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module 5 Flashcards

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1
Q

dynamic equilibrium

A

rates of forwards and reverse reaction are equal but non-zero
- concentrations are constant

eg. saturated solution of sodium chloride

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2
Q

what pracs can you use for reversibility

A

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

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3
Q

system

A

boundary of where the reaction occurs and composed of both substances and energy
- open
- closed

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4
Q

closed system

A
  • constant number of particles in a system (no matter transfer)
  • energy can be exchanged with the surroundings
    eg. sauce pan and lid
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5
Q

open system

A
  • system can interact with the surroundings allowing for the exchange of matter and energy
    eg. boiling water without a lid
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6
Q

static equilibrium

A

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

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7
Q

modelling dynamic equilibrium

A

counters or cards

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8
Q

enthalpy

A

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

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9
Q

entropy

A

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

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10
Q

difference between enthalpy and entropy

A

absolute value and change

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11
Q

law of thermodynamics

A
  1. Zeroth Law of Thermodynamics
    - if two systems are in thermal equilibrium with a third system, then they are in thermal equilibrium with each other
  2. The First Law of Thermodynamics
    - energy movement into or out of a system is in accordance with the law of conservation of energy
  3. The Second Law of thermodynamics
    - entropy of an isolated system not at equilibrium will increase over time, approaching maximum value at equilibrium
  4. The Third Law of Thermodynamics
    - the entropy of a system approaches a minimum as the temperature approaches zero
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12
Q

what is the enthalpy and entropy change in reverse reactions

A

H>0

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13
Q

collision theory

A
  • 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
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14
Q

how to increase rate of reaction through collision theory

A
  • frequency of collisions increase
  • individual collisions must have a higher chance of being successful
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15
Q

rate of reverse reaction is proportional to

A

products

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16
Q

rate of forward reaction is proportional to

A

reactants

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17
Q

Le Chateliers Principle

A

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
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18
Q

shift to the right

A

forward reaction begins to exceed rate of reverse reaction

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19
Q

shift to the left

A

reverse reaction begins to excess rate of forward reaction

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20
Q

concentration LCP

A
  • increase reactants –> forward reaction –> right by LCP
  • increase products –> forward reaction –> left by LCP
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21
Q

what has no effect if added or removed within a system on the equilibrium

A

liquids and solids (solvents accepted)

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22
Q

what happens if water is added to a system

A

concentration decreases –> favours forward reaction

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23
Q

endo or exo
reactants + heat –> products

A

endo
delta H > 0

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24
Q

endo or exo
reactantas –> products + heat

A

exo
delta H < 0

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25
Q

temperature LCP

A

increases temp –> favours endo side (usually left reverse)
decreases temp–> favours exo side (usually right forward)

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26
Q

pressure/volume LCP

A

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
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27
Q

partial pressure of a gas

A

hypothetical pressure if it were only the gas occupying the volume of the mixture, under the same conditions

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28
Q

total pressure of an ideal gas mixture

A

equal to the sum of the partial pressures of each component gas

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29
Q

inert gases

A

no effect on equilibrium

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30
Q

catalyst

A

substance which increases the rate of reaction by providing an alternate pathway with a lower activation energy

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31
Q

is a catalyst consumer or produced in a reaction

A

no

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32
Q

catalyst effect on equilibrium

A
  • hastens attainment of equilibrium
  • ratio of reactants and products is identical for catalysed and uncatlysed reactions
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33
Q

nitrogen dioxide

A

brown

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34
Q

dinitrogen tetroxide

A

colourless

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35
Q

equation for the interaction between nitrogen dioxide and dinitrogen tetroxide and colours and exo or endo

A

2NO2 –> N2O4
NO2 –> brown
N2O4 –> colourless
- exo

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36
Q

concentration via collision theory

A
  • 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
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37
Q

temperature via collision theory

A
  • 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
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38
Q

exothermic (activation energy)

A

lower for forward

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39
Q

endothermic (activation energy)

A

lower for reverse

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40
Q

pressure via collision theory
4:2 molar ratio

A

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

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41
Q

addition of inert gas via collision theory

A
  • 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
42
Q

equilibrium constant

A

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
43
Q

Q

A

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)

44
Q

if pressure, volume or concentration is changed

A

equilibrium will sift but old Keq can still be satisfied

45
Q

what is the only factor that can change the K value

A

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

46
Q

3 ways to do ICE tables

A
  • standard
  • quadratic
  • small change assumption
47
Q

small change assumption

A

relative large number is added or subtracted to a relatively small number of x FIXXX

48
Q

way to manipulate Keq

A
  • reciprocal
  • square
  • add multiple equations together
49
Q

catalyst via collision theory

A
  • 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
50
Q

lower activation energy of forward reaction

A
  • keq will be larger –> right by LCP
51
Q

iron (III) nitrate

A

pale-orange or yellow
Fe(NO3)3

51
Q

lower activation energy of reverse

A
  • keq will be smaller –> left by LCP
52
Q

potassium thiocyanate

A

colourless
KSCN

52
Q

iron (III) thiocyanate

A

blood red

53
Q

determining Keq prac

A

iron (III) thiocynate –> FeSCN

54
Q

ways to calculate keq

A
  • ICE tables
  • colorimetry
55
Q

colorimetry

A

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

56
Q

calibration curve axis

A

for colorimetry
x-axis: concentration
y-axis: absorbence

57
Q

method of colorimetry

A

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

58
Q

what does colorimetry rely on

A

beer lambert law:
- absorbance is directly proportional to the concentration of a substance
- when absorbance is plotted against concentration –> expect linear relationship

59
Q

change in gibbs free energy

A
  • 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

60
Q

gibbs free and Keq

A

large negative delta G –> Keq is large –> positioned to the right
- forward spontaneous

61
Q

rule of solubility

A

polar solutes dissolve in polar solvents
non-polar solutes dissolve in non-polar solvents

62
Q

when will a solute dissolve

A

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

63
Q

dissolution

A

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

64
Q

hydration

A

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

65
Q

salts that dissolve exothermically

A

CaCl2
NaCO3
NaOH

66
Q

salts that dissolve endothermically

A

KCl
NaHCO3 (bi - carb)
NH4O3

67
Q

entropy during dissolution

A

increases
- ionic lattice has low entropy –> ions are in a fixed position
- the hydrated ions are free to move –> higher entropy

68
Q

how does dissolution effect the ionic lattice

A

once an ion from one corner or edge is broken off –> more edges and corners are formed –> increasing number of sites for dissolution

69
Q

how to determine if dissolution of substance is static equilibrium or dynamic equilibrium

A

saturated (dynamic) or unsaturated solution (static)

70
Q

what equilibrium is unsaturated solution

A

static
- heavily favoured forward reaction

71
Q

what equilibrium is saturated solution

A

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

72
Q

unsaturated

A

salt will completely dissolve

73
Q

saturated

A
  • maximum amount of salt has been dissolved to reach dynamic equilibrium
    sufficient salt is added to dissolve
  • precipitate can form
  • eq constant = ksp
74
Q

supersaturated

A

formed when a solution is heated or cooled to allow for more solute to be dissolved than at standard conditions

75
Q

what is the concentration of a saturated solution

A

solubility

76
Q

what equilibrium is unsaturated solution

A

dynamic
- precipitate out any excess dissolved solute when standard conditions are re-esablished

77
Q

if the concentration of a saturated solution is:

A

soluble: >10g/L
partially soluble: 1 - 10g/L
insoluble: <1g/L

78
Q

relate gibbs free to dissolution

A

spontaneous (delta G < 0) = soluble salt

79
Q

islander fruit

A

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
80
Q

explain removing cycasin from cycad using solubility equilibrium

A
  • 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
81
Q

solubility of cycasin

A

56.6 L G L -1

82
Q

cobalt (II) chloride aqueous prac

A

endothermic
- mixture heated –> favours endo –> shifts right by LCP –> blue CoCl42-
- mixture cooled –> favours exo –> shifts left –> more pink Co(H2O062+

83
Q

Ksp/ dissociation constant

A

tendency of a substance to dissociate ??
- larger –> more likely to dissociate to respective ions
- lower –> less likely to dissociate

84
Q

precipitation

A

precipitation reaction involves two solutions mixing together resulting in the formation of an insoluble solid –> precipitate

85
Q

precipitate

A

insoluble solid
eg. silver chloride, lead (II) iodide, barium sulfate

86
Q

solubility rules

A

NAGSAG
N: nitrate
A: ammonium
G: group 1
S: sulfate except CaStroBar compounds
A: acetate
G: group 7

87
Q

what happens when a slightly ionic substance dissolves in water in sufficient quantities

A

some dissolves while some remains as a solid

88
Q

solubility measurements

A
  • g/100ml
  • mol L -1
89
Q

ionic product

A

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)

90
Q

common ion effect

A
  • 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
91
Q

is there a limiting reagent in dynamic equilibrium

A

no

92
Q

steps for finding K sp

A
  • identify salt precipitated (least soluble)
  • write equilibrium reaction for dissolution
  • ice table
93
Q

prac for precipitations

A

droppers and observe colours

94
Q

lead (II) iodide precipitate

A

yellow

95
Q

iron (II) hydroxide precipitate

A

green that tuns reddish brown as Fe2+ oxidises to Fe+

96
Q

silver chloride precipitate

A

white

97
Q

copper (II) hydroxide precipitate

A

blue-green

98
Q

barium sulfate precipitate

A

white

99
Q

reaction rate

A

dependent on frequency and success of collisions

chemistry (9 decks)

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