5D- Lattice enthalpy and free energy

define lattice enthalpy

the energy change that accompanies the formation of one mole of ionic compounds from gaseous ions under standard conditions

are lattice enthalpies exothermic or endothermic

exothermic (negative) as they involve bond formation

define enthalpy change of formation

the enthalpy change when one mole of a compound is formed from its elements under standard conditions, with all reactants and products in their standard states

define enthalpy change of atomisation

the enthalpy change that takes place when one mole of gaseous atoms forms from the element in its standard states

is enthalpy change of atomisation exothermic or endothermic

endothermic (positive) as bonds are being broken

define first ionisation energy

the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions

define first 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 mole of gaseous 1- ions

is first electron affinity exothermic or endothermic

exothermic (negative) because the electron being added is attracted to the positive nucleus. the nuclear attraction is stronger than the electron repulsion

are second/third electron affinities exothermic or endothermic

endothermic (positive) the electron is being added to a negative ion so is repelled. energy is required to overcome this repulsion

how can lattice enthalpy strengths be compared

based on strength of ionic bond. stronger attractions in lattices with smaller ions and higher charged ions

what do Born-Haber cycles look like

what does Hess’s Law state

if a reaction can take place through more than one route and the initial and final conditions are the same, the total enthalpy change is the same for each route

in a Born-Haber cycle for enthalpy of solution what should be at the top

gases as they have the most energy

define standard enthalpy change of solution

the enthalpy change that takes place when one mole of a compound is completely dissolved in water under standard conditions

define standard enthalpy change of hydration

the enthalpy change that takes place when one mole of isolated gaseous ions is dissolved in water forming one mole of aqueous ions under standard conditions

is standard enthalpy change of hydration exothermic or endothermic

exothermic as new electrostatic attractions are being made on hydration

is standard enthalpy change of solution exothermic or endothermic

either

what happens when a solid dissolves

the ionic lattice breaks down into gaseous ions, the ions become hydrated

what does the value of enthalpy change of solution depend on

the balance between the ionic lattice breaking down into gaseous ions and the ions becoming hydrated

what factors affect lattice enthalpy

ionic charge and ionic radius

how will lattice enthalpy change as radius of a cation decreases

it will become more exothermic as the charge will be distributed over a smaller volume so the attraction for the negative ion will increase

how will lattice enthalpy change as radius of an anion increases

lattice enthalpy will be less exothermic as the charge is distributed over a larger volume so the attraction for the positive ions will decrease

what factors affect enthalpy change of hydration

size of ions and charge of ions

what type of ions have the most exothermic lattice enthalpy/enthalpy of hydration

small ions with high charges

what is the process of hydration

forming bonds with water molecules, the stronger these bonds the more negative the enthalpy of hydration would be 

why do ionic charge and ionic radius affect the magnitude of lattice enthalpy and enthalpy of hydration

they both affect charge density- the amount of electric charge per volume, ions with small radii and high charges will have the greatest charge density

what is the relationship between the charge on the ions, the size of the ion and the charge density

charge density= ionic charge/ionic volume

how does lattice enthalpy change as the charge density on an ion decreases

lattice enthalpy becomes more negative

define entropy

a measure of the dispersal of energy is a system, the more disordered a system the greater the entropy

what is the symbol for entropy

S

what are the units for entropy

J K^-1 mol^-1

is entropy positive or negative

positive

how do you calculate entropy change

sum of entropy of products - sum of entropy of reactants

are changes in entropy positive or negative

either

are positive or negative changes in entropy favoured

positive- processes which involve an increase in disorder

What does Gibbs free energy show

ΔG, used to quantitatively decide if a reaction is feasible

what is a feasible reaction

one that occurs without a continuous supply of energy

what value of ΔG are processes feasible

if ΔG is negative

what is the equation for Gibbs free energy

ΔG = ΔH - (T x ΔS)

ΔG- Gibbs free energy (kJ mol^-1)

ΔH- enthalpy change (kJ mol^-1)

T- temperature (K)

ΔS- entropy change (kJ K^-1 mole^-1)

why might a feasible reaction not appear to occur

rate of reaction is too slow due to a high activation energy

what is there is a process requires a high or low temperature in order to be feasible

there will be a minimum or maximum temperature at which the process is feasible, the change occurs at the point that ΔG=0

enthalpy change and entropy change signs and temperature and whether reaction is feasible

how the Gibbs free energy equation be rearranged for graphical use

ΔG = -ΔS x T + ΔH

y = m x + c

in a graph of Gibbs free energy against temperature what does the y-intercept indicate

ΔH (enthalpy change)

in a graph of Gibbs free energy against temperature what does the x-intercept indicate

the temperature when the reaction become feasible

in a graph of Gibbs free energy against temperature what does the gradient represent

-ΔS (negative entropy change)