Enthalpy & entropy

• Lattice enthalpy --> the enthalpy change due to the formation of 1 mole of ionic lattice from gaseous ions.

  • Provides a measure of ionic bond strength
  • It is always exothermic (as it is bond formation)
  • Na+(g) + Cl-(g) --> NaCl(s)
• Standard enthalpy change of formation --> the enthalpy change due to 1 mole of a compound forming from its constituent elements under standard conditions

  • This is exothermic also
  • Na (s) + 1/2 Cl2(g) --> NaCl (s)
• Standard enthalpy change of atomisation --> the enthalpy change by forming of 1 mole of gaseous atoms from the element in its standard state under standard conditions

  • This is always endothermic (as bond breaking)
  • Na(s) --> Na(g)
  • However if the element is a gas this value is related to bond enthalpy i.e: 1 mole of bond breaking
  • 1/2 Cl2(g) --> Cl(g)
• First ionisation energy --> the enthalpy change required to remove one electron from each atom in 1 mole of gaseous atoms to form one mole of +1 gaseous ions

  • Always endothermic (as energy need to overcome nuclear pull)
  • Na (g) --> Na+(g) + e-
  • For positive ions with larger charges each successive ionisation energy is added to the previous
• First electron affinity --> the enthalpy change when an electron is added to each atom in 1 mole of gaseous atoms to form one mole of -1 gaseous ions

  • Always exothermic (as electron pulled towards nucleus)
  • Cl (g) + e- --> Cl- (g)
  • Successive electron affinities are needed for greater charges
  • Endothermic as electron is repelled by negative ion
  • O (g) + 2e- --> O2- (g) = O (g) + e- --> O- + O- (g) + e- --> O2-
• Born - Haber cycle --> can be used to determine enthalpy changes concerning ionic lattices
• Factors affecting lattice enthalpy:

  • As ionic radius increases the lattice enthalpy becomes less negative as the attraction between ions will decrease - lowering melting points
  • As ionic charges increase the lattice enthalpy becomes more negative as the attraction between ions will increase - raising melting points
• Successive electron affinities:

  • When an anion has a greater charge than -1, such as O2-, successive electron affinities are required
  • Second electron affinities are endothermic
  • A second electron is being gained by a negative ion, which repels the electron away. So energy must be put into force the negatively charged electron onto the negative ion
• Standard enthalpy change of solution --> the enthalpy change that takes place when one mole of a solute dissolves in a solvent

  • If the solvent is water, the ions from the ionic lattice finish up surrounded with water molecules as aqueous ions
  • Na+Cl- (s) + aq --> Na+ (aq) + Cl- (aq)
  • The enthalpy change of solution can be exothermic or endothermic
• • NaCl(s) : Na+ & Cl- ions attracted together in a giant ionic lattice
  • NaCl (aq): Na+ & Cl- ions are separate, but now surrounded by water molecules
• • In the aqueous ions, the delta positive and delta negative partial charges in the water molecules are attracted towards the positive & negative ions
  • The delta negative oxygen atom is attracted to the positive sodium ion
  • The delta positive hydrogen atom is attracted to the negative chloride ion
• The dissolving process:

  When a solid ionic compound dissolves in water, 2 processes take place:

  • the ionic lattice breaks up
  • water molecules are attracted to, and surround, the ions
  • Two types of energy changes are involved
• Energy changes in dissolving process:

  1. The ionic lattice is broken up forming separate gaseous ions. This is the opposite energy change from lattice energy, which forms the ionic lattice from gaseous ions
  2. The separate gaseous ions interact with polar water molecules to form hydrated aqueous ions. The energy change involved is called the enthalpy change of hydration.
• Enthalpy change of hydration --> enthalpy change due to dissolving gaseous ions in water to form 1 mole of aqueous ions.

  Na+ (g) + aq --> Na+ (aq)
• Factors affecting enthalpy of hydration:

  • As ionic radius increases the hydration enthalpy becomes less negative as the attraction between ion & water decreases
  • As ionic charge increases the hydration enthalpy becomes more negative as the attraction between ion & water increases
  • If hydration enthalpy is greater than the lattice enthalpy then dissolving is likely to happen, i.e: is the enthalpy of solution is exothermic
• Predicting solubility:

  • To dissolve an ionic compound in water, the attraction between the ions in the ionic lattice must be overcome. This requires a quantity of energy equal to the lattice enthalpy.
  • Water molecules are attracted to the positive & negative ions, surrounding them and releasing energy equal to hydration enthalpy
  • If the sum of the hydration enthalpies is larger than the magnitude of the lattice enthalpy, the overall enthapy change ( the enthalpy change of solution) will be exothermic and the compound should dissolve
• Entropy --> the dispersal of energy within the chemicals making up the chemical system

  • The greater the entropy, the greater the dispersal of energy and the greater the disorder
  • The units of entropy are JK^-1mol^-1. The greater the entropy value, the greater that energy is spread out per Kelvin per mole
• Entropies in states:

  • Solids have the smallest entropies
  • Liquids have greater entropies
  • Gases have the greatest entropies
• Predicting entropy changes:

  • Above 0K, energy becomes dispersed amongst the particles and all substances have positive entropy
  • Systems that are more chaotic have a higher entropy value
  • If a system changes to become more random. energy can be spread out more - there will be an entropy change which will be positive.
  • If a system changes to become less random, energy becomes more concentrated - the entropy change will be negative
• Changes of state:

  • Entropy increases during changes in state that give a more random arrangement of particles
  • Solid --> liquid --> gas
  • When any substance changes from solid to liquid to gas, its entropy increases
  • Melting & boiling increases the randomness of particles
  • Energy is spread out more and entropy change is positive
• Change in the number of gaseous molecules:

  • Reactions that produce gases result in an increase in entropy
  • CaCO3 (s) + 2HCl (aq) --> CaCl2 (aq) + CO2 (g) + H2O (l)
  • Production of a gas increases the disorder of particles
  • Energy is spread out more and entropy change is positive
  • N2 (g) + 3H2 (g) --> 2NH3 (g) = 4 moles of gas --> 2 moles of gas
  • There is a decrease in the randomness of particles
  • The energy is spread out less and entropy change is negative
• Standard entropies:

  • The standard entropy of a substance is the entropy of one mole of a substance, under standard conditions (100KPa and 298K)
  • Standard entropies have units of JK^-1mol^-1
  • Standard entropies are always positive
• Entropy changes = sum of entropy change of product - sum of entropy change of reactants
• • The feasibility of a reaction is used to describe if a reaction will be able to happen
• The free energy change ( delta G) during a reaction is the overall energy change caused by the:

  • Heat transfer between the system and surroundings
  • Dispersal of energy within the system at the temperature of the reaction
• Gibb's equation:

  • Free energy change = enthalpy change with surroundings - (temperature x entropy change of system)
  • delta G = delta H - T x delta S
• • A reaction is feasible when delta G is less than zero
  • Reaction occurs when delta G is a negative value
• • As enthalpy change is often a magnitude larger than entropy change it usually determines the feasibility
• • At room temperature the value for free energy change is more dependent on enthalpy change but as temperature increases T x delta S starts to dominate
• Limitations of predictions made for feasibility:

  • Large activation energies lead to slow rates
  • Doesn't take into account the kinetics or rate of reaction