pmt thermodynamics Flashcards
What does Hess’s Law state?
The enthalpy change for a reaction is independent of the route taken
Define standard enthalpy of formation.
The enthalpy change when one mole of a compound is formed from its constituent elements in standard conditions, with all products and reactants in their standard states.
What is the standard enthalpy of an element?
Zero, by definition.
Define standard enthalpy of combustion
The enthalpy change when one mole of a substance is completely burnt in (excess) oxygen
Define standard enthalpy of atomisation
Enthalpy change when one mole of gaseous atoms is formed from a compound in its standard state in standard conditions.
Define first ionisation energy
Enthalpy change when one mole of electrons is removed from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Define second ionisation energy.
Enthalpy change when one mole of electrons is removed from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions
Define first electron affinity
Enthalpy change when one mole of gaseous atoms gains one mole of electrons to form one mole of gaseous 1- ions.
Define second electron affinity.
Enthalpy change when one mole of gaseous 1- ions gains one mole of electrons to form one mole of gaseous 2- ions
Define lattice enthalpy of formation
Enthalpy change when one mole of solid ionic lattice is formed from its constituent gaseous ions.
Define lattice enthalpy of dissociation.
Enthalpy change when one mole of solid ionic lattice is dissociated (broken into) into its gaseous ions
Define enthalpy of hydration.
Enthalpy change when one mole of gaseous ions become hydrated/dissolved in water to infinite dilution [water molecules totally surround the ion]
Define enthalpy of solution
Enthalpy change when one mole of solute dissolves completely in a solvent to infinite dilution.
Define mean bond dissociation enthalpy
Enthalpy change when one mole of (a certain type of) covalent bonds is broken, with all species in the gaseous state
What factors affect the lattice enthalpy of an ionic compound?
Size of the ions, charge on the ions
How can you increase the lattice enthalpy of a compound? Why does this increase it?
Smaller ions, since the charge centres will be closer together. Increased charge, since there will be a greater electrostatic force of attraction between the oppositely charged ions. N.B. Increasing the charge on the anion has a much smaller effect than increasing the charge on the cation, since increasing anion charge also has the effect of increasing ionic size.
How can Born-Haber cycles be used to see if compounds could theoretically exist?
Use known data to predict certain values of theoretical compounds, and then see if these compounds would be thermodynamically stable. Was used to predict the existence of the first noble gas containing compound.
What actually happens when a solid is dissolved in terms of interactions of the ions with water molecules?
Break lattice → gaseous ions; dissolve each gaseous ion in water. The aqueous ions are surrounded by water molecules (which have a permanent dipole due to polar O-H bond)
What is the perfect ionic model?
Assumes that ions are perfectly spherical and that there is an even charge distribution (100% polar bonds). Act as point charges.
Why is the perfect ionic model often not accurate?
Ions are not perfectly spherical. Polarisation often occurs when small positive ions or large negative ions are involved, so the ionic bond gains covalent character. Some lattices are not regular and the crystal structure can differ.
Which kind of bonds will be the most ionic? Why?
Between large positive ions and small negative ions e.g. CsF
Define the terms spontaneous and feasible
If a reaction is spontaneous and feasible, it will take place of its own accord; does not take account of rate of reaction.
Is a reaction with a positive or negative enthalpy change more likely to be spontaneous?
Negative - exothermic
Define entropy
Randomness/disorder of a system. Higher value for entropy = more disordered
What units is entropy measured in?
JK-1mol-1
What is the second law of thermodynamics?
Entropy (of an isolated system) always increases, as it is overwhelmingly more likely for molecules to be disordered than ordered
Is a reaction with positive or negative entropy change more likely to be spontaneous?
Positive - reactions always try and increase the amount of disorder
Compare the general entropy values for solids, liquids and gases
Solids < liquids < gases
What does the value for Gibbs free energy for a reaction show?
If G < 0, reaction is feasible. If G = 0, reaction is JUST feasible. If G > 0, reaction is not feasible.
What is the significance of the temperature at which G = 0?
This is the temperature (in Kelvin) at which the reaction becomes feasible.
What are the limitations of using G as an indicator of whether a reaction will occur?
Gibbs free energy only indicates if a reaction is feasible. It does not take into account the rate of reaction (the kinetics of the reaction). In reality, many reactions that are feasible at a certain temperature have a rate of reaction that is so slow that effectively no reaction is occurring.
Why is entropy zero at 0K?
No disorder - molecules/atoms are not moving or vibrating and cannot be arranged in any other way. Maximum possible state of order
What are the two key things to look out for to decide if entropy increases/decreases/stays
relatively constant?
Number of moles - more moles made → increase in entropy Going from solid → liquid/gas or liquid → gas
How is it possible for the temperature of a substance undergoing an endothermic reaction to stay constant?
The heat that is given out escapes to the surroundings
