Module 5 - (chapter 22) Enthalpy and Entropy Flashcards
lattice enthalpy
- the enthalpy change the accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions
- it is exothermic and the value will always be negative
standard enthalpy change of formation
- enthalpy change that takes place when one mole of a compound is formed from its constituent elements under standard states
- the compound will always be an ionic compound in its solid lattice
standard enthalpy change of atomisation
- the enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions
- always endothermic because bonds are broken to form gaseous atoms
first ionisation energy
the enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to for one mole of gaseous 1+ ions
-endothermic because energy is required to overcome the attraction between negative electron and positive nucleus
electron affinity
- the enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one moles of gaseous 1- ions
- exothermic because the electron being added is attracted towards the nucleus
successive electron affinities
- occurs when an anion has a greater charge than 1- such as O2-, successive electron affinities are required
- they are endothermic as a second electron is being gained by a negative ion, which repels the electron away. energy must be put in to 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 ion from the ionic lattice finish surrounded with water molecules as AQ ions
- can be exothermic or endothermic
how does something dissolve? (e.g. sodium chloride)
- the partial + and - shares in the water molecules are attracted towards the positive and 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
what is important to note in MCAT
- make sure you consider what is changing temperature
- e.g. solution may be what is dissolved plus water
- if you don’t use what is dissolved you would have obtained finally enthalpy charge
what two types of energy are involved in the dissolving process
- the ionic late breaks up
- water molecules are attracted to and surround the ions (hydrated)
- 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
- the separate gaseous ions interact with polar water molecules to form hydrated aqueous ions. (ECoH)
enthalpy change of hydration
the enthalpy change that accompanies the dissolving of gaseous ions in water to form one moles of aqueous ions.
exothermic as we are forming bonds and energy is given out as water molecules bond to the metal ions
enthalpy change of solution + equation
can be exothermic or endothermic, depending on the relative sizes of the lattice enthalpy and the enthalpy change of hydration
is equal to (-lattice enthalpy) + hydration
borne- harbour route for dissolving
lattice enthalpy down arrow from gaseous to lattice
hydration down from gaseous to aqueous
dissolving up from lattice to aqueous
general properties of ionic compounds
-high melting and boiling points
some can be melted using a bunsen burned, others are so high that they can be used to coat the inside of furnaces
-soluble in polar solvents
-conduct electricity when molten or in aq solution
-some are soluble in polar solvents (e.g. water), others are insoluble
how does ionic size effect lattice enthalpy
- as you go down group 1 ionic radius increases
- attraction between ions decreases
- less energy needed to overcome forces
- lattice energy less exothermic (less negative9
- melting point decreases.
how does ionic charge impact lattice enthalpy
- as the ionic change increases
- the attraction between ions increases
- lattice energy becomes more negative (more exothermic)
- melting point increases
Na2+ to Al3+
- two supporting effects
- increasing charge gives more attraction
- decreasing size gives more attraction
cl- to S2-
- two opposing effects
- increase change gives more attraction
- increasing size gives less attraction
predicting melting points
- magnitude of lattice enthalpy indicated melting point
- MgO for example has very exothermic lattice enthalpies and very high mps
- these so state used to coat indie of furnaces
predicting solubility
- the attraction between the ions in the lattice must be overcome
- water molecules are attracted to the positive and negative ions, surrounding them and releasing energy equal to hydration enthalpy
- yet, compounds with + enthalpy changes are also soluble
entropy
used to describe the dispersal of energy within e chemical system which is greater the more disordered it is
- the greater the entropy, the greater the dispersal of energy and the greater the disorder and the greater the that energy is spread out per kelvin per mole
- units are JK-1mol-1
general entropy rules
- solids have the smallest entropies
- liquids have greater entropies
- gases have the greatest entropies
- this is because melting and boiling increase the randomness of particles meaning energy is spread out more and is thus more positive
- however even states have substances thatch have very different entropies
predicting entropy values
- at 0K there would be no energy and all substances would have zero entropy
- above 0K energy becomes dispersed
- if a system changes to become more random, energy can be spread out more and the entropy change will be positive
- if a system changes to become less random, energy becomes more concentrated meaning the entropy change will be negative
how does moles affect entropy
- if the products have more gas molecules entropy increases
- disorder increases so energy more spread out and more positive
standard entropies
entropy of one mole of a substance under standard conditions
-always positive
standard entry change
sum of standard entropy of products - sum of standard entropy of reactants
feasibility
used to describe whether a reaction is able to happen and is energetically feasible
-also known as spontaneous
-for a reaction to be feasible deltaG must be less than 0
-overall energy of the products must be lower than reactants
-just because reaction is feasible does t mean it reacts
spontaneously at room temp, activation energy must be put in
free energy
- the overall change in energy during a reaction
- delta G
- made of the enthalpy change (the heat transferred between chemical system and the surroundings) and entropy change (the dispersal of energy within a chemical system)
the Gibbs equation
deltaG = deltaH - Tx deltaS
conditions for feasibility
- the value of deltaH is usually much larger than deltaS so dominated the equation (usually given in KJ-1mol-1)
- as temperature increases the T x deltaS term becomes more important
delta H Negative
delta S positive
delta G is negative if T is low or high
reaction is feasible
delta H positive
delta S negative
dealt G positive if T is low or high
not feasible
delta H Negative
delta S negative
when T is low delta G is negative, so feasible
when T is high delta G is positive, so not feasible
delta H positive
delta S positive
when T is low delta G is positive, so not feasible
when T is high delta G is negative, so feasible
which reactions tend to be more feasible
exothermic
- delta H will be negative meaning delta G is ore likely to be negative
- if endothermic, the delta S or T must be very high to make delta G negative
limitations of Delta G
- indicated thermodynamic feasibility but doesn’t account for kinetics or rate of reaction
0e. g. hydrogen peroxide is negative but large activation energy so reaction very sow. MnO2 is added as a catalyst
