Entropy And Energetics

Ionic Bonding Recap

  • The bigger the charge on an ion the stronger the electrostatic force of attraction

  • More energy required to overcome these forces so they have a high M.P and B.P

  • The Smaller the ion the stronger the electrostatic force of attraction between ions

  • Smaller ions pack together more closely and more energy is required to overcome the stronger forces. Therefore, the M.P and B.P are higher for smaller ions

Theoretical and Experimental lattice Enthalpies

  • Theoretical lattice enthalpies are calculated assuming a perfectly ionic model

  • Most the time the positive ion distorts the charge distribution in the negative ion. We say that the positive ion polarises the negative ion

  • The more polarisation the more covalent character there will be

  • Below is an image of the perfectly ionic model

  • The difference between experimental and theoretical lattice enthalpy of formation can tell you how much covalent character an ionic compound has.

  • The bigger the difference in lattice enthalpy the more polarisation you have the greater the covalent character

Polarisability

  • Smaller cations are more polarising than larger ones

  • Smaller cations have a higher charge density as the charge is concentrated in a smaller area. The cation pulls electrons towards itself more readily

  • Large anions with a large charge are polarised much more easily than smaller, lower charged anions.

  • This is because the electrons are further away from the nucleus and there is more repulsion between the electrons in the ion. The electrons can be pulled towards the cations.

  • The more polarisation the more covalent character

  • Most polarisation occurs when the cation has a high charge but small radii and the anion has a large negative charge and large radii

Electronegativity

  • Electronegativity is the ability to attract electrons towards itself in a covalent bond

  • The bigger the difference in the electronegativity (using the Pauling’s scale) the more ionic the compound will be a difference of 0 is purely covalent

Enthalpy Change of Solution

  • The enthalpy change when 1 mole of an ionic substance is dissolved in the minimum amount of solvent to ensure no further enthalpy change is observed upon further dilution

  • For a substance to dissolve

    • Substance bonds must break (endothermic)

    • New bonds formed between the solvent and substance (exothermic)

  • Most ionic compounds dissolve in polar solvents like H2O. The H+ is attracted to the negative ions and the O- is attracted to the positive ions.

  • The water molecules surround the ions in a process called hydration

  • For this to happen the new bonds formed must be the same strength or greater than those broken.

  • If not then the substance is very unlikely to dissolve. Soluble substances tend to have exothermic enthalpies of solution for this reason

Enthalpy of Hydration

  • 2 things can affect the enthalpy change of hydration - Charge and Size of the ion

  • Charge:

    • Ions with a higher charge attract water molecules more strongly as the electrostatic attraction.

    • More energy energy is released when the bond is made which means they have a more exothermic enthalpy of hydration

    • The larger the charge the greater the enthalpy of hydration

  • Size:

    • Smaller ions have a higher charge density than larger ions

    • They can attract water molecules more strongly hence there is a more exothermic enthalpy of formation

    • The smaller the ion the greater the enthalpy of hydration

Entropy

  • Entropy is the measure of disorder in a system

  • It is the number of ways energy can be shared out between particles

  • The more disorder there is the higher the entropy

  • The number of particles also affects entropy change

  • If a reaction in the same state but more moles are produced then entropy increases. There are more ways energy can be distributed

  • A reaction can be spontaneous (feasible) even if it is enthalpically unfavourable for example endothermic reactions

  • A reaction will tend towards more disorder and hence increase entropy

  • Increasing entropy is energetically favourable and some reactions that are enthalpically unfavourable can still spontaneously react if changes in entropy overcome changes in enthalpy

  • This reaction is endothermic and enthalpically unfavourable (+164 kJmol-1)

  • 3 moles on the LHS and 13 moles on the RHS. Entropically favourable

  • Starting with 2 solids but making a gas and a liquid. Increased disorder so entropically favourable

Calculating Entropy

  • Entropy change (delta S) can be calculated between reactants and products.

  • Under standard entropy so:

    • 1 mole of substance

    • 100kPa

    • 298K

  • Total entropy change can be calculated using entropy changes of the system and surroundings

Gibbs free Energy

  • Tells us if a reaction is feasible or not

  • A reaction is feasible if the delta G is negative or zero

  • Even if the reaction is feasible you may not observe a reaction occurring. May be due to activation energy being too high or the rate of reaction being very slow

  • The temperature at which a reaction becomes feasible can be calculated

  • Just sub 0 in for delta G

  • Key point to remember is to convert delta H from kJ mol-1 to Jmol-1

  • Equilibrium constants are large when reversible reactions are feasible

  • When Delta G is negative and the reaction is theoretically feasible the equilibrium constants are greater than 1

  • When Delta G is negative and the reaction is not theoretically feasible the equilibrium constant are less than 1