module 5 Flashcards
dynamic equilibrium
rates of forwards and reverse reaction are equal but non-zero
- concentrations are constant
eg. saturated solution of sodium chloride
what pracs can you use for reversibility
reversibility of the dehydration of cobalt (II) chloride
- hexahydrate pink
- anhydroussky blue
reversibility of combustion metals –> shows the lack of reversibility
- combustion of magnesium and steel wool
- no changes occur when placed in ice bath
system
boundary of where the reaction occurs and composed of both substances and energy
- open
- closed
closed system
- constant number of particles in a system (no matter transfer)
- energy can be exchanged with the surroundings
eg. sauce pan and lid
open system
- system can interact with the surroundings allowing for the exchange of matter and energy
eg. boiling water without a lid
static equilibrium
the rates of forward and reverse reaction are equal and zero
- irreversible reaction at completion eg. dissolution of salt in unsaturated solution
- irreversible reaction before initiation
eg. combustion of fuel without th initial spark
- reversible reaction with insurmountable activation energy
eg. diamond and graphite
modelling dynamic equilibrium
counters or cards
enthalpy
amount. of stored heat energy within a substance
H reaction = h products - h reactants
- reactants > products (exo) (more energy for bond forming) forward
- reactants < products (endo) ( more energy for bond breaking) reverse
- kJ mol-1
-J mol-1
combustion: negative enthalpy change
photosynthesis: positive enthalpy change
entropy
measure of state of disorder in a chemical system
- J mol-1K
kJ mol-1K
delta s = sum of s products - sum of s reactants
S<0 reverse
S>0 forward
combustion: positive entropy change
photosynthesis: negative entropy change
difference between enthalpy and entropy
absolute value and change
law of thermodynamics
- Zeroth Law of Thermodynamics
- if two systems are in thermal equilibrium with a third system, then they are in thermal equilibrium with each other - The First Law of Thermodynamics
- energy movement into or out of a system is in accordance with the law of conservation of energy - The Second Law of thermodynamics
- entropy of an isolated system not at equilibrium will increase over time, approaching maximum value at equilibrium - The Third Law of Thermodynamics
- the entropy of a system approaches a minimum as the temperature approaches zero
what is the enthalpy and entropy change in reverse reactions
H>0
collision theory
- based on the principle that all matter is made up of tiny particles which are in constant motion and reactant particles have successful or unsuccessful collisions
- collision theory explains that chemical reactions take place when molecules with sufficient energy collide at a correct orientation (successful collision)
- high conc (reactants) – collide at high frequency –> rate of reactants formed high –> conc of reactants decreases –> forward reaction reduces over time
- increased conc of products as they are being formed –> collision of reverse increases
- reaches dynamic equilibrium
- conc of both is rarely equal
how to increase rate of reaction through collision theory
- frequency of collisions increase
- individual collisions must have a higher chance of being successful
rate of reverse reaction is proportional to
products
rate of forward reaction is proportional to
reactants
Le Chateliers Principle
if a system is at dynamic equilibrium is disturbed, then the system will shift so to minimise the change until a new equilibrium is reached
- change in: volume/pressure, concentration, temperature
- the change will never be completely nullified
shift to the right
forward reaction begins to exceed rate of reverse reaction
shift to the left
reverse reaction begins to excess rate of forward reaction
concentration LCP
- increase reactants –> forward reaction –> right by LCP
- increase products –> forward reaction –> left by LCP
what has no effect if added or removed within a system on the equilibrium
liquids and solids (solvents accepted)
what happens if water is added to a system
concentration decreases –> favours forward reaction
endo or exo
reactants + heat –> products
endo
delta H > 0
endo or exo
reactantas –> products + heat
exo
delta H < 0
temperature LCP
increases temp –> favours endo side (usually left reverse)
decreases temp–> favours exo side (usually right forward)
pressure/volume LCP
partial pressure of reactants pumped
3:2
- shifts to the right by LCP to lower the number of moles of substance per unit volume
volume decreases –> pressure increases –> concentration increases
- 3:2
- shifts to the rights by LCP to decrease the number of moles present
- inert gas has no effect on equilibrium –> make time taken to reach longer
partial pressure of a gas
hypothetical pressure if it were only the gas occupying the volume of the mixture, under the same conditions
total pressure of an ideal gas mixture
equal to the sum of the partial pressures of each component gas
inert gases
no effect on equilibrium
catalyst
substance which increases the rate of reaction by providing an alternate pathway with a lower activation energy
is a catalyst consumer or produced in a reaction
no
catalyst effect on equilibrium
- hastens attainment of equilibrium
- ratio of reactants and products is identical for catalysed and uncatlysed reactions
nitrogen dioxide
brown
dinitrogen tetroxide
colourless
equation for the interaction between nitrogen dioxide and dinitrogen tetroxide and colours and exo or endo
2NO2 –> N2O4
NO2 –> brown
N2O4 –> colourless
- exo
concentration via collision theory
- in a system at equilibrium, when the concentration of reactants is increased, there is an increased number of collisions between reactant particles –> rate of forward reaction increases as per collision theory
- As new product particles form due to the increased rate of reaction the reverse reaction also increases as the concentration of products increase
- over time a new equilibrium is re-established where the forward and reverse rates of reaction are equal
- overall the equilibrium shifted to the right by LCP
- reverse reaction is also higher than at old equilibrium
- means forward rate is higher
- therefore not all of new added reactant is removed
temperature via collision theory
- endothermic has higher activation energy for forward rather than reverse and exothermic has a lower activation energy for the forward reaction
- increasing temp increases rate of any reaction as the proportion of particles with enough energy to overcome the activation energy increases and collision frequency increases
- increasing temp has a greater effect on reaction rate when activation energy is high –> greater relative proportion of molecules will have enough energy to overcome activation energy
- thus increasing temp for endo leads to an increased forward and reverse rate
- however, since forward has higher activation energy, rate of forward is increase to greater than rate of reverse —> more product molecules
- therefore endo is favoured in the increase in temp
exothermic (activation energy)
lower for forward
endothermic (activation energy)
lower for reverse
pressure via collision theory
4:2 molar ratio
4:2
- if we decrease total volume –> increase total pressure
- reactant and product molecules are brought together –> increasing collision frequency –> increasing rate of reaction of both forward and reverse
- however, more moles of reactant gas –> more frequent collisions in reactant particles
- rate of forward exceed rate of reverse until enough product molecule is formed to make both rates equal
- more products were formed –> because it shifted right to the side with fewer moles of gas by LCP
addition of inert gas via collision theory
- wont change partial pressure of reactant and product gases
- collision between inert gas and molecules taking place in chemical reaction is unsuccessful
- addition of unsuccessful collisions will obstruct potentially sucessful
-reduce both forward and reverse rates of reaction - hence equilibrium will take a longer time to reach
equilibrium constant
quantitative relationship between concentration of reactants and
products at equilibrium
- dictates the position of equilibrium
- Keq is large = equilibrium position
is side of products ie. right
- Keq has a middling value = (0.1 -
10) = equilibrium exists as a
medium measure
- Keq is small = equilibrium position
is side of reactants ie. left
- K
- Keq
K = (products)^coefficients/(reactants)^coefficients - must be gases or aqueous
Q
concentrations of reactants and products for reaction is not yet at equilibrium
Q = (same equation as K)
- Q < K = system will proceed to the right to increase concentration of products (more reactants)
- Q = K = system is at dynamic equilibrium
- Q > K = system will proceed to the left to increase concentration of reactants (more products)
if pressure, volume or concentration is changed
equilibrium will sift but old Keq can still be satisfied
what is the only factor that can change the K value
temperature
exo
- increase in temp –> favours endo –> shift left by LCP–> more reactants –> smaller Keq
- decrease temp –> favours exo –> shifts right by LCP –> more products –> larger Keq
endo
- increase temp –> favours exo –> shifts right by LCP –> more products –> larger Keq
- decrease temp –> favours exo –> shifts left by LCP –> more reactants –> smaller Keq
3 ways to do ICE tables
- standard
- quadratic
- small change assumption
small change assumption
relative large number is added or subtracted to a relatively small number of x FIXXX
way to manipulate Keq
- reciprocal
- square
- add multiple equations together
catalyst via collision theory
- catalyst provides and alternate pathway for a reaction through lowering the activation energy
- allows for energetically weaker collisions which were deemed unsuccessful to be considered successful
- rate of forward and reverse simultaneously increase therefore relative concentration of products and reacts will not change at equilibrium
- therefore changes to activation energy will never effect the value of K eq however it hastens the time taken to reach it
lower activation energy of forward reaction
- keq will be larger –> right by LCP
iron (III) nitrate
pale-orange or yellow
Fe(NO3)3
lower activation energy of reverse
- keq will be smaller –> left by LCP
potassium thiocyanate
colourless
KSCN
iron (III) thiocyanate
blood red
determining Keq prac
iron (III) thiocynate –> FeSCN
ways to calculate keq
- ICE tables
- colorimetry
colorimetry
method of analysis which relates the absorbance of light of a specific wavelength by a colour solution to a quantitative measurement of the concentration of the solute
- standard solutions with a known solute at certain concentrations analysed with a colorimeter to determine absorbance
- relies on beer lambert
calibration curve axis
for colorimetry
x-axis: concentration
y-axis: absorbence
method of colorimetry
1.colorimeter passes light through a sample and measures the intensity of light of a specific wavelength received by the detector on the other side
2. gives a reading of the absorbency of the standard solution (unit less) between 0 - 1
3. performed number of times and calibration curve is formed
- sample solutions with the same solute can be analysed through finding the absorbance and the concentration an be read of the curve
what does colorimetry rely on
beer lambert law:
- absorbance is directly proportional to the concentration of a substance
- when absorbance is plotted against concentration –> expect linear relationship
change in gibbs free energy
- refers to the maximum amount of work that can be performed by a thermodynamic system during a chemical process
delta G = delta H -T delta S - usually J
delta G < 0 = spontaneous
- reaction will go to completion without an external energy input (usually a flame at the start) (exergonic)
delta G > 0 = non-spontaneous
- continuous energy input is required to force the reaction to occur (endergonic)
delta G = 0
- neither spontaneous or non-spontaneous
gibbs free and Keq
large negative delta G –> Keq is large –> positioned to the right
- forward spontaneous
rule of solubility
polar solutes dissolve in polar solvents
non-polar solutes dissolve in non-polar solvents
when will a solute dissolve
if the formation of intermolecular forces between the solute and the solvent are more favourable than the existing intermolecular forces between solute molecules and solvent molecules
dissolution
solute -solute bonds: broken (requires energy –> endothermic):
- ionic lattice held together by ionic bonds –> forces
of electrostatic attraction between positively and
negatively charged ions
solvent-solvent bonds: broken (requires energy –> endothermic):
- liquid state (water molecules are held together by
intermolecular forces eg. hydrogen bonding, dipole
- dipole interactions, dispersion forces)
solute-solvent interactions: formed (between ions and water molecules –> energy is released –> exothermic)
- solvation/hydration(case of water) –> solvent
molecules form hydration spheres around
individual ions
- the ion - dipole forces are formed between the
ions and the water molecules are. formed –>
stronger than hydrogen bonding in water
molecules
hydration
in aqueous solutions
- NaCL dissolves in water, slightly positive hydrogen in the polar H2O orients towards the Cl- anions while the slightly negative oxygens orient themselves towards the Na+ cations
salts that dissolve exothermically
CaCl2
NaCO3
NaOH
salts that dissolve endothermically
KCl
NaHCO3 (bi - carb)
NH4O3
entropy during dissolution
increases
- ionic lattice has low entropy –> ions are in a fixed position
- the hydrated ions are free to move –> higher entropy
how does dissolution effect the ionic lattice
once an ion from one corner or edge is broken off –> more edges and corners are formed –> increasing number of sites for dissolution
how to determine if dissolution of substance is static equilibrium or dynamic equilibrium
saturated (dynamic) or unsaturated solution (static)
what equilibrium is unsaturated solution
static
- heavily favoured forward reaction
what equilibrium is saturated solution
dynamic
- both rates of forward and reverse gradually become equal
- dissociation is occurring at the same rate as precipitation
- when solute added to saturated solution –> total mass of dissolved solute wont increase
- dynamic eq is still formed because ions are in constant motion
unsaturated
salt will completely dissolve
saturated
- maximum amount of salt has been dissolved to reach dynamic equilibrium
sufficient salt is added to dissolve - precipitate can form
- eq constant = ksp
supersaturated
formed when a solution is heated or cooled to allow for more solute to be dissolved than at standard conditions
what is the concentration of a saturated solution
solubility
what equilibrium is unsaturated solution
dynamic
- precipitate out any excess dissolved solute when standard conditions are re-esablished
if the concentration of a saturated solution is:
soluble: >10g/L
partially soluble: 1 - 10g/L
insoluble: <1g/L
relate gibbs free to dissolution
spontaneous (delta G < 0) = soluble salt
islander fruit
cycad –> toxin: cycasin (C8H16N2O7)
- carcinogenic
to detoxify
- crushed increases surface are
two methods of removing toxins:
- roasting (increases rate of reaction) –> leaching
- leaching for long periods of time
- water is drained –> resultant extracted –> pounded into fine powder again –. leached again
- repeated leaching and pounding ensures little cycasin left as possible
explain removing cycasin from cycad using solubility equilibrium
- cycasin (s) –> <— cycasin (aq)
- when seed is leached –> solid cycasin dissolves into the water in its aqueous form –> reaches dynamic eq
- water is drained –> dissolved toxin is removed (some solid cycasin may be left)
- process repeated with more clean water
- some of the toxin will produce more cycasin (aq) due to dynamic equilibrium
- the toxin will be leached out again by draining
- boiling water aids this –> increase temp –> increase rate of reaction
- more Cycasin (aq) toxin produced and drained away
solubility of cycasin
56.6 L G L -1
cobalt (II) chloride aqueous prac
endothermic
- mixture heated –> favours endo –> shifts right by LCP –> blue CoCl42-
- mixture cooled –> favours exo –> shifts left –> more pink Co(H2O062+
Ksp/ dissociation constant
tendency of a substance to dissociate ??
- larger –> more likely to dissociate to respective ions
- lower –> less likely to dissociate
precipitation
precipitation reaction involves two solutions mixing together resulting in the formation of an insoluble solid –> precipitate
precipitate
insoluble solid
eg. silver chloride, lead (II) iodide, barium sulfate
solubility rules
NAGSAG
N: nitrate
A: ammonium
G: group 1
S: sulfate except CaStroBar compounds
A: acetate
G: group 7
what happens when a slightly ionic substance dissolves in water in sufficient quantities
some dissolves while some remains as a solid
solubility measurements
- g/100ml
- mol L -1
ionic product
Qsp: when concentrations are not necessarily at equilibrium
Qsp < Ksp = unsaturated
- more ionic solid will dissolve if added (right)
Qsp = Ksp
- at equilibrium and saturated
Qsp > Ksp = solution may be supersaturated and ionic solid will precipitate (left)
common ion effect
- possible to dissolve salts into solutions that already contain one or more of the same dissolved ions
- decreases solubility of salts when dissolved in comparison to dissolution in pure water (DUE TO LCP increased conc of an ion causes the reverse reaction eg precipitation)
- if insoluble or sparingly soluble is added to the solution the concentration of the common ion does not change very much –> small change assumption
is there a limiting reagent in dynamic equilibrium
no
steps for finding K sp
- identify salt precipitated (least soluble)
- write equilibrium reaction for dissolution
- ice table
prac for precipitations
droppers and observe colours
lead (II) iodide precipitate
yellow
iron (II) hydroxide precipitate
green that tuns reddish brown as Fe2+ oxidises to Fe+
silver chloride precipitate
white
copper (II) hydroxide precipitate
blue-green
barium sulfate precipitate
white
reaction rate
dependent on frequency and success of collisions
