycles
atoms loses 1 e- to form 1 mole of = the overall energy changes that take place when an ionic compound is made from
gaseous 1+ atoms (endo) its elements
e.g. Arrow going down = -ve (exo)
Second ionisation = enthalpy change Arrow going up = +ve (endo)
when each ion in 1 mole of gaseous 1+ Lattice enthalpyinoftheir
1. Elements formation
standard=states
enthalpy change when 1 mole of solid ionic lattice is
and balanced
ions loses 1 e- to form 1 mole of formed
2. Atfrom its constituent
the bottom ions
= final stage in theenthalpy
of lattice gaseousof states under SC (exo)
gaseous 2+ atoms (endo) formation = solid ionic lattice. From elements in
e.g. standard states → solid ionic lattice = ∆H f (-ve
Lattice enthalpy of formation = enthalpy ∴ go down)
3. Turn both elements (separately) into 1 mole of
change when 1 mole of solid ionic
gaseous atoms using ∆Hat (+ve)
lattice is formed from its constituent ions 4. Remove 1 e- from your gaseous metal to form
in the gaseous states under SC (exo) a +ve metal ion using ∆H1stIE (+ve)
e.g. ⤷ can follow this by removing a 2nd/3rd e-
Lattice enthalpy of dissociation = 5. Add 1 e- to your gaseous non-metal to form a -
enthalpy change when 1 mole of a solid ve ion using using ∆H1stEA (-ve∴ go down)
⤷ can follow this by adding a 2nd/3rd e-
ionic lattice is broken up into its
⤷ 2nd & 3rd EA are normally +ve ∴ go up
constituent ions in the gaseous states 6. From ions in gaseous states to solid ionic If there are balancing numbers, multiply by them
(endo) e.g. lattice = ∆Hlatt (-ve ∴ go down)
, Factors affecting lattice enthalpy Standard enthalpy of solution
Theoretical = the enthalpy change when 1 mole of an ionic compound is dissolved in
- Perfectly symmetrical water to produce an infinitely dilute solution of aq ions (under SC)
- Even charge distribution E.g. 1 mole of NaCl(s) → Na+(aq) + Cl-(aq)
- No covalent character (perfectly touching no This happens in two steps:
overlap) When an ionic solid dissolves in aq solution:
Experimental (measured/actual) lattice energy 1. ΔHLED - Lattice breaks down into gaseous ions
- Calculated from experimental values - Opposite of ΔHLEF (opposite sign, endo = +ve)
Agreement - ΔHLED = -ΔHLEF
Nearer the actual is to the theoretical = more Standard enthalpy of lattice dissociation (ΔH LED)
perfect ionic character - ‘agreement’ = enthalpy change when 1 mole of an ionic compound is broken up into its
Further away = ↑ covalent constituent ions (under SC)
↑ covalent character if anion is large & highly NaCl(s) → Na+(g) + Cl-(g)
charged 1. ΔHhyd - Gaseous atoms become hydrated by water
- Large charge & size of anion = more easily - Always exothermic
polarised Standard enthalpy of hydration (ΔHhyd)
↑ covalent character if cation is small & highly = the enthalpy change when 1 mole of isolated gaseous ions is dissolved in
charged water, converting it to 1 mole of aq ions (under SC)
- Polarising power = ability of a cation to Na+(g) → Na+(aq)
attract electrons Cl-(g) → Cl-(aq)
- ↑ charge density = ↑ polarising power = ↑ Factors affecting hydration enthalpy:
polarisation of the anion - Charge density (ionic radius & charge)
Exp-theo = % difference - Smaller ionic radii = ↑ charge density = stronger ion-dipole attractions
Theo between water and the ions in solution = more energy released when
Covalency in bonding they are hydrated
- Caused by the polarisation of the anion by - Large ionic charges = ↑ charge density = stronger ion-dipole
the cation attractions between water and the ions in solution = more energy
- ∴ distortion of e- density in anion (↑ e- near released when they are hydrated
cation)
- ∴ some e- density exists between the 2 ions
- ∴ a degree of covalency
If ∆Hlatt is large & -ve (more exo) = stronger ionic
If there’s a bigger difference between exp & theo =
more covalent character
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