module 5

5.1 static and dynamic equilibrium

enthalpy (ΔH): energy required by/released in a reaction.

• exothermic (-ΔH): energy released/graph ends lower than start

• endothermic (+ΔH): energy absorbed/graph ends higher than start

• ΔH reaction = ΔH products - ΔH reactants

entropy (ΔS): amount of disorder in system, increases as particles become more randomly dispersed

• 2nd law of thermodynamics: total entropy of Universe is increasing

Gibbs Free Energy (ΔG): thermodynamically favoured/spontaneous

• calculated by ΔG = ΔH - TΔS, where T = temperature

• ΔG < 0, products are favoured, spontaneous reaction

• ΔG > 0, reactants are favoured, non spontaneous

• ΔG = 0 is equilibrium

chemical equilibrium: rate of forward reaction equal to rate of reverse reaction

static equilibirum: very rare, forward and back rates both equal to 0. equilibirum of irreversible reactions

dynamic equilibrium: forward and reverse reaction rates equal, not equal to 0. can appear unchanging

a system at equilibrium can’t produce any energy. a non-equilibrium system is deliberately and permanently prevented from reaching equilibrium eg. combustion

all reactions will move towards equilibrium

homogenous equilibrium: all same state

collision theory

collision theory: particles must collide to react with each other

successful collision: must collide with force higher than required activation energy, and at correct orientation for bonds to be formed

rate of reaction = number of successful reactant collisions = Δconc/time

conc, pressure, volume, temp, SA, catalysts affect rate

activation energy = Ea

Maxwell-Boltzmann distribution curve: probability curve. shows number of particles that have a particular kinetic energy.

small area under curve = small number of molecules

reaction rate

temperature: increase = increase in rate, increases kinetic energy

concentration: increase = increase, more particles increase chance of collision

pressure: increase = increased reaction rate as greater collisions with walls. increasing temp/conc is same as increasing pressure

surface area: more exposed surface area = more successful collisions = greater reaction rate

catalysts: alternative pathway for reaction to occur, reduces activation energy required. therefore more successful collisions, but not more overall collisions = increased reaction rate. often also impact orientation, therefore more likely to be correct collision

known non-equilibrium reactions

• generally non-equilibrium systems are open and consist of irreversible processes under the conditions at which they operate

• non-equilibrium systems are perpetually moving away from equilbrium. ΔG is positive

• combustion: reverse reaction has a highly positive ΔG at any temperature; therefore reverse reaction won’t occurr

• photosynthesis: large overall negative change in enthalpy and large positive change in entropy = overall negative gibbs free energy, can occur spontenously

5.2 factors that affect equilibrium

dynamic equilibrium: ΔG=0. constant concentrations, but not necessarily equal

Le Chatelier’s Principle

stress: change imposed on equilibrium that changes conditions of the system and alters position

equilibrium position: relative amounts of reactants/products

if stress disturbs equilibrium balance, it will shift to restore balance (in direction that counteracts the stress)

any stress applied to system at equilibrium can be viewed as an effect on concentration or temperature

le chatlier’s principle states the system will shift to restore balance/shift in the direction that will counteract the stress

temperature: depends on the forward/reverse enthalpies. if temperature is increased, the reaction will move in the direction of +ΔH to ‘cool’ and vice versa

concentration: if concentration of all gases increases (eg. volume decrease, pressure increase), will shift to produce less gas. if concentration of one gas increases, will shift to produce less of that specific gas. if inert gas is added, no change due to no impact on concentration overall

eg. N2 + 3H2 → 2NH3

increase volume = shift left

N2/H2 added = shift right

equilibrium & collision theory

collision theory allows an explanation of why le chatlier’s principle occurs

temperature:

when heated, all reaction rate will increase, but proportionally more in reverse reaction (endothermic reaction). this is because the higher temperature will increase the number of particles with sufficient energy by a greater proportion for the direction with the highest activation energy

decreasing will favour the direction with the lower activation energy (exothermic reaction). this is because the lower temperature will decrease the number of particles with sufficient energy by a greater proportion for the direction with the highest activation energy

concentration: increase in concentration of reactants shifts equilibrium to favour products/foward reaction increases compared to the reverse. percentage of successful collisions remains constant!!!! total number of collision increases. decreasing concentration reactants = products are more likely to collide, therefore reverse reaction is favoured until new equilibrium is found

pressure:

• if increased due to temperature, use previous explanation

• for gases,concentration is considered in terms of partial pressure. partial pressure determined by imagining one gas is only gas in container ie. how many times will it collide with walls. increasing number of particles of any of the reactant/product gases will increase partial pressure/concentration and collision theory can be applied. converse applies for removing.

• adding inert (non-reacting) gas will not alter

volume:

• if gas only on one side of reaction: decreasing volume will increase concentration of products/reactants and equilibrium will shift accordingly

• if gas on both sides: changing volume will change concentration of products and reactants by same factor. the side with greater stoichiometric change will have more frequent succesful collisions and the equilibrium will shift away

  eg. N2O4 (g) → 2NO2

  2 mol 4 mol

if you reduce volume of both by half, goes up by factor of 2. therefore 8 mol NO2 and 4 mol N2O4. as NO2 gone up by 4, equilibrium shifts left/reverse

catalyst: lower activiation energy. equilibrium will be reached more quickly, but will not change position

equilibrium graphs

• 0 gradient indicates equilibrium has been established

• vertical line indicates stress has ultimately altered concentration, directly (adding/removing) or indirectly (altering pressure or volume)

• smooth curved sections = le chatlier change

• temperature = smooth curved response

5.3 calculating the equilibrium constant

equilibrium expressions: shows relationship between concentration of reactants and products

aA + bB → cC+ dD

Keq = [C]c[D]d/[A]a[B]b

• aqueous must be in mol/L

• treat as unitless

• when referring to reverse reaction, inverse

• pure solid/liquid present has concentrtion of 1 and disappears from equilibrium expression

big Keq (>10^3) favours products, small (<10^-3) favours reactants

RICE boxes

to determine Keq with incomplete data

Reaction: A → 2B

Initial conc: 3M 0

Change in conc: -x +2x

equilibrium: 3-x 2M

determine using quotient Q

aA + bB → cC+ dD

Q = [C]c[D]d/[A]a[B]b

compare Q and Keq

Q = K at equilibrium

Q < K forward reaction favoured

Q > K reverse reaction favoured

ΔG = 0 then Q = K

ΔG>0 then Q>K

ΔG<0 then Q<K

impact of temperature of Keq

more products = higher Keq

therefore impact of temperature depends on whether system is exothermic or endothermic, and therefore which way heating will shift

different uses of Keq

for ionic: dissociation of ionic compounds can be explored through Keq (Ksp)

for acid/bases: see module 6

solution equilibria

dissolution of ionic compounds in water

• water molecules are polar (oxygen partially negative, hydrogen partially positive)

• when ionic substances added to water, hydrogen atom is attracted to negatively charged ions, oxygen to positively charged = ion-dipole attraction

• when these ion-dipole attractions are strong enough, ions are removed from the ionic lattice and enter solution

• when an ion is surrounded by water = hydrated

• dissociation: solid ionic compound separates into charged ions

saturated solutions: unable to dissolve more solute at a particular temp

aboriginal people

detoxifying food works due to constant concentration gradient/le chatliers

cycads and other plants eg. yellow walnut: seeds contain the toxin cycasin, solubility of 56.6 g/L. toxins can be leached in multiple ways: roasted (breakdown chemical strucutre) and then leached in still or running water for at least 1-3 days

bitter yams (toxic oxalates) and saltbush also use ;leaching via soaking/placing in running water

conduct a practical to analyse composition of substances are mixed

potassium chloride and silver nitrte

potassium iodide and lead nitrate

sodium sulfate and barium nitrate

ksp

solubility product of dissociation

low ksp = not much solute can be dissolved before solution is saturated

can be used to calculate if a precipitate will form eg. by calculating Qsp, if greater than Ksp precipitate will form

common ion effect: greatly decreases the solubility of ionic compounds due to Le Chateliers

eg. AgCl added to NaCl, there are already Cl ions, so will shift away from products

quantitative analysis of an equilibrium reaction

(Fe3+HNO3)(aq) + KSCN-(aq) ⇌ [FeSCN]2+ (aq)

Keq = [[FeSCN]2+]/[Fe3+][SCN-]

turns red →

exothermic reaction in forward direction! therefore forward direction favours lower temp.

intro:

• can be challenging, most methods of detecting inadvertantly involve disturbing equilibrium eg. if you titrate, will react and remove chemical species of interest and equilibrium will shift and mol will be too high

• therefore UV spectrophotometry in situ used. involves using absorbance of photons by speciees of interest (in this case, iron thiocynate) to measure concentration

• prepare a series of standard solutions with excess Fe(NO3)3, so KSCN is limiting reagent. forced into completion, therefore in static equilibrium

• for the second part, variety of volumes were combined and allowed to reach dynamic equilibrium (not limiting reagent). recorded against curve

risk assessment:

• uv spectrometer: light can cause eye damage. keep lid closed whilst in operation.

• iron thiocynate: corrosive to skin, eye damage, inhalation

• nitric acid: corrosive (can severely irritate and burn the skin and eyes with possible eye damage), can irritate nose and throate, can irritate lungs (cause build up of fluid), nausea/vomiting/abdominal pain. not combustible but strong oxidiser that enhances combustion of other substances. explosion risk, releases dangerous gases in fire

• potassium thiocynate: irritate respiratory tract (long term exposure may result in dizziness/general weakness), may irritate eyes, skin, cause nausea. non-combustible

constructing a standard curve:

• 100x more FeNO3 than KSCN, to ensure all KSCN is consumed in the reaction. forces reaction RIGHT, only product and excess iron remain

• HNO3 solution is present to stabilise the Fe ions in the 3+ oxidation state. added in different volumes to ensure each solution ends up with the same total volume

questions:

• experiment run at 447nm = blue region colour. as solution is orange-red, absorbs this complimentary colour best

• NOTE: equilibrium position different to equilibrium constant.

  Equilibrium constant is the number that gives the relationship between amounts of products and reactant of a reaction mixture at its equilibrium. Equilibrium position is the moment at which the forward reaction of the equilibrium is equal to the backward reaction

• concentration and absorbance directly related

• Fe concentration much higher than KSCN to drive to completion

• reaction is exothermic forward ie. heat is a product. therefore warmer day will drive left and produce less FeSCN