3.6 enthalpy changes for solids and solutions Flashcards
in ionic bonding, the ions arrange themselves into a _____?
lattic
- so that the ions of opposite charge are next to one another
enthalpy change of formation definition
- the energy transferred when 1 mole of the compound is formed from its elements under standard conditions (298K and 100KPa) with all the reactants and products being in their standard states
is the enthalpy change of formation a positive or negative process?
negative
enthalpy of atomisation definition
- the enthalpy change when 1 mole of gaseous atoms is formed from its own elements in its standard state
is the enthalpy of atomisation a positive or negative process?
positive
first ionisation enthalpy definition
the enthalpy change when one mole of gaseous atoms forms one mole of gaseous 1+ ions
is the first ionisation enthalpy a positive or negative process?
positive
first electron affinity definition
the enthalpy change that occurs when 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms
is the first electron affinity a positive or negative process?
negative
second electron affinity definition
the enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ions
is the second electron affinity a positive or negative process?
positive
enthalpy of lattice formation definition
the standard enthalpy change when 1 mole of an ionic crystal lattice/solid ionic compound is formed from its constituent ions in gaseous form
is the enthalpy of lattice formation a positive or negative process?
negative
enthalpy of lattice dissociation definition
the standard enthalpy change when 1 mole of an ionic crystal lattice (ionic compound) is separated into its constituent ions in gaseous form
is the enthalpy of lattice dissociation a positive or negative process?
positive
categorise these:
1. Na (s) —> Na (g)
2. Na (g) —> Na+ (g) + e-
3. Na+ (g) + Cl- (g) —> NaCl (s)
4. O2 (g) —> 2O2 (g)
- enthalpy of atomisation
- first ionisation enthalpy
- enthalpy of lattice formation
- 2x enthalpy of atomisation
what are born-haber cycles?
- thermochemical cycles that include all of the enthalpy changes involved in the formation of an ionic compound
- lattice enthalpy cant be calculated directly so a born-haber cycle can be used to calculate the lattice enthalpy by applying hess’s law
drawing born-haber cycles:
- start with the elements in their standard states - these have zero enthalpy by definition
- the zero line will beer to be in the middle of the paper
- draw a line upwards from the zero line to the solid ionic compound made. this represents the enthalpy change of formation
- draw a lone upwards to usually atomise the elements individually
- positive (endothermic) changes are shown by an arrow pointing upwards
- negative (exothermic) changes are shown by an arrow pointing downwards
- need to include state symbols in born-haber cycles *
why is the second electron affinity positive while the first is negative?
because of the repulsion, so energy is required for the 2nd electron affinity a
what are some factors affecting the strength of the lattice formation enthalpy?
- size of ions
- the charges on the ions
how does the size of the ion affect the strength of the lattice formation enthalpy?
- the larger the ions, the less negative the enthalpies of lattice formation (weaker lattice)
- larger ions means the charges become further apart and have a weaker attractive force between them
- also larger ions means a smaller charge density
how does the charges on the ion affect the strength of the lattice formation enthalpy?
- the bigger the charge of the ion, the greater the attraction between the ions so the stronger the lattice enthalpy (more negative values)
- this means more energy will be required to separate the ions
what do born-haber cycles assume?
that an ionic compound is ‘perfectly ionic’:
- ions are treated as point charges, with the charge centred in the middle of the ion and 100% spherical
- attractions are assumed to be purely electrostatic
(the born-haber cycles use theoretical data)
why might there be a difference between theoretical and experimental values of enthalpy?
- theoretical values assumes an ionic model
how do you use the enthalpies of formation as a measure of a compound’s stability?
e.g calcium chloride - why does it have the formula CaCl2 and not CaCl or CaCl3
- need to calculate an enthalpy of formation for each case
- the one with the most exothermic enthalpy of formation will be the one that forms as it will be the most thermodynamically stable
do ionic solids dissolve in polar solvents?
yes
what is the lattice enthalpy?
- in order to dissolve an ionic compound, the lattice must be broken up
- this requires an input of energy - the lattice enthalpy
what is hydration in terms of enthalpies of solution?
the separate ions (once the lattice is dissolved) are then solvated by solvent molecules, usually water and cluster around the ions
- this is called hydration when the solvent is water
enthalpy of hydration definition
the enthalpy change when one mole of gaseous ions become aqueous ions
e.g Cu2+ (g) —> Cu2+ (aq)
is the enthalpy of hydration a positive or negative process?
negative
(it gives out energy (exothermic) because bonds are made between the ions and the water molecules)
the higher the charge density, the (greater/lower) the hydration enthalpy
greater
- as the ions attract the water molecules more strongly
e.g compare the enthalpies of hydration of Na+ and Mg2+
- Mg2+ has stronger attraction
- more exothermic
- bc bond making is exothermic
??? i think
e.g compare the enthalpies of hydration of Ca2+ and Mg2+
- Ca2+ less exothermic
- Ca2+ has lower charge density (bc bigger)
- weaker attraction between water and Ca2+
enthalpy of solution definition
- the standard enthalpy change when 1 mole of an ionic solid dissolves in a large enough amount of water to ensure that the dissolved ions are well separated and do not interact with one another
e.g NaCl (s) —> Na+ (aq) + Cl- (aq)
you can link the enthalpy of solution, the enthalpy of hydration and lattice enthalpy in a Hess’s law cycle
we can think of dissolving an ionic compound as the sum of 3 processes:
- breaking the ionic lattice up to give gaseous ions
- hydrating the positive ions
- hydration the negative ions
therefore, what can the equation be for dissolving an ionic compound?
∆H soln = -∆H latt + ∆H hyd
exam questions tend to focus on you working out the lattice formation/dissociation enthalpy in a Hess’s law cycle and then use that value to calculate solution/hydration enthalpies
e.g calculate the enthalpy of solution of NaCl given that the lattice enthalpy of formation of NaCl is -771 kJmol^-1 and the enthalpies of hydration of sodium and chloride ions are -406 and -364 kJmol^-1 respectively
∆H soln = -∆H latt + ∆H hyd
= - (-771) + (-406 + - 364)
= 1 kJmol^-1
what does ∆H solution tell us?
- in general, the substance is more likely to be soluble if the ∆H solution is exothermic
- however its not always the case as e.g NaCl = 1 and can dissolve in water
- (bc whether a compound is soluble or not also depends on other thermodynamic quantities such as entropy change
- when a solid dissolves into ions the entropy increases
- this positive entropy can make ∆G negative even if ∆H solution is endothermic, especially at higher temperatures )
what type of enthalpy change is taking place in the following reaction?
Na+ (g) —> Na+ (aq)
enthalpy change of hydration
what type of enthalpy change is taking place in the following reaction?
Mg2+ (g) + 2Cl- (g) —> MgCl2 (s)
lattice enthalpy of formation
what type of enthalpy change is taking place in the following reaction?
1/2Cl2 (g) —> Cl (g)
enthalpy of atomisation
what affects the solubility of ionic compounds in water?
- the balance between the hydration enthalpy of the ions in the compound and the lattice dissociation enthalpy of the compound
what factors affect enthalpy of hydration?
- depends on the amount of attraction between the ions and the water molecules
- the bigger the charge and the smaller the ion, the larger the enthalpy of hydration
what is Hess’s law?
the enthalpy change of a reaction is independent of the route taken
how can the enthalpy change of a solution be calculated using an energy cycle?
- lattice breaking enthalpy and enthalpy change of hydration can be used in an energy cycle to calculate the enthalpy change of solution using hess’s law
give the equations for the first electron affinity of oxygen and the first ionisation energy of magnesium:
first electron affinity of oxygen:
O (g) + e- —> O- (g)
first ionisation energy of magnesium:
Mg (g) —> Mg+ (g) + e-
what is the perfect ionic model?
- the perfect ionic model assumes that all ions are perfectly spherical and that the bonds have no covalent character
why are theoretical lattice enthalpies often different from experimental ones?
- theoretical lattice enthalpies use calculations which are based on a perfect ionic model of the lattice
- the experimental values often differ bc most ionic compounds have some covalent character
- the more polarisation in the ionic bond, the more covalent character of the compound
how does exothermicity or endothermicity of ∆fH (enthalpy change of formation) act as a qualitative indicator of a compound’s stability?
- for an exothermic reaction the products are more stable than the reactants
- for an endothermic reaction the reactants are more stable than the products
